study guide-for-chemical-principles-5-by-atkins-slideshare

ChemistryAtkins 5th Edition

Ch.0 - Intro to General Chemistry

Ch.1 - Atoms: The Quantum World

Ch.2 - Chemical Bonds

Ch.3 - Molecular Shape & Structure

Ch.4 - The Properties of Gases

Ch.5 - Liquids & Solids

Ch.6 - Inorganic Materials

Ch.7 - Thermodynamics: The First Law

...and more!

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CONCEPT: TEMPERATURE

Temperature is a measure of thermal energy in a substance, which is __________________ on the amount of matter.

Heat is a form of thermal energy, which is __________________ on the amount of matter.

EXAMPLE 1: Which of the following has the greatest amount of heat?

30 g H2O at 50oC 300 g H2O at 50oC 30 g H2O at 100oC

Temperature Conversions

Temperature conversions are easy, as long as you know how to solve for x.

• You only need to know 2 equations to convert from the 3 different units of temperature:

K = 273.15 + oC

oF = 1.8 (oC) + 32

EXAMPLE 2: Convert the following units of temperature

(1) -115°C into Fahrenheit (2) 73.23 K into Fahrenheit

CHEMISTRY - ATKINS 5E

CH.0 - INTRO TO GENERAL CHEMISTRY

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16. Which of the following is an intensive property of platinum metal? a) The molar mass of platinum is 195.08 g/mol.

b) The volume of 110 grams of platinum is 5.14 mL.

c) The density of platinum is 21.4 g/mL.

d) 0.800 moles of platinum.





17. Which of the following is an extensive property?

a) Hardness b) Color c) Volume d) Melting Point e) Temperature





5. At what temperature is the numerical value the same, where the units are in Celsius or Fahrenheit?

a. 32o

b. 0 o

c. – 40 o

d. – 273 o





CONCEPT: GROUP NAMES AND CLASSIFICATIONS

Ever wonder where did this periodic table ever come from?

• At the end of the 18th century, Lavoisier compiled a list of the 23 elements known at the time.

• In 1869, Dmitri Mendeleev coined the term “Periodic Table” .

• Today the total is 114 and still counting!

Now, to understand chemistry fully it will be imperative that you memorize and learn the different portions of the Periodic

Table.

Phase Differences

At room temperature (between 20 oC to 25oC), all elements are _______________ except:

• Mercury and bromine are _______________ . • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine and the Noble Gases are ____________.




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CONCEPT: CHARGE DISTRIBUTIONS OF THE PERIODIC TABLE

A majority of the elements on the periodic table are reactive because they all want to be like the _____________________.

• They have the perfect number of electrons in their outer atomic shells.

1. Metals tend to __________ electrons to become positively charged ions called _____________.

• Metals that have ONLY one charge are referred to as ________________ metals.

• Metals that have MORE THAN one charge are referred to as ________________ metals.

2. Nonmetals tend to __________ electrons to become negatively charged ions called _____________.





CONCEPT: ELEMENT SYMBOLS

Some of the names and symbols for the elements are easy to recognize like Aluminum is Al, but some others aren’t.

EXAMPLE 1: Identify the elements by their given symbols.

a. Au b. Hg c. Pb d. Fe e. Ag

Some elements exist in nature connected to their exact double.

We call these chemical Siamese twins ________________________________.

To recall them just remember this funny phrase:

Have No Fear Of Ice Cold Beer

Some elements exist in nature as monoatomic elements such as _______________ & _______________.

Some elements exist in nature as polyatomic molecules such as _______________ & _______________.





7. Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound. a) Iodine b) CH3OH c) Lead

d) NaClO2





CONCEPT: STRUCTURE OF THE ATOM

We learned that the basic functional unit in chemistry is the _____________ .

• Now it’s time to go into an atom to figure out its components: _________ subatomic particles.

In the center of an atom there is the _____________,

• It contains the subatomic particles: _____________ and _____________.

• Spinning around it we find the third subatomic particle: the _____________.

• PROTONS are _________________ charged subatomic particles.

• ELECTRONS are _________________ charged subatomic particles.

• NEUTRONS are _________________ charged subatomic particles.

ATOMIC NUMBER equals the number of _________________ and determines _________________ of an element.

ATOMIC MASS equals the number of ____________________________ in an element.

EXAMPLE: Identify the unknown element.

a. Element X (8 protons, 8 electrons, 8 neutrons) b. Element Y (35 protons, 36 electrons, 46 neutrons) c. Element Z (12 protons, 10 electrons, 13 neutrons)

!




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CONCEPT: MODERN ATOMIC THEORY

According to the Law of ________________________________ in a reaction matter is neither created nor destroyed.

• Originated in 1789 by Antoine Lavoisier.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

According to the Law of ________________________________ all samples of a compound, no matter on their origin or

preparation has the same ratio in terms of their elements.

• Originated in 1797 by Joseph Proust.

Mass Ratio = 12.0gC)((32.0gO)

= 0.375

According to the Law of ________________________________ when two elements (A & B) form different compounds, the

masses of element B that combine with 1 g of A are a ratio of whole numbers.

• Originated in 1804 by John Dalton.

Mass Ratio = 16.0gO)((14.0gN )

=1.143


= 2.286

The ratio of the two mass ratios obtained then gives us a whole number:

2.2861.143

= 2.0

CO2

NO

NO2




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CONCEPT: MODERN ATOMIC THEORY (PRACTICE)

EXAMPLE 1: A 15.39 g sample of iodine reacts with 62.92 g of chlorine to form iodine pentachloride, ICl5. If iodine

pentachloride is the only product formed calculate its mass.

EXAMPLE 2: Two samples sodium fluoride decompose into their constituent elements. The first sample produces 15.8 kg

of sodium and 20.1 kg of fluorine. If the second sample produces 192.0 g of sodium, how many grams of fluorine were also

produced?

PRACTICE: Which of the following is an example of the law of multiple proportions?

a. A sample of bromine (Br) contains equal amounts of its two isotopes.

b. Two different samples of H2O have the same mass ratio.

c. The atomic mass of sodium (Na) is 22.99 amu.

d. Two different compounds composed of sulfur (S) and oxygen (O) have different mass ratios: 2.48 g O: 1 g S and

1.24 g O: to 1 g S.




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CONCEPT: STRUCTURE OF THE ATOM

We learned that the basic functional unit in chemistry is the _____________ .

• Now it’s time to go into an atom to figure out its components: _________ subatomic particles.

In the center of an atom there is the _____________,






ATOMIC NUMBER equals the number of _________________ and determines _________________ of an element.

ATOMIC MASS equals the number of ____________________________ in an element.

EXAMPLE: Identify the unknown element.

a. Element X (8 protons, 8 electrons, 8 neutrons) b. Element Y (35 protons, 36 electrons, 46 neutrons) c. Element Z (12 protons, 10 electrons, 13 neutrons)

!




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CONCEPT: MODERN ATOMIC THEORY

According to the Law of ________________________________ in a reaction matter is neither created nor destroyed.

• Originated in 1789 by Antoine Lavoisier.

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g)

According to the Law of ________________________________ all samples of a compound, no matter on their origin or

preparation has the same ratio in terms of their elements.

• Originated in 1797 by Joseph Proust.

Mass Ratio = 12.0gC)((32.0gO)

= 0.375

According to the Law of ________________________________ when two elements (A & B) form different compounds, the

masses of element B that combine with 1 g of A are a ratio of whole numbers.

• Originated in 1804 by John Dalton.


=1.143


= 2.286

The ratio of the two mass ratios obtained then gives us a whole number:

2.2861.143

= 2.0

CO2

NO

NO2




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CONCEPT: ATOMIC MASS

Whether you call it atomic mass or weight both terms tell us the combined mass of the protons and neutrons in an element.

• The atomic masses listed for the elements on the periodic table are the ____________________ of their isotopes.

• Isotopes are elements with the _______________ number of protons, but _______________ number of neutrons.

Atomic Mass = [(Mass of Isotope 1) x (Fractional Abundance 1)] + [(Mass of Isotope 2) x (Fractional Abundance 2)]

EXAMPLE 1: Antimony has two common isotopes. If one of the isotopes 121Sb has an isotopic mass of 120.9038 amu and

a natural abundance of 57.25%, what is the isotopic mass (to 4 significant figures) of the other isotope? The atomic mass of

antimony is 121.8 g/mol.

EXAMPLE 2: The atomic mass of an imaginary element A is 251.7 amu. If element A consists of two isotopes that have

atomic masses of 250 and 253 respectively, what is the natural abundance of each isotope?




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CONCEPT: MASS CONVERSIONS

The _______________ is the chemical unit for the amount of a substance.

One mole (1 mol) contains 6.022 x 1023 entities, which is known as _________________________________________.

• Entities means ______________________ , ______________________ or ______________________.

• We use ______________________ when dealing with a single, individual element.

• We use _________________ or _________________when dealing with more than one element or a compound.

6.022 x 1023 atoms of Fe is equal to 1 mole of Fe and has a mass of 55.85 amu

Atoms Moles Grams

EXAMPLE: Determine the mass (in grams) found in 7.28 x 1028 nitrogen atoms.

6.022 x 1023 molecules of H2O is equal to 1 mole of H2O and has a mass of 18.016 amu

Molecules Moles Grams

EXAMPLE: Determine how many molecules of carbon dioxide, CO2, are found in 75.0 g CO2.





1. How many atoms of oxygen are there in 53.2 g Al2(SO4)3?





2. How many moles of C3H8 contain 4.95 x 1024 hydrogen atoms?





9. Which of the following amounts has the largest mass? a) 1.186 x 1023 atoms I

b) 5.932 x 1022 molecules I2

c) 3.954 x 1022 molecules I3

d) All have relatively the same mass.





CONCEPT: MASS CONVERSIONS (PRACTICE)

PRACTICE 1: If the density of water is 1.00 g/mL at 25oC calculate the number of water molecules found in 1.50 x 103 µL

of water.

PRACTICE 2: Calculate the number of oxygen atoms found in 783.9 g CuSO4 · 5 H2O.

PRACTICE 3: The density of the sun is 1.41 g/cm3 and its volume is 1.41 x 1027 m3. How many hydrogen molecules are in

the sun if we assume all the mass is hydrogen gas?

PRACTICE 4 (CHALLENGE): A cylindrical copper wire is used for the fences of a house. The copper wire has a diameter

of 0.0750 in. How many copper atoms are found in 5.160 cm piece? The density of copper is 8.96 g/cm3. ( V = π · r2 · h ).





CONCEPT: STOICHIOMETRIC REACTIONS

2 H2 (g) + 1 O2 (g) 2 H2O (g)

In the above equation the numbers that are in bold are called _______________________.

• They tell us the number of ______________ of each compound that reacts.

• This numerical relationship between compounds in a balanced equation is called __________________________.

STOICHIOMETRIC CHART

Before we get into solving stoichiometric reactions lets work out a plan of attack.

Entities means ______________________ , ______________________ or ______________________.

Entities of Given

Grams of Given

Entities of Unknown

Grams of Unknown

Moles of Given Moles of Unknown

Use this chart when given a chemical equation with the ____________ quantity of a compound or element and asked to find

the ____________ quantity of another compound or element.

EXAMPLE: How many grams of H2O are produced when 12.3 g H2 reacts?

2 H2 (g) + 1 O2 (g) 2 H2O (g)




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The reaction of Potassium chlorate and sucrose is given below: 8 KClO3 + C12H22O11 8 KCl + 12 CO2 + 11 H2O

a) If 76.13 grams of potassium chlorate reacted with excess sucrose, C12H22O11, how many moles of water would be produced?

b) How many oxygen atoms from sucrose reacted if 235.23 g of carbon dioxide were produced?





PRACTICE: STOICHIOMETRIC REACTIONS

EXAMPLE 1: The oxidation of chromium solid is represented by the following equation:

4 Cr (s) + 3 O2 (g) 2 Cr2O3 (s)

a. How many moles of chromium (III) oxide are produced when 34.69 g Cr reacts with excess oxygen gas?

b. How many grams of O2 were needed to produce 4.28 x 103 molecules Cr2O3?

EXAMPLE 2: If the density of ethanol, CH3CH2OH, is 0.789 g/mL, how many milliliters of ethanol are needed to produce 4.8

g of H2O in the following reaction?

CH3CH2OH (l) + 3 O2 2 CO2 (g) + 3 H2O (l)

PRACTICE: Dinitrogen monoxide gas decomposes to form nitrogen gas and oxygen gas. How many molecules of oxygen

are formed when 8.00 g of dinitrogen monoxide decomposes?





CONCEPT: LIMITING REAGENT

In a chemical reaction the reactant that is consumed when a reaction occurs and determines the maximum amount of

product formed is called the _________________________________.

• The amount of product it forms is called the _________________________ yield.

The reactant that remains after the completion of the chemical reaction is called the ____________________ reactant.

EXAMPLE: Chromium (III) oxide reacts with hydrogen sulfide (H2S) gas to form chromium (III) sulfide and water:

Cr2O3 (s) + 3 H2S (g) Cr2S3 (s) + 3 H2O (l) [Balanced]

a. What is the mass of chromium (III) sulfide formed when 14.20 g Cr2O3 reacts with 12.80 g H2S?

b. Identify the limiting reactant, excess reactant and theoretical yield.

c. What mass of excess reactant remains?





The reaction between aluminum and iron (III) oxide can generate temperatures reaching 3000oC and is used in welding metals:

2 Al + Fe2O3 Al2O3 + 2 Fe

a) If 150 g of Al are reacted with 432 g of Fe2O3 what is the mass of Al2O3 produced?

b) Identify the limiting reactant, excess reactant and theoretical yield.





Based on the answer of the previous question, how much of the excess reactant is left at the end of the reaction?





CONCEPT: PERCENT YIELD

The percent yield of a reaction is used to determine how effective the chemist was in creating their desired products.

• A high percent yield would signify that the reaction is _______________________________________.

EXAMPLE 1: A scientist performs an experiment in the laboratory and obtains 13.27 g Cr2S3. If his calculations on scratch

paper give him a theoretical yield of 18.23 g what is the percent yield?

PRACTICE: Consider the following balanced chemical reaction:

2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (l)

a. If a 2.6 g sample of C6H6 reacted with excess O2 to produce 1.25 g of water, what is the percent yield of water?

b. If the above reaction only went to 75% completion, how many moles of CO2 would be produced if 1.57 x 10-5 molecules of C6H6 were reacted with excess oxygen?





Many kinds of metals react with oxygen gas to form metal oxides. For example, calcium reacts as follows:

2 Ca (s) + O2 (g) 2 CaO (s) If the percent yield for the above reaction is 67.33%, how many grams of O2 gas would be required to produce 37.51 g of calcium oxide?





CONCEPT: MOLARITY

Molarity (M) can serve as the connection between the interconversion of ____________ to ____________ and vice versa.

For example, a 5.8 M NaCl solution really means __________________________ per __________________________.

Molarity =MolesSolute)(LitersSolution)(

A typical mixture consists of a smaller amount of one substance, the ________________, dissolved in a larger amount of

another substance, the __________________. Together they form a __________________.




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20. Which of the following solutions will have the highest concentration of chloride ions? a) 0.10 M KCl

b) 0.10 M SrCl2

c) 0.10 M AlCl3

d) 0.05 M PbCl4

e) All of these solutions have the same concentration of bromide ions.





21. Calculate the molarity of bromide ions in a solution if you mixed 12.86 g calcium bromide,

CaBr2, in enough water to make 305.0 mL of solution.





PRACTICE: MOLARITY

EXAMPLE 1: 2.64 grams of an unknown compound was dissolved in water to yield

150 mL of solution. The concentration of the solution was 0.075 M. What was the

molecular weight of the substance?

EXAMPLE 2: A solution is prepared by dissolving 0.1408 mol calcium nitrate, Ca(NO3)2, in enough water to make 100.0 mL

of stock solution. If 20.0 mL of this solution is then mix with an additional 90 mL of deionized water, calculate the

concentration of the calcium nitrate solution.

PRACTICE 1: What is the molarity of calcium ions of a 650 mL solution containing 42.7 g of calcium phosphate?

PRACTICE 2: A solution with a final volume of 750.0 mL was prepared by dissolving 30.00 mL of benzene (C6H6, density =

0.8787 gmL

) in dichloromethane. Calculate the molarity of benzene in the solution.

M =MolesSolute)(LitersSolution)(





CONCEPT: MOLARITY & CHEMICAL REACTIONS

Whenever we are provided given information in a reaction we use ___________________ to find any unknown information.

• In aqueous reactions, this given information is typically in units of __________________ or __________________ .

Entities means ______________________ , ______________________ or ______________________.

Volume of Given Moles of Given Moles of Unknown Volume of Unknown

Entities of Unknown

Grams of Unknown

Use this chart when given a chemical equation with the known quantity in either ________ or ________ of a compound or

element and asked to find the unknown quantity of another compound or element.

EXAMPLE: How many grams of sodium metal are needed to react with 38.74 mL of 0.275 M NaOH?

2 Na (s) + 2 H2O (l) H2 (g) + 2 NaOH (aq)





PRACTICE: MOLARITY & CHEMICAL REACTIONS

PRACTICE 1: How many milliliters of 0.325 M HCl are needed to react with 16.2 g

of magnesium metal?

2 HCl (aq) + Mg (s) MgCl2 + H2 (g)

PRACTICE 2: What is the molarity of a hydrobromic acid solution if it takes 34.12 mL of HBr to completely neutralize 82.56

mL of 0.156 M Ca(OH)2?

2 HBr (aq) + Ca(OH)2 (aq) CaBr2 (aq) + 2 H2O (l)

PRACTICE 3 (CHALLENGE): Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation:

Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O

If it takes 30.0 mL of 0.100 M K2Cr2O7 to titrate a 25 mL Fe2+ solution, what is the molar concentration of Fe2+?

M =MolesSolute)(LitersSolution)(





CONCEPT: AQUEOUS SOLUTIONS

The ___________________________ of a compound represents the maximum amount of solute that dissolves in a solvent.

SOLUBILITY RULES

SOLUBLE IONIC COMPOUNDS

INSOLUBLE IONIC COMPOUNDS

1. Group 1A ions (Li+, Na+, K+, etc.) and ammonium ion (NH4+) are soluble.

1. (Hydroxides) OH- and (Sulfides) S2-, are insoluble

except when with Group 1A ions (Li+, Na+, K+, etc.),

ammonium ion (NH4+) and Ca2+, Sr2+, Ba2+.

2. (Nitrates) NO3- , (acetates) CH3COO- or C2H3O2-,

and most perchlorates (ClO4-) are soluble.

2. (Carbonates) CO32- and (Phosphates) PO43- are

insoluble except when with Group 1A ions

(Li+, Na+, K+, etc.), ammonium ion (NH4+).

3. Cl- , Br- , and I- are soluble, except when paired

with Ag+ , Pb2+ , Cu+ and Hg22+.

4. (Sulfates) SO42- are soluble, except those of Ca2+ ,

Sr2+ , Ba2+ , Ag+ , and Pb2+ .

When we classify a compound as soluble it means that the compound is _______________________, it is also known as

a(n) _______________________ because it conducts electricity.

NaNO3 (s) H2O

Na+ (aq) + NO3– (aq)

When we classify a compound as insoluble it means that the compound is a _______________________, it is also known

as a(n) _______________________ because it doesn’t conduct electricity.

CH3OH (l) H2O

CH3OH (aq)

BaSO4 (s) H2O

BaSO4 (aq)




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35. Which of the following reagents could be used to separate the two anions from a solution containing magnesium nitrate and cesium hydroxide? a) NH4CN

b) NaCl

c) KNO3

d) ZnBr2

e) CsBrO3





36. Which of the following reagents could be used to separate the two cations from a solution containing Lead (IV) acetate and cesium permanganate? a) Sr(NO3)2

b) TiC2H3O2

c) K2S

d) NaClO4

e) KNO3





CONCEPT: ELECTROLYTES

Whenever we add a solute into a solvent three outcomes are possible:

• the solute will _________________ dissolve ( STRONG electrolytes).

• the solute will _________________ dissolve ( WEAK electrolytes).

• the solute will _________________ dissolve ( NON electrolytes).

Classification of Solutes in Aqueous Solution

STRONG ELECTROLYTES

WEAK ELECTROLYTES

NONELECTROLYTES

1. STRONG ACIDS: HCl, ______ , HI ,

HNO3 , _______ , _______ , _______ .

2. STRONG BASES:

Group 1A Metal with OH-, H-, O2- or

NH2-

Groups 2A Metal, Calcium or Lower, with

OH-, H-, O2- or NH2-

3) SOLUBLE IONIC COMPOUNDS:

1. WEAK ACIDS: HF, ____________ ,

________ , ________ , ________ .

2. WEAK BASES: Be(OH)2 , Mg(OH)2 ,

_________ , _________ .

1. MOLECULAR

COMPOUNDS:

______________

C6H12O6 (glucose)

C12H22O11 (sucrose)

______________





PRACTICE: ELECTROLYTES

EXAMPLE: Each of the following reactions depicts a solute dissolving in water. Classify each solute as a strong electrolyte,

a weak electrolyte or a non-electrolyte.

a. PbSO4 (s) PbSO4 (aq)

b. HC2H3O2 (aq) H+ (aq) + C2H3O2– (aq)

c. CaS (s) Ca2+ (aq) + S2- (aq)

d. Hg (l) Hg (aq)

PRACTICE: Classify each of the following solutes as either a strong electrolyte, a weak electrolyte or a non-electrolyte.

a. Perbromic acid, HBrO4

b. Lithium chloride, LiCl

c. Formic Acid, HCO2H

d. Methylamine, CH3NH2

e. Zinc bromide, ZnBr2

f. Propanol, C3H8OH





33. Which of the following is NOT a strong electrolyte? a) NaOH b) BaCl2 c) MgSO4 d) LiNO3 e) K2SO4





34. Which of the following is considered a STRONG electrolyte? a) NH4NO3 b) Be(OH)2 c) PbCl2 d) HF e) CH3OH





CONCEPT: WRITING CHEMICAL REACTIONS

EXAMPLE: Predict whether a reaction occurs, and write the balanced molecular equation.

a. LiOH (aq) + MgSO4 (aq)

EXAMPLE: Predict whether a reaction occurs, and write the balanced molecular equation, the total and net ionic equations.

Molecular: Na2CO3 (aq) + HBr (aq)

Total Ionic:

Net Ionic:





PRACTICE: Predict whether a reaction occurs, and write the balanced total and net ionic equations.

Total Ionic:

Net Ionic:

Molecular: Ag2SO4 (aq) + KCl (aq)

PRACTICE: Predict whether a reaction occurs, and write the balanced total and net ionic equations.

Total Ionic:

Net Ionic:

Molecular: MgBr2 (aq) + NaC2H3O2 (aq)





CONCEPT: OXIDATION-REDUCTION REACTIONS

Chemists use some important terminology to describe the movement of electrons.

• In ______________ reactions we have the movement of electrons from one reactant to another.

L

E

O

G

E

R

Agent Agent

Rules for Assigning an Oxidation Number (O.N.)

A. General Rules

1. For an atom in its elemental form (Na, O2, S8, etc.): O.N. = 0

2. For an ion the O.N. equals the charge: Na+ , Ca2+ , NO3 –

B. Specific Rules

1. Group 1A: O.N. = +1

2. Group 2A: O.N. = +2

3. For hydrogen: O.N. = +1 with nonmetals

O.N. = -1 with metals and boron

4. For Fluorine: O.N. = -1

5. For oxygen: O.N. = -1 in peroxides (X2O2 , X = Group 1(A) element)

O.N. = − 12

in superoxides (XO2 , X = Group 1(A) element)

O.N. = - 2 in all other compounds

6. Group 7A O.N. = -1 (except when connected to O)




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CONCEPT: OXIDATION-REDUCTION REACTIONS (PRACTICE)

EXAMPLE: In the following reaction identify the oxidizing agent and the reducing agent:

a. 2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (g)

PRACTICE: What is the oxidation number of each underlined element?

a. P4 b. BO33-

c. AsO42- d. HSO4

–

PRACTICE: In the following reaction identify the oxidizing agent and the reducing agent:

a. Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O





6. Determine the best possible answer(s) for each of the following questions based on the periodic table.

Which of the following statements is/are true? a) Phosphorus is a homonuclear diatomic molecule.

b) Zinc is a Type I Metal

c) Tin is a Type II metal.

d) Europium (Eu) is an actinide metal

Which nonmetal exists as a diatomic liquid at room temperature?

Bromine Tellurium Sulfur Chlorine Iodine




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6. Determine the best possible answer(s) for each of the following questions based on the periodic table.

Which of the following statements is/are true? a) Phosphorus is a homonuclear diatomic molecule.

b) Zinc is a Type I Metal

c) Tin is a Type II metal.

d) Europium (Eu) is an actinide metal

Which nonmetal exists as a diatomic liquid at room temperature?

Bromine Tellurium Sulfur Chlorine Iodine





8. The barium content of a metal ore was analyzed several times by a percent composition process.

Method

Barium (weight %)

1

0.012

0.016

0.010

0.012

a) Calculate the mean, median and mode. b) Calculate the standard deviation.





13. Determine which of the following represents a physical change and a chemical change.

a)

b)

c)





14. Answer each of the following questions based on the images provided below.

a) Which of the following images represents an elemental gas? b) Which of the following images represents a heterogeneous mixture? c) Which of the following images represents a liquid?





15. Which of the following represents a physical property? a) Mercury is a liquid at room temperature.

b) Alkanes burn spontaneously.

c) CO2 (s) CO2 (g)

d) 2 H2 (g) + O2 (g) 2 H2O (g)

e) The rusting of a car.





18. A cigarette lighter contains the substance butane, C4H10, with the given properties. a) The lighter contains 10 g of butane b) the density of butane is 0.57 g/cm3 c) The freezing point of the butane is -138oC d) Butane is combustible in air. Which of the following features is a chemical change? Which of the following properties are intensive properties?





19. Which of the following properties about an unknown metal represent chemical properties? I. It has a bluish-green color. II. Upon exposure to air the metal experiences corrosion. III. It has a density of 4.36 g/cm3. IV. It has a boiling point of 522oC. V. It conducts electricity.

a. IV b. V c. I, II, III d. II e. IV and V





25. An empty flask has a volume of 250 mL. When water is placed inside the flask (d = 0.9983 g/mL) the volume is 124.3 mL. What is the mass (in grams) of the water inside the flask?





26. If a statue at an art gallery is covered in 279 kg of copper. If the copper on the statue has a thickness of 0.0055 cm,

what surface area is covered (in square meters)? Copper has a density of 8.96 g/cm3.





30. A Volkswagen diesel engine consumes diesel at a rate of 20.17 L per hour. If the density of the diesel is 0.850

g/mL. What is the mass (in mg) of diesel needed to drive for a continuous 4.3 hours?





32. If the charge and mass of one proton is 1.60218 x 10-19 C and 1.673 x 10-30 g respectively, what is the charge of

378 kg of protons.





32. A cylindrical tube has a length that is 15.2 cm and is filled with bromine liquid. If the mass of the bromine is 112.3 g and has a density of 25.3 g/mL. What is the inner diameter of the glass tube? ( V = π• r2 •h )





33. A large body of water contaminated with lead has an average depth of 750 m, a total area of 1.25 x 105 km2, and

an average of 5.8 x 10-12 g/L of dissolved lead. How many milligrams of lead are in this large body of water?





34. Calculate the mass (in grams) of a golden sphere with a diameter of 30.0 mm. The density of gold is 19.3 g/cm3.

(Vsphere = π ! r3).





3. How many molecules of hexane are contained in 55.0 mL of hexane? The density of hexane is 0.6548 g/mL and the molar mass is 86.17 g/mol.





4. How many SO3 ions are contained in 120.0 mg of Na2SO3? The molar mass of Na2SO3 is 126.05 g/mol.





5. What mass of phosphorus pentafluoride, PF5, has the same number of fluorine atoms as 50.0 g of oxygen difluoride, OF2?





6. How many bromide ions are there in 4.50 moles of gallium bromide?





7. How many moles of oxygen atoms are required to combine with 3.05 moles of Pb to create lead (IV) phosphate?





8. How many cations are there in 100.0 g of lithium nitride?





10. Which of the following amounts would contain the least atoms? a) 10.0 g Sr

b) 10.0 g Br

c) 10.0 g Mg

d) 10.0 g Li





11. Which of the following amounts have the most molecules?

a) 15.0 g N2

b) 15.0 g Br2

c) 15.0 g O2

d) 15.0 g I2





16. How many milligrams of NaCN are required to prepare 712 mL of 0.250 M NaCN?




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17. What volume (in µL) of 0.100 M HBr contains 0.170 moles of HBr?





18. How many moles of Ca2+ ions are in 0.100 L of a 0.450 M solution of Ca3(PO4)2?





19. How many chloride ions are present in 65.5 mL of 0.210 M AlCl3 solution? a) 4.02 × 1023 chloride ions

b) 5.79 × 1024 chloride ions

c) 2.48 × 1022 chloride ions

d) 8.28 × 1021 chloride ions

e) 1.21 × 1022 chloride ions





22. To what final volume would 100 mL of 5.0 M KCl have to be diluted in order to make a solution that is 0.54 M KCl?





23. If 880 mL of water is added to 125.0 mL of a 0.770 M HBrO4 solution what is the resulting molarity?





PRACTICE: REACTION PATHWAYS (CALCULATIONS)

PRACTICE 1: In each of the following reactions determine if the catalyst (the one in bold) is either homogeneous or

heterogeneous?

a. C2H4 (g) + H2 (g) C2H6 (s)

b. H2O (l) + SO3 (g) H2SO4 (aq)

PRACTICE 2: What is the rate law to the elementary steps given?

Cl2 (g) + NO (g) Cl2NO (g) [SLOW]

Cl2NO (g) + NO (g) 2 ClNO (g) [FAST]

PRACTICE 3: What is the rate law from the following elementary steps?

NO (g) + Br2 (g) NOBr2 (g) [FAST]

NOBr2 (g) + NO (g) 2 NOBr (g) [SLOW]

a) Rate = k [NO] [Br]

b) Rate = k [NOBr2] [NO]

c) Rate = k [NO]

d) Rate = k [NO]2 [Br2]

e) Rate = k [NOBr]2





26. Consider the following balanced redox equation:

H2O + 2 MnO4 – + 3 SO32- 2 MnO2 + 3 SO42- + 2 OH – How many grams of MnO2 (MW: 86.94 g/mol) are produced when 32.0 mL of 0.615 M MnO4- (MW: 118.90 g/mol) reacts with excess water and sulfite?





27. Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation:

Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O

If it takes 35.0 mL of 0.250 M FeCl2 to titrate 50 mL of a solution containing Cr2O72-, what

is the molar concentration of Cr2O72-?





28. Vinegar is a solution of acetic acid, CH3COOH, dissolved in water. A 5.54 g sample of

vinegar was neutralized by 30.10 mL of 0.100 M NaOH. What is the percent by weight of

acetic acid in the vinegar?





29. What is the molar mass of a 0.350 g sample of a monoprotic acid if it requires 50.0 mL of 0.440 M Ca(OH)2 to completely neutralize it?





30. Give the complete ionic equation for the reaction (if any) that occurs when aqueous solutions of sodium sulfide and copper (II) nitrate are mixed. a) Na+ (aq) + SO42-(aq) + Cu+(aq) + NO3-(aq) CuS(s) + Na+(aq) + NO3-(aq)

b) Na+ (aq) + S-(aq) + Cu+(aq) + NO3-(aq) CuS(s) + NaNO3(aq)

c) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) Cu2+(aq) + S2-(aq) + 2 NaNO3(s)

d) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) CuS(s) + 2 Na+(aq) + 2 NO3-(aq)

e) No reaction occurs.





31. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed. a) H+(aq) + OH-(aq) H2O(l)

b) 2 K+(aq) + SO42-(aq) K2SO4(s)

c) H+(aq) + OH-(aq) + 2 K+(aq) + SO42-(aq) H2O(l) + K2SO4(s)

d) H22+(aq) + OH-(aq) H2(OH)2(l)






32. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of Na2CO3 and HCl are mixed. a) 2 H+(aq) + CO32-(aq) H2CO3(s)

b) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 NaCl(aq)

c) 2 H+(aq) + CO32-(aq) H2O(l) + CO2(g)

d) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq)






CONCEPT: SCIENTIFIC NOTATION

We use scientific notation to turn small or large, inconvenient numbers into manageable ones.

6.88 x 10-12

The first number 6.88 is called the __________________________.

It must be greater than or equal to 1 and less than 10.

The second number is known as the ______________________.

It must always be 10 in scientific notation.

In the number 6.88 x 10-12, the number -12 is referred to as the ____________________.

EXAMPLE 1: Convert the following numbers into scientific notation

a. 377,000 b. 0.000101

c. 707.82 d. 161.00 x 107

e. 0.0628 x 10-9


CH.1 - ATOMS: THE QUANTUM WORLD


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CONCEPT: TYPES OF ERRORS

Even though we try to be as accurate as possible, there is always some level of uncertainty called ____________________.

When we investigate the quality of an experimental decision or calculation we take into account two major principles:

• __________________ deals with the reproducibility of our calculations.

• __________________ deals with how close our measured calculation is to the “actual” value.

EXAMPLE: Which of the 4 following images is not precise and not accurate?

EXAMPLE: A student must measure the weight of a sodium bicarbonate compound, NaHCO3, and obtains the following measurements: 23.12 g, 23.08 g and 23.17g. If the true weight of the compound is 18.01 g what can be said about the student’s results?

a. They are accurate and precise. b. They are accurate, but not precise. c. They are not accurate, but precise. d. They are neither accurate or precise.




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2. A student must measure the weight of a sodium bicarbonate compound and obtains the following measurements: 23.12 g, 23.26 g, 23.08 g and 23.17g. If the true weight of the compound is 18.01 g what can be said about the student’s results?

a. They are accurate and precise. b. They are accurate, but not precise. c. They are not accurate, but precise.

d. They are neither accurate or precise.




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CONCEPT: EVALUATING YOUR MEAN VALUE

The ____________________________ measures how close data results are in relation to the mean, or average, value.

• To check if the results are close to the “true” value we can merely look, but sometimes determining their accuracy may require more work.

Σ(x1 − x )2

n−1

x =

x =

n =




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CONCEPT: METRIC PREFIXES

pico nano micro milli centi deci 0 deca hecto kilo mega giga tera

(p) (n) (µ) (m) (c) (d) (da) (h) (k) (M) (G) (T)

10-12 10-9 10-6 10-3 10-2 10-1 100 101 102 103 106 109 1012




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3. Convert the following to the desired units.

a) 5.62 pL to L b) 407 Gg to cg





4. Convert the following to the desired units.

a) 1.17 x 10-6 ns to ks b) 3.77 x 103 dm3 to µm3





CONCEPT: SIGNIFICANT FIGURES Significant figures are necessary to communicate the level of accuracy with which values are recorded. It can be easy if you remember these 2 rules.

1. If your number has a decimal point move from ______________________________. o Start counting once you get to your first non-zero number and keep counting until you get to the end.

2. If your number has NO decimal point move from ______________________________.

o Start counting once you get to your first non-zero number and keep counting until you get to the end.

EXAMPLE 1: How many sig figs does each number contain?

(1) 0.0000185 m (2) 749 mol (3) 17.3 x 103 mL (4) 100. min (5) 0.0010050 kg (6) 1560 mol

EXAMPLE 2: Read the length of the metal bar to the correct number of significant figures.

a) 15 cm b) 15.000 cm c) 20 cm d) 15.0 cm e) 15.00 cm




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1. What is the SI unit for each of the following? a) Mass kilograms grams milligrams pounds

b) Length miles kilometers meters inches

c) Time hours seconds minutes days

d) Temperature Celsius Fahrenheit Kelvin Degrees





11. Determine the correct number of significant figures for each of the following images.

a) 5 mL b) 5.3 mL c) 5.32 mL d) 5.320 mL e) 5.3200 mL





CONCEPT: CALCULATIONS & SIGNIFICANT FIGURES

• Multiplication/Division: Measurement with least _______________________________ determines final answer.

• Addition/Subtraction: Measurement with least _______________________________ determines final answer. EXAMPLE 1: Perform the following calculation to the right number of sig figs:

(3.16) x (0.003027) x (5.7 x 10-3) EXAMPLE 2: Perform the following calculation to the right number of sig figs:

2.628 x 106

6.281 x 104

+ 0.827 x 107

EXAMPLE 3: Perform the following calculation to the right number of sig figs:

(42.00 – 40.915) • (25.739 – 25.729) (11.50 • 1.001) + (0.00710 • 700.)





9. Determine the answer to the correct number of significant figures.

(8.123− 0.891) ⋅ (0.053+ 0.947)(0.312− 0.23)

a) 90.4 b) 9 c) 9.0 x 101 d) 9.04 x 101 e) 9 x 101





10. Determine the answer to the correct number of significant figures.

(19.4+13.7+ 0.0092+1503.017)(103.6+ 4.)(0.24+ 2.0046)

x10−5 + (6.039 x10−2 )

a) 1 b) 2 c) 3 d) 4 e) 5





CONCEPT: CONVERSIONS

Length 1 km = 0.6214 miles 1 m = 1.094 yards 1 in = 2.54 cm 1 mile = 5280 ft

Volume 1 gallon = 3.785 L 1 L = dm3 1 mL = 1 cm3 1 L = 1.057 qt

Mass 1 kg = 2.205 lbs 1 lb = 454 g 1 oz = 28.35 g

EXAMPLE 1: Every Saturday morning Gregor has to travel from Main Campus to his parents’ home. If his car gets 58.5 km/L how many L will his car need to travel the 19.3 miles? EXAMPLE 2: A backyard swimming pool holds 315 cubic yards (yd3) of water. What is the mass of the water in pounds? PRACTICE: An intravenous bag delivers medication to a patient at a rate of 2.75 drops a second. If a drop weighs 42 mg, how many grams of solution are delivered in 7.0 hours?




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27. The area on a piece of standard paper is 93.5 square inches. What is the area of the paper in units of cm2?





CONCEPT: DIMENSIONAL ANALYSIS We use dimensional analysis as a fail proof process to convert from one unit to another.

- Design the problems to ______________ with your known, and to _____________ with the unit of your unknown. - Be sure ALL of your units cancel out! Just follow the units.

EXAMPLE 1: Natty Light contains 4.2% alcohol. Steve from Kappa Epsilon Gamma, KEG for short, wants to get tanked tonight, and he is aiming to down at least 175 ml of alcohol in one night. If each can of Natty light contains 355 mL of beer, how many cans of Natty Light must Steve consume at minimum to reach his goal? EXAMPLE 2: A Volkswagen diesel engine consumes diesel at a rate of 25.83 L per hour. If the density of the diesel is 0.850

g/mL, what is the mass (in mg) of diesel needed to drive for a continuous 8.5 hours?

PRACTICE: An acetaminophen suspension for toddlers contains 95 mg/0.85 mL suspension. The recommended dose is 22 mg/kg body weight. How many liters of this suspension should be given to a toddler weighing 30.5 lbs?





28. During the annual New York Marathon it is suggested that runners drink about 1.2 quarts of water for every 2.2 miles. How many gallons of water would a person require if they wish to run 42.17 kilometers?





29. If a person is taking a prescription of Betaxol after surgery and the recommended dosage is 50.0 mg/kg of body mass. Calculate the recommended dosage for a person who weighs 100 lbs.





CONCEPT: DENSITY

We use density to understand the relationship between ______________ and ______________.

• We can use it within dimensional analysis to go from one unit to the other or vice versa.

Density =

EXAMPLE 1: If the density of an unknown metal is 21.4 g/cm3, express its density in lb/ft3. Remember that 1 in = 2.54 cm.

EXAMPLE 2: A piece of unknown metal weighs approximately 0.45 lbs. When a scientist places it in a glass beaker the water level increases from 1.85 L to 2.13 L. What is the density of the palladium metal in g/mL?

PRACTICE: The U.S. Environmental Protection Agency sets the maximum safe level of lead in blood at 24 µg per dL of blood. The average person has 60 mL of blood per kilogram of body weight. For a 63.7-kg (140.459 lb) person, what is the total maximum safe content of lead in blood?





23. An empty container weighs 16.55 g. When the container is filled with an unknown liquid it weighs 37.82 g. If the

unknown liquid has a volume of 5.0 mL. What is the density of the unknown liquid in g/cm3?





24. An unknown compound is placed into a 500 mL graduated cylinder that contains 130 mL of water. If the solid weighs 55.3 grams and causes the final volume to be 198.7 mL, will the solid float or sink in the graduated cylinder?





CONCEPT: THE NATURE OF LIGHT

Visible light represents a small portion of the continuum of radiant energy known as _______________________________.

The visible light spectrum ranges from ______________ to ______________ .

Its wave properties of electromagnetic radiation are described by two independent variables:

_________ (ν, Greek mu) is the number of waves you have per second and is expressed in units of ______ or ________.

__________ (λ, Greek lambda) is the distance from one crest of a wave to the other and is expressed in units of _______.

Relationship between frequency & wavelength




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PRACTICE: THE NATURE OF LIGHT

A. Based on the images of different electromagnetic waves, answer each of the following questions.

I. II.

III.

a) Which electromagnetic wave has the longest wavelength?

b) Which electromagnetic wave has the greatest energy?

c) Which electromagnetic wave has the lowest frequency?

d) Which electromagnetic wave has the largest amplitude?





7. Place the following types of electromagnetic radiation in order of increasing wavelength. UV light Gamma Rays Microwave Blue Light Violet Light





CONCEPT: INTERCONVERSION OF LIGHT UNITS

The speed of a wave, is the product of ν and λ. In a vacuum, all forms of electromagnetic radiation travel at 3.00 x 108 ,

which is a physical constant called the _________________________________ (c).

c = ν · λ

EXAMPLE: Even the music we listen to deals with how energy travels to get to our car radio. If Power 96 broadcasts its

music at 96.5 MHz (megahertz, or 106 Hertz) find the wavelength in μm and Ao

of the radio waves.

PRACTICE: Calculate the frequency of the red light emitted by a neon sign with a wavelength of 663.8 nm.





CONCEPT: ENERGY AND MATTER

Light travels at different speeds as it passes through different media in a phenomenon known as _____________________.

Light passing through the opening of a slit creates a semicircular wave in a phenomenon known as ___________________.

If the light wave passes through two adjacent slits then the semicircular waves can interact with one another .

• ___________________ inteference ______ amplitude. � ___________________ inteference ______ amplitude.




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CONCEPT: THE PHOTOELECTRIC EFFECT

Albert Einstein theorized that light was quantiized into small “packets” or “bundles” of energy.

• A single particle of this quantized “packet” of electromagnetic energy was later named a

________________.

According to the Photoelectric effect, when photons with enough energy hit the surface of a metal electrons are emitted.

– Energy is directly proportional to ____________________ rather than its ____________________.

– The Photoelectric Effect only happens with photons over a certain _______________ frequency.

EXAMPLE: Illustrate what happens when a photon of sufficient energy strike the surface of a metal.

Real-life Application:




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CONCEPT: THE WAVE NATURE OF LIGHT

Up to this point we have discussed light as “packets” or particles of energy that travel through a given space, now we will

look at light as it travels as a uniform wave through a given space.

According to the ______________________ equation matter behaves as though it moves in a wave. To calculate the

wavelength of matter we simply use the following equation:

EXAMPLE: Find the wavelength (in nm) of a proton with a speed of 7.33 x 109 . (Mass of an proton = 1.67 x 10-27 kg)

PRACTICE: What is the speed of an electron that has a wavelength of 895 μm? (Mass of a electron = 9.11 x 10-31 kg)

λ =hmν

λ =

h =

m =

ν =




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6. What is the speed, in ms

, of an electron that has a wavelength of 3.13 x 105 pm? (Mass of electron = 9.11 x 10-31

kg).




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CONCEPT: QUANTUM MECHANICAL PICTURE OF THE ATOM

The main atomic sub-levels are the s, p, d and f. Each atomic sub-level has a set number of atomic or electron orbitals.

Each electron orbital can hold up ________ electrons.

The s sub-level contains one electron orbital _______

The p sub-level contains three electron orbitals

_______ _______ _______

The d sub-level contains five electron orbitals

_______ _______ _______ _______ _______

The f sub-level contains seven electron orbitals

_______ _______ _______ _______ _______ _______ _______




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CONCEPT: QUANTUM NUMBERS OF AN ATOMIC MODEL

An atomic orbital is characterized by three quantum numbers.

The __________________ quantum number deals with the atomic orbital’s size and energy. It tells us the relative distance

of the electron from the nucleus. It uses the variable ___________ and provides the shell number of the electron.

EXAMPLE: Calculate the principal quantum number of each atomic sublevel.

a. 7p b. 5s c. 3d d. 4f

The electron capacity of each shell can be determined by using the formula: ____________________ .

Electron Shell (n) Maximum Number of Electrons

1

2

3

4

The _______________________ quantum number deals with the shape of the atomic orbital. Each atomic orbital has a

specific shape.

• It uses the variable ___________ and formula _______________________.

Each atomic sub-level has an L value associated with it.

Sublevel s p d f g

L value 0 1 2 3 4

The ________________________________ quantum number deals with the orientation of the orbital in the space around

the nucleus. It is a range of the previous quantum number: -l to +l. It uses the variable ___________.

Sublevel s p d f

L value 0 1 2 3

ML value





PRACTICE: QUANTUM NUMBERS OF AN ATOMIC MODEL

EXAMPLE 1: What l or ml values are allowed if n = 2? How many orbitals exist for n = 2?

EXAMPLE 2: How many electrons can have the following quantum sets?

a) n = 4

b) n = 3, l = 1

c) n = 4, mL = -2

d) n = 5, l = 2, mL = -2

PRACTICE 1: Provide the n, l and ml value for each of the given orbitals.

a. 6p n = l = mL =

b. 4d n = l = mL =

c. 5f n = l = mL =

PRACTICE 2: State all the l and mLvalues possible if the principle quantum number is equal to 3.





14. How many electrons can have the following quantum sets? a) In the 7th shell of an atom. (n = 7) b) n = 5, l = 2

c) n = 6, l = 3, mL = -2

d) n = 4, l = 2, mL = 0 , mS = – 12





CONCEPT: THE ATOMIC MODEL

An atom is composed of __________ subatomic particles.

In the center of an atom there is the ________________ .






________ Model helped to explain what happened when an electron absorbed or released energy within a hydrogen atom.

After the hydrogen electron absorbed sufficient energy and becomes __________ it would jump to a higher energy level.

• Eventually it would return to its _____________________ and release the energy it absorbed as heat or light.

!




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PRACTICE: THE ATOMIC MODEL

EXAMPLE: Calculate the energy of the 4th electron found in the n = 2 state of the boron atom in kilojoules per mole.

PRACTICE 1: Which of the following transitions (in a hydrogen atom) represents emission of the longest wavelength?

a) n = 4 to n = 2 b) n = 3 to n= 4 c) n = 1 to n = 2 d) n = 6 to n = 5 e) n = 2 to n = 5

PRACTICE 2: Which of the following transitions represents absorption of a photon with the largest energy?

a) n = 3 to n = 1 b) n = 2 to n = 4 c) n = 1 to n = 2 d) n = 6 to n = 3 e) n = 1 to n = 4





5. Determine the end (final) value of n in a hydrogen atom transition, if the electron starts in n = 4 and the atom releases a photon of light with an energy of 4.084 × 10-19 J.





CONCEPT: ATOMIC EMISSION

When an electron absorbs enough energy it goes from a ___________ numbered shell to a ___________ numbered shell.

• The electron eventually releases or emits the energy it took in and goes from a ___________ numbered shell to a

___________ numbered shell.

If the electron goes from a higher numbered shell to the 1st shell it is referred to as a _____________________ Series.

1

∞

If the electron goes from a higher numbered shell to the 2nd shell it is referred to as a _____________________ Series.

2

∞

If the electron goes from a higher numbered shell to the 3rd shell it is referred to as a _____________________ Series.

3

∞





PRACTICE: ATOMIC EMISSION

EXAMPLE: What is the wavelength of a photon (in nanometers) emitted during a transition from n = 4 to n = 2 state in the

hydrogen atom?

PRACTICE: Classify each of the following transitions as either a Lyman, Balmer or Paschen series.

a) n = 3 to n = 1 b) n = 6 to n = 1 c) n = 3 to n = 2 d) n = 6 to n = 3 e) n = 4 to n = 2





4. What is the wavelength of a photon (in nanometers) absorbed during a transition from n = 5 to n = 2 state in the hydrogen atom?





CONCEPT: ELECTRON CONFIGURATIONS

In this chapter we will focus on how an element’s ________________________________________ - the distribution of

electrons within the orbitals of its atoms – relates to its chemical and physical properties.

History Lesson: In 1870, Dmitri Mendeleev arranged 65 elements into a ___________________________________ .

• He summarized their behavior in the _______________________________.

• When arranged by atomic mass, the elements exhibit a periodic recurrence of similar properties.

The Electron Configuration

According to the _______________ Principle you first have to totally fill in the lowest energy level before moving to the next.

1s 2s 2p

1s 2s 2p

Hund’s Rule states that electron orbitals that are _______________________ are first half-filled before they are totally filled.

F (9 electrons)

1s2s$$$$$$2p3s$$$$$$3p$$$$$$3d4s$$$$$$4p$$$$$$4d$$$$$4f5s$$$$$$5p$$$$$$5d$$$$$5f$$$$$5g6s$$$$$$6p$$$$$$6d$$$$$6f$$$$$6g$$$$6h$7s$$$$$$7p$$$$$$7d$$$$$7f$$$$$7g$$$$7h




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20. Which electron configuration represents a violation of the Pauli Exclusion Principle?

a. 1s 2s 2p

b. 1s 2s 2p

c. 1s 2s 2p

d. 1s 2s 2p

e. 1s 2s 2p





21. Which electron configuration represents a violation of Hund’s Rule?

a.

b.

c.

d.

e.





CONCEPT: CONDENSED ELECTRON CONFIGURATION

EXAMPLE: Write the condensed configuration for each of the following elements:

a. Co (27 electrons)

b. Se (34 electrons)

PRACTICE: Write the condensed configuration for each of the following elements:

a. Ag (47 electrons)





CONCEPT: PARAMAGNETISM Vs. DIAMAGNETISM

EXAMPLE: Write the condensed electron configuration of each ion and state if the ion is paramagnetic or diamagnetic.

a. Ni3+

b. S2-

PRACTICE: Write the condensed electron configuration of each ion and state if the ion is paramagnetic or diamagnetic.

a. Cu+





23. Give the electron configuration for the following element and its ion. For the ion, state if it is paramagnetic or diamagnetic:

a. S

S2-






a. Mo

Mo3+





CONCEPT: INNER CORE & VALENCE ELECTRONS

EXAMPLE: How many core (inner) and valence electrons are present in each of the following elements?

a. P

b. Al

c. Mn





CONCEPT: ATOMIC ORBITAL SHAPE The _______________________ quantum number deals with the shape of the atomic orbital. Each atomic orbital has a

specific shape.

• It uses the variable ___________ and formula _______________________.

Each atomic sub-level has an L value associated with it.

Sublevel s p d f g

L value 0 1 2 3 4

EXAMPLE: Based on the following atomic orbital shape, which of the following set of quantum numbers is correct:

a) n = 8, l = 1, ml = 12

b) n = 8, l = 2, ml = -2

c) n = 8, l = 0, ml = 1

d) n = 8, l = 0, ml = 0

PRACTICE: Based on the following atomic orbital shape, which of the following set of quantum numbers is correct:

a) n = 2, l = 1, ml = +1 , ms = - 1

b) n = 4, l = 1, ml = - 2 , ms = +12

c) n = 3, l = 1, ml = - 1, ms = 0

d) n = 2, l = 1, ml = + 1 , ms = – 12




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CONCEPT: TRENDS IN ATOMIC RADIUS

Atomic radius is defined as half the distance between the nuclei in a molecule of two identical elements.

• Generally, it ____________ going from left to right across a period and ______________ going down a group.

ATOMIC RADIUS

EXAMPLE: If the sum of the atomic radii of diatomic carbon is 154 pm and of diatomic chlorine is 198 pm, what is the sum

of the atomic radii between a carbon and a chlorine atom.

PRACTICE: Which one of the following atoms has the largest atomic radius?

A) K B) Rb C) Y D) Ca E) Sr




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CONCEPT: TRENDS IN IONIZATION ENERGY

Metals tend to lose electrons to become positive ions called ____________ .

• Therefore they have ____________ ionization energies.

Nonmetals tend to gain electrons to become negative ions called ___________ .

• Therefore they have ______________ ionization energies.

Ionization energy (IE) is the energy (in kJ) required to remove an electron from a gaseous atom or ion.

• Generally, it ________________ going from left to right of a period and ________________ going down a group.

Atom (g) ion+ (g) + e – ∆E = IE1 > 0

Exceptions:

• When in the same period, Group ______ elements have lower ionization energy than elements in Group ______ .

O 1s 2s 2p 1s 2s 2p

N 1s 2s 2p 1s 2s 2p

• When in the same period, Group ______ elements have lower ionization energy than elements in Group ______ .

B 1s 2s 2p 1s 2s 2p

Be 1s 2s 1s 2s

IONIZATION ENERGY





PRACTICE: TRENDS IN IONIZATION ENERGY

EXAMPLE: Of the following atoms, which has the smallest second ionization energy?

a. Al b. Li c. Rb d. Mg e. Be

PRACTICE 1: Of the following atoms, which has the smallest third ionization energy?

a. Al b. Ca c. K d. Ga e. Cs

PRACTICE 2: Which of the following statements is/are true?

a. Sulfur has a larger IE1 than phosphorus

b. Boron has a lower IE1 than Magnesium

c. Magnesium has a higher IE1 than Aluminum

PRACTICE 3: Shown below are the numerical values for ionization energies (IE’s). Match the numerical values with each of

the following elements provided in the boxes.

Na Mg Al Si P S Cl Ar

Numbers: 496, 578, 738, 786, 1000, 1012, 1251 & 1521.





CONCEPT: TRENDS IN ELECTRON AFFINITY

Electron Affinity (EA) is the energy change (in kJ) from the addition of 1 mole of e – to 1 mol of gaseous atoms or ions.

• Generally, it ________________ going from left to right across a period and ______________ going down a group.

Atom (g) + e – ion – (g) ∆E = - EA1

ELECTRON AFFINITY

EXAMPLE: Rank the following elements in order of increasing electron affinity.

a. Cs, Hg, F, S

b. Se, S, Si

PRACTICE: Shown below are the numerical values for electron affinities (EA’s). Match the numerical values with each of the following elements provided in the boxes.

Li Be B C N O F Ne

Numbers: - 328, -141, -122, -60, -27, > 0, > 0, > 0.





CONCEPT: TRENDS IN IONIC RADIUS

Ionic Size estimates the size of an ion in an ionic compound.

__________________ (POSITIVE IONS) tend to be smaller than their parent atoms.

Lithium ( 3 Electrons)

1s 2s 1s 2s

__________________ (NEGATIVE IONS) tend to be larger than their parent atoms.

Fluorine ( 9 Electrons)

1s 2s 2p 1s 2s 2p

The pattern for ionic size correlates with the following trend when comparing ions with the same number of electrons:

-3 > -2 > -1 > 0 > +1 > +2 > +3

EXAMPLE: Rank each set of ions in order of increasing ionic size.

a) K+ , Ca2+, Ar

b) Sr2+, Na+, I –

c) V5+, S2-, Cl –





8. Which of the following transitions (in a hydrogen atom) represent emission of the smallest or shortest wavelength?

a. n = 4 to n = 2

b. n = 3 to n= 4

c. n = 1 to n = 2

d. n = 7 to n = 5

e. n = 2 to n = 5





9. Which of the following transitions represent absorption of a photon with the highest frequency? a. n = 3 to n = 1

b. n = 2 to n = 4

c. n = 1 to n =2

d. n = 6 to n = 3

e. n = 1 to n = 3





10. Provide the n, l and ml value for each of the given orbitals. a) 7s n = b) 5d n =

l = l =

ml = ml =

c) 2p n = d) 4f n =

l = l =

ml = ml =





11. Which statement about the four quantum numbers is false? a. n = principal quantum number, n = 1 to ∞

b. l = azimuthal quantum number, l = 0,1,2, . . ., (n+1)

c. mL = magnetic quantum number, mL = (-l), . . .,0,. . ., (+l)

d. ms = spin quantum number, ms = + 12or − 1

2

e. The first three quantum numbers deal with the atomic orbitals except for the ms quantum

number, which deals with the electrons in the atomic orbitals.





12. Each of the following sets of quantum numbers gives information on a specific orbital. Find the error in each.

a. n = 4, l = 0 , ml = 1, ms = – 12

b. n = 5, l = 2 , ml = - 1, ms = 1

c. n = 7, l = 7, ml = - 5, ms = – 12

d. n = 0, l = 5, ml = - 3, ms = 12





14. How many electrons can have the following quantum sets? a) n = 4, mL = -1

b) n = 5, mL = 0 , mS = – 12

c) n = 9, l = 4, mS = – 12

d) n = 2, mS = 12





19. For n = 2, what are the possible sublevels? a) 0

b) 0, 1

c) 0, 1, 2

d) 0, 1,2, 3





16. Based on the following atomic orbital shape, which of the following set of quantum numbers is correct:

a) n = 2, l = 1, ml = 0

b) n = 3, l = 2, ml = –1

c) n = 4, l = 0, ml = +1

d) n = 1, l = 1, ml = 0






a) n = 3, l = 2, ml = 0, ms = – 12

b) n = 3, l = 1, ml = - 3, ms = 1

c) n = 4, l = 0, ml = 0, ms = – 12

d) n = 4, l = 2, ml = - 3, ms = 12






a) n = 3, l = 3, ml = 0, ms = 12

b) n = 1, l = 3, ml = -3, ms = 1

c) n = 7, l = 3, ml = - 4, ms = – 12

d) n = 6, l = 3, ml = -3, ms = – 12






a. Ag

Ag+






a. Cl

Cl+





27. Which of the following represents an “excited” state?

a) Cl: 1s22s22p63s23p5

b) Be: 1s22s2

c) Na: 1s22s2-2p63p1

d) N: 1s22s22p3





28. Give the set of four quantum numbers that represent the indicated electron in the following element:

a. Br (33rd electron) n = , l = , ml = , ms =






a. Ca (19th electron) n = , l = , ml = , ms =






a. Cu (27th electron) n = , l = , ml = , ms =






a. Mo3+ (38th electron) n = , l = , ml = , ms =





32. For a multi-electron atom, arrange the electron subshells of the following listing in order of increasing energy: 6s, 4f, 2p, 5d.





CONCEPT: MATTER

Chemistry is the study of matter and the changes it undergoes, with the _________________ being its basic functional unit. When two or more of these elements chemically bond together they form an independent structure called a molecule.

Classification of Matter

Under appropriate conditions of pressure and temperature, most substances can exist in 3 states of matter:

• _______________, _______________ and ________________.

• _______________ have a fixed shape and volume.

• ______________ take up the shape and volume of a container.

• _______________ conform to the shape of a container, but not the volume.

__________________ __________________ __________________

Microscopic Explanation for the Behavior of Gases, Liquids and Solids

Gas Liquid Solid

Assumes the

___________ and ___________

of its container.

Assumes the __________ of the

portion of its container it occupies,

but not the __________ .

Maintains a fixed

___________ and ___________

___________________ compressible

Viscosity Viscosity Viscosity

____________________ Viscous

___________________ compressible ___________________ compressible

____________________ Viscous ____________________ Viscous


CH.2 - CHEMICAL BONDS


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CONCEPT: MIXTURES

Most matter consists of mixtures of pure substances.

• __________________________ mixtures have no indistinguishable parts.

• _____________________________ mixtures have some distinguishable parts.

Changes of matter

• ______________________ changes are changes in the form of the substance, but not in its chemical composition.

• ______________________ changes create new substances with different properties and different chemical

compositions.

PRACTICE: Which of the following represents a physical change?

a) Alkanes burn spontaneously.

b) The sublimation of CO2.

c) 2 H2 (g) + O2 (g) 2 H2O (g)

d) The rusting of a car.





22. Which of the following is not a mixture? a. Air

b. Gasoline

c. Sodium

d. Bronze





20. Which of the following descriptions represents a physical property for a given nonmetal? a. It undergoes combustion.

b. It vaporizes to a gas at 80oC.

c. It reacts with a metal to create an ionic compound.

d. None of the above.





21. Which of the following physical changes involves the releasing of energy? a. Freezing of liquid water. b. Fusion of ice. c. Vaporization of water. d. All of the above. e. None of the above.





CONCEPT: EMPIRICAL FORMULA

The empirical formula is also known as the __________________________.

• It represents the _______________________________ ratio of moles

of each element in the compound.

The molecular formula is also known as the __________________________.

• It represents the _______________________________ ratio of moles of each element in the compound.

EXAMPLE: What is the empirical formula of dimethylhydrazine, C2H8N2, a colorless liquid used as a rocket fuel?

EXAMPLE: Elemental analysis of a sample of an ionic compound showed 2.82 g of Na, 4.35 g of Cl, and 7.83 g of O. What

is the empirical formula and name of the compound?

EXAMPLE: After a workout session, lactic acid (M = 90.08 g/mol) forms in muscle tissue and is responsible for muscle

soreness. Elemental analysis shows that this compound contains 40% C, 6.7% H and 53.3% O. Determine the molecular

formula.

C6H12O6 =




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CONCEPT: COMBUSTION ANALYSIS

Under a combustion reaction a compound made of _______________ or _______________ reacts with ______________ .

• The products formed will be ________________ and ________________ .

EXAMPLE: A 0.2500 g sample contains carbon, hydrogen and oxygen and undergoes complete combustion to produce

0.3664 g of CO2 and 0.1500 g of H2O. What is the empirical formula of the compound?




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CONCEPT: CHEMICAL COMPOSITION

The mass per mole of an element is called its _____________________________ (M).

The mass per mole of a compound is called its ____________________________ (M).

• They both have the units of _______________.

1. Elements. To find mass of an element just look up its atomic mass in the periodic table.

EXAMPLE: What is the total mass of each of the following elements?

a. Sodium b. Gold c. Mercury 2. Compounds. The mass of a compound is the sum of the individual masses of the elements in the chemical formula.

EXAMPLE: What is the total mass of each of the following compounds?

a. N2O5 b. C12H22O11 c. (NH4)3PO4




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CONCEPT: MASS PERCENT

Mass or weight percent is the percentage of a given element in a compound.

EXAMPLE: What is the percentage of carbon in sodium hydrogen carbonate, NaHCO3?

EXAMPLE: A sample of toothpaste contains tin (II), SnF2. Analysis of a 5.25 g sample contains 8.77 x 10-3 g of F. What is

the percentage of tin (II) fluoride in the sample?

PRACTICE: Hemoglobin contains 0.33% iron and has a molecular weight of 68 kg. How many iron atoms are in each

molecule of hemoglobin?

Mass Percent (%) =

MassComponent)(TotalMass)(

•100




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CONCEPT: BALANCING CHEMICAL REACTIONS

When balancing an equation always make sure the ________ and ________ of atoms on both sides of the arrow are equal.

EXAMPLE: Write balanced equations for each of the following by inserting the correct coefficients in the blanks:

a. ____ Al (s) + ____ Cl2 (g) ____ AlCl3 (s)

b. ____ Ba3(PO4)2 (s) + ____ KOH (aq) ____ K3PO4 (aq) + ____ Ba(OH)2 (aq)

c. ____ C4H10 (aq) + ____ O2 (g) ____ CO2 (g) + ____ H2O (l)




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CONCEPT: GROUP REACTIVITY

The central principle of Organic Chemistry is based on the _________________________________. • The reactivity of an organic compound is all based on which type is present.

C

CC

C

CC

CO

C

C

CC

CC

CH3

H3C

H3C

OH

Tetrahydrocannabinol (THC)

CH2CH2CH2CH2CH3

Alkane

Alkene

Alkyne

Alcohol

Amine

Aldehyde

Ketone

Carboxylic Acid

Ester




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CONCEPT: GROUP REACTIVITY (PRACTICE)

EXAMPLE: In each of the following molecules, identify the type(s) of functional groups present.

a. b.

O

OH

PRACTICE: In each of the following molecules, identify the type(s) of functional groups present.

a. b.

O

O




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CONCEPT: ATOMIC PROPERTIES AND CHEMICAL BONDS

Before we examine the types of chemical bonding, we should ask why atoms bond at all.

• Generally, the reason is that ionic bonding ____________ the potential energy between positive and negative ions.

• Generally, the reason covalent bonds form is to follow the ____________ rule, in which the element is then

surrounded by 8 valence electrons.

There are three models of chemical bonding:

In ____________________ bonding, metals connect to non-metals.

• __________________ transfers an electron to the ________________ , creating ions with opposite charges that

are attracted to each other.

Li F Li F Li F

In _______________ bonding, non-metals connect to non-metals.

• In it the nonmetals __________________ electron pairs between their nuclei.

ClCl

In _______________ bonding, metal atoms “pool” their valence electrons to form an electron “sea” that holds the metal-ion

together




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CONCEPT: CHEMICAL BONDS (PRACTICE)

EXAMPLE: Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound.

atomic element ––

molecular element ––

molecular compound ––

ionic compound ––

a. Iodine

b. NH3

c. Graphite

d. Na3P

e. Ag2(SO4)2





CONCEPT: THE IONIC-BONDING MODEL

The central idea of ionic bonding is that the metal transfers an electron(s) to a nonmetal.

• The metal then becomes a(n) ____________ (positive ion). and the nonmetal becomes a(n) _____________ (negative ion).

• Their opposite charges cause them to combine into a crystalline solid.

PRACTICE: Determine the molecular formula of the compound formed from each of the following ions.

a. K+ & P3-

b. Sn4+ & O2-

c. Al3+ & CO32-





CONCEPT: COMMON POLYATOMIC IONS

Polyatomic ions are compounds made up of different elements, usually only ____________, and possess a ____________.

Singly Charged Cation (Positive Ion)

NH4+

Ammonium

Doubly Charged Cation (Positive Ion)

Hg22+

Mercury (I)

Singly Charged Anions (Negative Ions)

CH3CO2– or C2H3O2

–

Acetate

CN–

Cyanide

OH–

Hydroxide

MnO4–

Permanganate

NO3–

Nitrate

Nitrite

Doubly & Singly Charged Anions (Negative Ions)

HPO42–

Hydrogen Phosphate

H2PO4–

Dihydrogen Phosphate

HCO3–

Hydrogen Carbonate or Bicarbonate

HSO4–

Hydrogen Sulfate or Bisulfate

Doubly Charged Anions (Negative Ions)

CO32–

Carbonate

CrO42–

Chromate

Cr2O72–

Dichromate

O22–

Peroxide

SO42–

Sulfate

Sulfite

Triply Charged Anions (Negative Ions)

PO43–

Phosphate

Phosphite




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CONCEPT: POLYATOMIC IONS w/ HALOGENS

Polyatomic ions containing halogens are sometimes referred to as __________ halogens or halogen ________________.

These compounds share 4 common characteristics:

1.

2.

3.

4.

These compounds use the same system for naming:

PRACTICE: Name each of the following compounds.

a. BrO4 – b. FO2 –

c. ClO – d. IO3 –




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CONCEPT: IONIC COMPOUNDS

In the early days of chemistry, newly discovered compounds were given fancy names such as morphine, quicklime and

muriatic acid. Since then thousands of new compounds have been discovered and named under a system called

_____________________________.

Metals tend to __________ electrons to become positively charged ions called _______________.

Nonmetals tend to __________ electrons to become negatively charged ions called _______________.




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CONCEPT: NAMING BINARY IONIC COMPOUNDS Features: ___________________ & ___________________

Rules for Naming: a. The metal is named and written first.

• If the metal is a transition metal we must use a _________________________ to describe its positive charge.

b. The nonmetal keeps its base name but has its ending changed to ___________________.

EXAMPLE: Provide the molecular formula or name for each of the following compounds.

a. Calcium phosphide b. CoO

PRACTICE: Provide the molecular formula or name for each of the following compounds.

a. AlBr3 b. Lead (IV) sulfide c. SnO2





CONCEPT: NAMING IONIC COMPOUNDS w/ POLYATOMICS Features: _________________ & _________________

Rules for Naming:

a) The metal keeps its name and is named and written first.

• If the metal is a transition metal we must use a _____________________ to describe its positive charge.

b) Name the polyatomic as you would normally.

EXAMPLE: Write the formula for each of the following compounds:

a. Iron (III) Acetate b. Copper (I) phosphate

c. Strontium Carbonate d. Ammonium Nitrite

EXAMPLE: Give the systematic name for each of the following compounds:

a. Pb(CrO4)2 b. Ga(ClO4)3

c. Mn(HSO4)2 d. Ba(CN)2





CONCEPT: NAMING IONIC HYDRATES Features: _________________ & _________________

CuSO4 5 H2O

Rules for Naming the Ionic Compound portion: a. The metal is named normally and written first.

• If the metal is a transition metal we must use a ________________________ to describe its positive charge.

b. The nonmetal keeps the first part of its name but has its ending changed to ___________________.

c. Name the polyatomic as you would normally.

Rules for Naming the H2O portion:

a. The H2O portion will be called ___________________ .

b. To describe the number of H2O molecules use these prefixes.

EXAMPLE: Write the formula for each of the following compounds.

a. Calcium carbonate hexahydrate

b. Lead (IV) Sulfate pentahydrate

PRACTICE: Give the systematic name for each of the following compounds:

a) K2Cr2O7 · 3 H2O b) Sn(SO3)2 · 4 H2O





CONCEPT: NAMING MOLECULAR COMPOUNDS Features: _________________ & _________________

Because molecular compounds combine in different proportions to form different compounds, we must use numerical

prefixes.

Rules for Naming: a. The first nonmetal is named normally and uses all numerical prefixes except ___________________. b. The second nonmetal keeps its base name but has its ending changed to _____________________. EXAMPLE: Write the formula for each of the following compounds.

a. Disulfur monobromide b. Iodine Tetrachloride


a. CO b. N2S4 c. IO5





CONCEPT: NAMING ACIDS

1. BINARY ACIDS Features: _______________________ + _______________________

Rules for Naming:

a. The prefix will be ___________________ .

b. Use the base name of the nonmetal.

c. The suffix will be ___________________ .

EXAMPLE: Write the formula for each of the following compounds:

a. Hydroiodic acid b. Hydroselenic acid c. Hydrofluoric acid


a. HBr b. H2S c. HCN

2. OXOACIDS or OXYACIDS Features: _______________________ + _______________________

Rules for Naming: a. If the polyatomic ion ends with –ate then change the ending to _____________________. b. If the polyatomic ion ends with –ite then change the ending to ______________________.

EXAMPLE: Give the systematic name or formula for each of the following compounds:

a. H2CO3 b. Nitric acid c. H2SO4

PRACTICE: Give the systematic name or formula for each of the following compounds:

a. Hypobromous acid b. HClO3 c. Acetic acid





CONCEPT: ENERGY CONSIDERATIONS IN IONIC BONDING

________________________ is the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions. It

tells us the strength of ionic interactions and has an influence in melting point, hardness, solubility and other properties.

Li+ (g) + F – (g) LiF (s) ∆H = 1050 kJ/mol

In order to calculate the energy of an ionic bond we use the following equation;

Ionic Bond Energy =

Radius = __________________________________

EXAMPLE: For each pair, choose the compound with the lower lattice energy.

a. BaO or MgO

b. LiCl or CaS

PRACTICE 1: Choose the compound with the lower lattice energy.

a. AlN or KBr

PRACTICE 2: Choose the compound with the higher lattice energy.

a. CsF or LiCl




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CONCEPT: DIPOLE ARROWS

Before drawing covalent compounds we first need to understand the idea of polarity and its connection to electronegativity.

• Polarity arises whenever two elements are connected to each other and there is a significant difference in their

electronegativities.

• Generally, electronegativity ________________ going from left to right of a period and ________________ going

down a group.

To show this difference in electronegativity we use a dipole arrow.

The dipole arrow points towards the ________________ electronegative element.

The Effect of Electronegativity Difference on Bond Classification

Electronegativity Difference (ΔEN)

Bond Classification

Example

Zero (0.0)

Pure Covalent

Small (0.1 – 0.4)

Nonpolar Covalent

Intermediate (0.4 – 1.7)

Polar Covalent

Large (Greater than 1.7)

Ionic




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PRACTICE: DIPOLE ARROWS

EXAMPLE: Based on each of the given bonds determine the direction of the dipole arrow and the polarity that may arise.

a. H Cl

b. S O

c. Br B Br

PRACTICE 1: Based on the given bond determine the direction of the dipole arrow and the polarity that may arise.

a. H C


a. N F


a. H N H





CONCEPT: CHEMICAL BOND IDENTIFICATION

PRACTICE: Answer each of the following questions dealing with the following compounds.

KBr NH3 F2 CaO NaClO

a. Which of the following compound(s) contains a polar covalent bond?

b. Which of the following compound(s) contains a pure covalent bond?

c. Which of the following compound(s) contains a polar ionic bond?

d. Which of the following compound(s) contains both a polar ionic bond and a polar covalent bond?





CONCEPT: ELECTRON-DOT SYMBOLS

Before we look at the first two bonding models, we have to figure out how to depict the valence electrons of bonding atoms.

• In the _________ electron-dot symbol, the element symbol represents the nucleus and inner electrons, and the

surrounding dots represent the ________________ electrons.

EXAMPLE: Draw the electron-dot symbol for each of the following elements.

1A 2A 3A 4A 5A 6A 7A 8A

Li

Be

B

C

N

O

F

Ne

It’s easy to write the Lewis symbol for any Main-Group element:

1) Remember that Group Number equals Valence Electron Number.

2) Place one dot at a time on the four sides (top, right, bottom, left) of the element symbol.

3) Keep adding dots, pairing them up until you have reach the number of total valence electrons for that element.

PRACTICE 1: Draw the electron-dot symbol for the following ion.

Mg2+


N3-


Cr1+




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CONCEPT: CHEMICAL BONDING I

Rules for Drawing

1. Least electronegative element goes into the center. Important Facts to Know:

(a) Electronegativity increases across any Period going from left to right and up any Group going from bottom to top.

(b) Hydrogen and Fluorine ________________ go in the center and they only make _________ BOND.

2. Number of valence electrons equals group number.

3. Carbon must make _____ bonds, except in rare occasions when it makes _____ bonds.

• If the carbon atom were positive or negative then it would make _____ bonds

4. Nitrogen likes to make _____ bonds.

5. Oxygen likes to make _____ bonds.

6. Halogens (Group 7A), when not in the center, make _____ bond.

7. Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence

electrons around them when in the center.





CONCEPT: INCOMPLETE OCTETS

Nonmetals form covalent bonds to generally follow the ___________ rule, in which the element is surrounded by 8 valence

electrons.

• Sometimes elements form compounds in which they have ____________________ 8 valence electrons.

• These elements are said to have an incomplete octet or to be ________________________________________ .

EXAMPLE: Draw the following molecular compound.

BH3

PRACTICE: Draw the following molecular compound.

BeCl2





CONCEPT: EXPANDED OCTETS

Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence


EXAMPLE: Draw each of the following molecular compounds.

IF3 KrF5+

PRACTICE 1: Draw the following molecular compound.

SBr4


I3–





CONCEPT: EXPANDED OCTETS

Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence



IF3 KrF5+


SBr4


I3–





CONCEPT: POLYATOMIC IONS

Shortcut: If you have _____, _____, _____, _____, __________________ or __________________ connected to oxygen

then the negative charge tells us how many oxygens are single bonded.

• The remaining oxygens are _______________________ bonded to the central element.


SO42-

PO43- H2SO4


SeO42-


XeO64-





CONCEPT: FORMAL CHARGE

Structures and polyatomic ions that break the octet rule often have ________________ Lewis Structures.

• The purpose of using the formal charge formula is to determine which Lewis structure is the best answer.

Formal Charge =

a) Use formal charge formula to check to see if you drew your compound correctly.

b) Formal charges must be either _____, ______, ______.

c) If you add up all the formula charges in your compound that will equal the overall charge of the compound.

EXAMPLE: Calculate the formal charge for each of the following element designated for each of the following.

a. The carbon atom in

b. The sulfur atom in

PRACTICE: Calculate the formal charge for each of the following element designated in the following compound.

a. Both oxygen atoms in:

!A B




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CONCEPT: RESONANCE STRUCTURES

Resonance structures are used to represent bonding in a molecule or ion when a single Lewis structure cannot correctly

describe the Lewis structure.

EXAMPLE: Determine all the possible Lewis structures possible for NO2–. Determine its resonance hybrid.

EXAMPLE: Determine the remaining resonance structures possible for the following compound, CO32-.

O

C OO




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CONCEPT: ELECTRONIC GEOMETRY

When drawing a compound you have to take into account two different systems of geometrical shape.

• The simpler system known as electronic geometry or __________ shape treats lone pairs (nonbonding electrons)

and surrounding elements as the same.

Key: A = Central Element

X = Lone Pairs and Surrounding Elements

O C O H C N

AX2 = Linear (2 Groups)

AX3 = Trigonal Planar (3 Groups)

F

BF F

SnF F

AX4 = Tetrahedral (4 Groups)

NH

HH C

Cl

ClCl

Cl

Cl

PCl

ClCl

Cl

AX5 = Trigonal Bipyramidal (5 Groups)

Xe FF

Cl

SCl Cl

Cl Cl

AX6 = Octahedral (6 Groups)

Cl

XeH

H

H

H

!

AX

XX

!


CH.3 - MOLECULAR SHAPE & STRUCTURE


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PRACTICE: ELECTRONIC GEOMETRY

EXAMPLE: Draw each of the following compounds and determine their electronic geometries.

PH3 BeCl2

PRACTICE 1: Draw the following compound and determine its electronic geometry.

SBr4


IF3


H2S


PO43-





CONCEPT: MOLECULAR GEOMETRY

When drawing a compound you have to take into account two different systems of geometrical shape.

• With the molecular geometry you treat lone pairs (nonbonding electrons) and

surrounding elements as different.

Key: A = Central Element X = Surrounding Element E = Lone Pair

O C O H C N

3 Groups2 Groups F

BF F

SnF F

AX3 - Trigonal Planar AX2E1 - Bent, Angular or V-ShapedAX2 - Linear

4 Groups NH

HHC

Cl

ClClCl H

OH

AX4 - Tetrahedral AX2E2 - Bent, Angular or V-ShapedAX3E1 - Trigonal Pyramidal

5 Groups

Xe FF

AX5 - Trigonal Bipyramidal

ClP

ClCl

Cl Cl

F Cl F

FAX4E1 - Seesaw AX2E3 - Linear

FS FF

F

AX3E2 - T-Shaped

6 Groups ClSCl Cl

Cl ClClXe

H

H

H

HAX6 - Octahedral AX4E2 - Square Planar

F

SFF F

F

AX5E1 - Square Pyramidal

AX

XX





PRACTICE: MOLECULAR GEOMETRY

EXAMPLE: Draw each of the following compounds and determine their molecular geometries.

PH2 – XeCl2

PRACTICE 1: Draw the following compound and determine its molecular geometry.

OBr2

PRACTICE 2: Draw the following compound and determine its molecular geometry.

SO42-





CONCEPT: IDEALIZED BOND ANGLES

According to the ___________________________________________ (VSEPR) model bond and lone electron pairs will

position themselves around the central element so that they are as far apart as possible.




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PRACTICE: IDEALIZED BOND ANGLES

EXAMPLE: Determine the bond angles of each of the following compounds.

CO2 BrF4+

PRACTICE 1: Determine the bond angle of the following compound.

AsCl5

PRACTICE 2: Determine the bond angle of the following compound.

IF3




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CONCEPT: HYBRIDIZATION

Covalent bonds are formed when atomic orbitals on different atoms overlap and electrons are shared.

Cl Cl

H H H H

ClCl

But what happens when we need to form covalent bonds with different atomic oribitals, for example BeCl2?

Be Cl

[He]2s2 [Ne]3s23p5

2

2p 2s 2p2s

Promotion

& Hybridization2p sp 2p2s

Be




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PRACTICE: HYBRIDIZATION

EXAMPLE: For each of the given covalent compounds draw out the Lewis Structure and answer the questions

CH2Cl2 Hybridization: XeCl5+ Hybridization:

Unhybridized Orbitals: Unhybridized Orbitals:

Bonding orbitals (C – Cl): Bonding orbitals (Xe – Cl):

PRACTICE 1: For the given covalent compound draw out the Lewis Structure and answer the questions.

IF5 Hybridization:

Unhybridized Orbitals:

Bonding orbitals (I – F):


NH3 Hybridization: Unhybridized Orbitals:

Bonding orbitals (N – H):





PRACTICE: HYBRIDIZATION

EXAMPLE: For each of the given covalent compounds draw out the Lewis Structure and answer the questions

CH2Cl2 Hybridization: XeCl5+ Hybridization:

Unhybridized Orbitals: Unhybridized Orbitals:

Bonding orbitals (C – Cl): Bonding orbitals (Xe – Cl):


IF5 Hybridization:

Unhybridized Orbitals:

Bonding orbitals (I – F):


NH3 Hybridization: Unhybridized Orbitals:

Bonding orbitals (N – H):





CONCEPT: MO THEORY In the molecular orbital theory electrons are seen as being _______________, or spread out over a molecule instead of concentrated in a covalent bond.

• A(n) _______________ orbital is the region of high electron density between nuclei where a bond forms.

• A(n) _______________ orbital is the region that has zero electron density (a node) between the nuclei where a

bond can’t form.

EXAMPLE: Use a MO diagram to write the electron configuration of each of the following:

a. C22-

b. F2+




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PRACTICE: MO THEORY

The MO diagram can be connected to the MO bond order:

Bond Order = 12

(# of e – in bonding MO – # of e – in anti-bonding MO)

A bond order __________________________ zero means that the compound is stable and exists.

A bond order __________________________ zero means the compound is unstable and does not exist.

• In general, the _______________ the bond order, the _______________ the bond.

PRACTICE: Use a MO diagram to determine if the following compound exists or not.

a. O22-

b. B2-




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CONCEPT: UNITS OF PRESSURE

Pressure is defined as the force exerted per unit of surface area.

The SI unit for Pressure is the ________________, which has the units of ________________.

The SI Unit for Force is the ________________, which has the units of ________________.

Despite a gradual change to SI units for pressure, many chemists still express pressure in _______ , _______ and _______

Unit Name

Pressure Unit

Pascal (Pa); Kilopascal (kPa)

1.01325 x 105 Pa; 101.325 kPa

Atmosphere (atm) 1 atm

Millimeters of mercury (mmHg) 760 mmHg

Torr 760 torr

Bar 1.01325 bar

Pounds per square inch (lb/in2 or psi) 14.7 lb/in2

Psi 14.696 psi

ForceArea Pressure =


CH.4 - THE PROPERTIES OF GASES


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PRACTICE: UNITS OF PRESSURE

EXAMPLE: A geochemist heats a limestone (CaCO3) sample and collects the CO2 released in an evacuated flask. The CO2

pressure is 283.7 mmHg. Calculate the CO2 pressure in torrs and atmospheres.

PRACTICE: If the barometer in a laboratory reads 34.2 inHg what is the pressure in bars? (1 in = 2.54 cm)




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CONCEPT: PARTIAL PRESSURES OF GASES

_________________ Law states that in a container of unreacting gases, the total pressure of the container is the sum of the

partial pressures of the individual gases.

PTOTAL = PGas 1 + PGas 2 + PGas 3 + ……

The total pressure is due to the total number of moles.

• The partial pressure of each gas molecule is the total pressure multiplied by the mole fraction of each gas molecule.

PGas1 = XGas 1)( • PTotal( )

XGas1 =molesGas1)(TotalMoles)(

X = ___________________________ PGas1 = ___________________________

EXAMPLE: A container has 16.7 g O2, 8.1 g H2 and 35.2 g N2 and contains a total pressure of 0.83 atm. Calculate the mole

fraction of O2 and its partial pressure.

PRACTICE: A gas mixture with a total pressure of 812 mmHg contains the following gases at with their partial pressures:

Cl2 = 210 mmHg, H2 = 180 mmHg, CO2 = 215 mmHg. If argon gas is also present calculate its mole fraction.




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27. A mixture of gases contains 7.5 g H2, 29.6 g N2 and 27.1 g O2 with a total pressure of 127 torr. Calculate the mole fraction

and partial pressure of N2 in atm.




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CONCEPT: THE KINETIC-MOLECULAR THEORY

To better understand the behavior of a gas we use the kinetic-molecular theory.

• Under this theory, a gas is seen as a collection of molecules or individual atoms that are in constant motion.

An ideal gas has the following characteristics:

1. The size of the particle is ____________________________________ when compared to the volume of the container.

2. The average kinetic energy of a particle is directly proportional to the temperature of the container in ______________ .

3. The collision of a particle with another gas particle or with the walls of a container are completely _________________ .

EXAMPLE: Two identical 10.0 L flasks each containing equal masses of O2 and N2 gas are heated to the same

temperature. Which of the following statements is/are true?

a) The flask with the oxygen gas will have a greater overall pressure.

b) The nitrogen and oxygen gases will have the same average speed or velocity

c) The nitrogen and oxygen gases will have the same average kinetic energy.




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30. Which of the following statements would correctly explain the non-ideal behavior of a gas based on the Kinetic Molecular Theory (KMT)?

i. At high temperatures the attractive forces between molecules becomes negligible.

ii. At high pressure the volume of gas molecules become significant.

iii. An increase or decrease in the moles of gas causes the gas constant value to change.





31. Which statement is TRUE about kinetic molecular theory? a) A single particle does not move in a straight line.

b) The size of the particle is large compared to the volume.

c) The collisions of particles with one another is completely elastic.

d) The average kinetic energy of a particle is not proportional to the temperature.





2. Answer each of the following questions:

For image A, the gas container is connected to an open-end U-tube manometer. The mercury in the manometer is 5.0 cm higher on the side open to the atmosphere. If the atmospheric pressure is 759 mmHg, what is the pressure of the gas in atm?






For image B, the gas container is connected to an open-end U-tube manometer. The mercury in the manometer is 7.6 cm lower on the side open to the atmosphere. If the atmospheric pressure is 1080 mmHg, what is the pressure of the gas in atm?






For image C, the gas container is connected to a closed end U-tube manometer. If the pressures of the gas and the atmosphere are initially 800 mmHg and 1200 mmHg respectively, what will be the pressure of the gas if mercury in the manometer is 5.0 cm higher on the side closer to the atmosphere?





7. A sealed container with a moveable piston contains a gas with a pressure of 1380 torr, a volume of 820 mL and a temperature of 31oC. What would the volume be if the new pressure is now 2.83 atm, while the temperature decreased to 25oC?





8. The pressure in a system is said to be 5.83 atm. What would be the new pressure if the number of moles of gas were quadrupled and the volume were tripled while maintaining constant temperature?





9. The pressure in a system is said to be 6.11 atm. What would be the new pressure if the number of moles of gas were cut by a third and the volume was cut by a fourth while maintaining constant temperature?





10. A bicycle tire is filled with air to a pressure of 4.25 atm at a temperature of 19oC. Riding the bike on a hot day increases the temperature of the tire to 52oC. The volume of the tire also increases by 5.0%. What is the new pressure in the bicycle tire?





12. What is the volume, in mL, occupied by 132.7 g CO2 (MW: 44.01 g/mol) at STP?





13. The volume of O2 gas collected at 24oC and an atmospheric pressure of 702 mmHg is 192 mL. Calculate the mass of the dry oxygen gas collected if the pressure of water vapor at 24oC is 22 mmHg.





16. What is the density (in g/L) of phosphorus pentachloride at 1157.3 mmHg and 32oC?





17. A gaseous compound of nitrogen and hydrogen is found to have a density of 0.977 g/L at 528 torr and 100oC. What is the molecular formula of the compound? a) N2H4 b) NH3 c) HN3 d) HN e) N4H8





18. Consider two containers of gases at the same temperature. One has helium at a pressure of 1.00 atm. The other contains carbon dioxide with the same density as the helium gas. What is the pressure of the carbon dioxide gas sample? a) 0.023 atm b) 1.00 atm c) 0.091 atm d) 0.18 atm e) 2.12 atm





19. Determine the molecular formula of a gaseous compound that is 49.48% carbon, 5.19% hydrogen, 28.85% nitrogen, and 16.48% oxygen. At 27oC, the density of the gas is 1.45 g/L and it exerts a pressure of 0.092 atm.





22. Calculate the molar mass, in gmol

, of a gaseous compound with a velocity of 312 ms

at 35oC.





25. How many times faster will H2 gas pass through a pin hole into an area of vacuum than O2 gas? a) 32 b) 2 c) 2.5 d) 4 e) 8





26. Rank the following in order of increasing rate of effusion:

O2 AlF5 CO2 Xe





29. Three identical flasks contain equal moles of three different gases all at standard temperature and pressure. Flask A contains C2H4, Flask B contains CO2 and Flask C contains Cl2. Answer each of the following questions:

a) Which flask will have the greatest overall pressure?

b) Which flask has the greatest average speed of velocity?

c) Which flask has the greatest average kinetic energy? a) Which flask has the greatest density?

b) Which flask has the most molecules?

c) Which flask contains the most number of atoms?

d) Which flask has the greatest momentum?





32. Which of the following statements is TRUE? a) Particles of different masses have the same average speed at a given temperature.

b) The larger a molecule, the faster it will effuse.

c) At very high pressures, a gas will occupy a larger volume than predicted by the ideal gas law.

d) For a given gas, the lower the temperature, the faster it will effuse.

e) None of the above statements are true.





33. Which conditions of P, T and n make for the most ideal gas? a) High P, high T, high n

b) Low P, low T, low n

c) Low P, high T, low n

d) Low P, high T, high n





34. Two identical 10.0 L flasks each containing equal masses of O2 and N2 gas are heated to the same temperature. Which of the following statements is/are true?

a) The flask with the oxygen gas will have a greater overall pressure.

b) The nitrogen and oxygen gases will have the same average speed or velocity

c) The nitrogen and oxygen gases will have the same average kinetic energy.





CONCEPT: POLARITY

Molecule that have _______________ sharing of electrons contain a molecular polarity.

• In these molecules both _________________ and _________________ determine the molecular polarity.

POLARITY RULES TO BEING NON-POLAR:

1) If central element has NO lone pair(s): a. Central element must be connected to the ___________ elements.

b. Central element must be ______________ electronegative than the surrounding elements.

PRACTICE 1: Determine if carbon dioxide, CO2, is polar or nonpolar.

2) If central element has lone pair(s): a. Central element must be connected to the ___________ elements.


c. Use dipole arrows to point to the ________ electronegative element. These dipole arrows must cancel out.

d. Dipole arrows extend _____________________ lone pairs. These lone pair dipole arrows must cancel out.

PRACTICE 2: Determine if xenon tetrafluoride, XeF4, is polar or nonpolar.


CH.5 - LIQUIDS & SOLIDS


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CONCEPT: POLARITY

Molecule that have _______________ sharing of electrons contain a molecular polarity.

• In these molecules both _________________ and _________________ determine the molecular polarity.


1) If central element has NO lone pair(s): a. Central element must be connected to the ___________ elements.


PRACTICE 1: Determine if carbon dioxide, CO2, is polar or nonpolar.

2) If central element has lone pair(s): a. Central element must be connected to the ___________ elements.


c. Use dipole arrows to point to the ________ electronegative element. These dipole arrows must cancel out.

d. Dipole arrows extend _____________________ lone pairs. These lone pair dipole arrows must cancel out.

PRACTICE 2: Determine if xenon tetrafluoride, XeF4, is polar or nonpolar.





CONCEPT: POLARITY

Molecule that have unequal sharing of electrons contain a molecular polarity. In these molecules both _________________ and _________________ determine the molecular polarity.


1) If central element has NO lone pair(s): a. Central element must be connected to the same elements.

b. Central element must be less electronegative than the surrounding elements.

EXAMPLE: Determine if silicon tetrachloride , SiCl4, is polar or nonpolar.

2) If central element has lone pair(s): a. Central element must be connected to the same elements.

b. Central element must be less electronegative than the surrounding elements.

c. Use dipole arrows to point to the more electronegative element. These dipole arrows must cancel out.

d. Dipole arrows extend out from lone pairs. These lone pair dipole arrows must cancel out.

EXAMPLE: Determine if phosphorus trihydride, PH3, is polar or nonpolar.





PRACTICE: POLARITY PRACTICE 1: Determine if the following compound is polar or nonpolar.

a. SiBr42- PRACTICE 2: Determine if the following compound is polar or nonpolar.

a. H2S PRACTICE 3: Determine if the following compound is polar or nonpolar.

a. PCl2F3 PRACTICE 4: Determine if the following compound is polar or nonpolar.

a. IF2 –





CONCEPT: INTERMOLECULAR FORCES

When looking at a molecular substance such as H2O you will discover two types of electrostatic forces at work:

______________ forces exist within a molecule and influences the _____________ properties of the substance.

______________ forces exist between molecules and influence the _____________ properties of the substance.

___________________ is the force that exists between an ion and a polar compound. (Strongest)

Ex:

___________________ is the force that exists when H is directly connected F, O, N. (2nd Strongest)

Ex:

____________________ is the force that exists when two polar covalent compounds interact. (3rd Strongest)

Ex:

____________________is the force that exists when a nonpolar covalent compound interacts with a polar

covalent compound. (4th Strongest)

Ex:

____________________ is the force that exists when two nonpolar covalent compounds interact. (Weakest)

Ex:




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PRACTICE: INTERMOLECULAR FORCES

EXAMPLE: Based on the given compounds, answer each of the following questions:

a. CH3CH3 b. KBr c. C6H5OH d. CaS e. Ne

a) Which compound will have the lowest boiling point?

b) Which compound will have the highest surface tension.

c) Which compound will have the highest vapor pressure.

PRACTICE 1: The predominant intermolecular force in C6H5NH2 is:

a. London Dispersion b. Hydrogen Bonding c. Ion-Dipole d. Dipole-Dipole e. Dipole-induced Dipole

PRACTICE 2: The predominant intermolecular force in HBr is:


PRACTICE 3: The predominant intermolecular force in ZnBr2 with H2O is:


PRACTICE 4: The predominant intermolecular force in Ne with H2O is:






PRACTICE: INTERMOLECULAR FORCES

EXAMPLE: Based on the given compounds, answer each of the following questions:

a. CH3CH3 b. KBr c. C6H5OH d. CaS e. Ne

a) Which compound will have the lowest boiling point?

b) Which compound will have the highest surface tension.

c) Which compound will have the highest vapor pressure.

PRACTICE 1: The predominant intermolecular force in C6H5NH2 is:


PRACTICE 2: The predominant intermolecular force in HBr is:


PRACTICE 3: The predominant intermolecular force in ZnBr2 with H2O is:


PRACTICE 4: The predominant intermolecular force in Ne with H2O is:






CONCEPT: SOLUBILITY

According to the theory of __________ dissolves ____________ compounds with the same intermolecular

force or polarity will dissolve into each other.

EXAMPLE: Identify the intermolecular forces present in both the solute and the solvent, and predict whether a solution will

form between the two.

a. CCl4 and P4

b. CH3OH and C6H6

c. C6H5CH2NH2 and HF

d. IF4 – and NH3

PRACTICE: Which of the following statements is/are true?

a. Methane will dissolve completely in acetone, CH3COCH3.

b. Hydrofluoric acid (HF) will form a heterogeneous mixture with tetrachloride, CCl4.

c. Pentane will form a homogeneous mixture with CBr4.

d. Methanethiol (CH3SH) is miscible in fluoromethane (CH3F).





CONCEPT: PHASE DIAGRAMS

Under appropriate conditions of pressure and temperature, most substances can exist in 3 states of matter: ___________,

_______________ and ________________.

Microscopic Explanation for the Behavior of Gases, Liquids and Solids

Gas Liquid Solid

Assumes the

___________ and ___________

of its container.

Assumes the __________ of the

portion of its container it occupies,

but not the __________ .

Maintains a fixed

___________ and ___________

___________________ compressible

Viscosity Viscosity Viscosity

____________________ Viscous

___________________ compressible ___________________ compressible

____________________ Viscous ____________________ Viscous

Now, a convenient way to show the effect that temperature and pressure has on a pure substance in a closed system

without any air is to use a phase diagram.




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PRACTICE: PHASE DIAGRAMS

a) At what temperature can we no longer tell the difference between the liquid and the gas? _____________

b) Which point represents an equilibrium between the solid, liquid and gas phase? _____________

c) Which line segment represents an equilibrium between fusion and freezing? _____________

d) Which line segment represents an equilibrium between sublimation and deposition? _____________

e) Which line segment represents an equilibrium between condensation and vaporization? _____________

f) What is the normal freezing point of this unknown substance? _____________

g) What is the normal boiling point of this unknown substance? _____________




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CONCEPT: CLASIUS-CLAPEYRON EQUATION

By using the Clasius-Clapeyron equation a quantitative relationship between _____________________ and _____________________ can be established.

ln P2P1= −

ΔHvap

R1T2−1T1

#

$%

&

'(

`

EXAMPLE : The heat of vaporization (∆Hvap) of water is 40.3 kJ/mol at its normal boiling point at 100oC. What is the vapor pressure (in mmHg) of water at 60oC?




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CONCEPT: ENERGY CHANGES AND ENERGY CONSERVATION

_______________________ is the branch of physical science concerned with heat and its transformations to and from

other forms of energy.

_______________________ is the branch of chemistry that deals with the heat involved in chemical and physical changes.

Energy Changes and Energy Conservation

• The ______________________ is the specific part of the universe that we are focused on.

• The _________________________ deals with everything outside of it.

When talking about the movement of energy or heat between the ____________________ & ____________________ we

use the terms: Endothermic & Exothermic.

2 H2 (g) + O2 (g) 2 H2O (l) +

HEAT

HEAT

+ 2 HgO (s) 2 Hg(l) + O2 (g)

EXAMPLE: Classify each of the following process as either exothermic or endothermic: a) Fusion of Ice.

b) Sublimation of CO2.

c) Vaporization of aqueous water.

d) Deposition of chlorine gas.

e) Condensation of water vapor.


CH.7 - THERMODYNAMICS: THE FIRST LAW


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1. The First Law of Thermodynamics states that:

a. all product-favored processes happen quickly.

b. The total entropy in the universe is always increasing.

c. The entropy of pure crystalline substances at 0 K = 0.

d. The total energy in the universe is a constant.





CONCEPT: ENERGY CHANGES AND ENERGY CONSERVATION

_______________________ is the branch of physical science concerned with heat and its transformations to and from

other forms of energy.

_______________________ is the branch of chemistry that deals with the heat involved in chemical and physical changes.

Energy Changes and Energy Conservation

• The ______________________ is the specific part of the universe that we are focused on.

• The _________________________ deals with everything outside of it.

When talking about the movement of energy or heat between the ____________________ & ____________________ we

use the terms: Endothermic & Exothermic.

2 H2 (g) + O2 (g) 2 H2O (l) +

HEAT

HEAT

+ 2 HgO (s) 2 Hg(l) + O2 (g)

EXAMPLE: Classify each of the following process as either exothermic or endothermic: a) Fusion of Ice.




e) Condensation of water vapor.





CONCEPT: ENERGY FLOW TO AND FROM A SYSTEM

The _______ Law of Thermodynamics states that energy cannot be created nor destroyed, but only converted from one

form to another.

In chemistry, we are normally concerned with the energy changes associated with the system, not with its surroundings.

∆E = q + w q = ∆H (enthalpy) w = - P∆V

∆E =

q = * For q: (+) when system __________, __________, __________, heat or energy,

(-) when system __________, __________, __________, __________ heat or energy.

w = * For w : (+) when work done _____ system _____ the surroundings. Key word: volume ______________

(-) when work done _____ system _____ the surroundings. Key word: volume _______________

EXAMPLE: Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as

losing heat to the surroundings?

a) q = - , w = - b) q = +, w = + c) q = -, w = + d) q = +, w = -





3. Which process, among the listed chemical and physical changes, has ΔHo > 0?

a) Lake Michigan cools during the fall from 75 oC to 65 oC.

b) Magnesium burns in air to form magnesium oxide, MgO (s).

c) Liquid nitrogen boils at -196 oC, forming nitrogen gas.

d) Water condenses to a liquid. e) None of the above





PRACTICE: ENERGY FLOW TO AND FROM A SYSTEM

EXAMPLE: An unknown gas expands in a container increasing the volume from 4.3 L to 8.2 L at a constant pressure of 931

mmHg.

a. Calculate the work done (in kJ) by the gas as it expands. (1 L · atm = 101.3 J)

b. Using part A, calculate the internal energy of the system if the system absorbs 2.3 kJ of energy.

c. Using part B, calculate the internal energy of the system if the system does work against a vacuum.

PRACTICE: The reaction of nitrogen with hydrogen to make ammonia has an enthalpy, ∆H = - 92.2 kJ:

N2 (g) + 3 H2 (g) 2 NH3 (g)

What is in the internal energy of the system if the reaction is done at a constant pressure of 20.0 atm and the volume compresses from 10 L to 5 L?





5. A gas reaction is allowed to take place in a canister while submerged in water at a temperature of 25oC. The gas

expands and does P-V work on the surroundings equal to 385 J. At the same time, the temperature of the water

decreases to 20oC as the energy in the gas reaction reaches 364 J. What is the change in energy of the system?





6. When NH4NO3 dissolves in water the solution becomes colder. Based on this observation, circle the correct choices.

a) Endothermic or Exothermic

b) ΔHo > 0 or ΔHo < 0





CONCEPT: CONSTANT-VOLUME CALORIMETRY

Every object has its own _________________________ (C), the quantity of heat required to change its temperature by 1 K.

C = qΔT

[in units of JK ]

_________________________________ (c), the quantity of heat required to change 1 gram of a substance by 1 degree K.

c = [in units of J

g•K ]

If we know c of a substance, we can algebraically solve the amount of heat absorbed or released:

q =

EXAMPLE: At constant volume, the heat of combustion of a particular compound is – 4621.0 kJ/mol. When 2.319 grams of

this compound (molar mass = 192.75 g/mol) was burned in a bomb calorimeter, the temperature of the calorimeter

(including its contents) rose by 3.138oC. What is the heat capacity of the calorimeter in J/K?




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PRACTICE: CONSTANT-VOLUME CALORIMETRY

EXAMPLE: In an experiment a 9.87 carat (1 carat = 0.200g) diamond is heated to 72.25oC and immersed in 22.08 g of

water in a calorimeter. If the initial temperature of the water was 31.0oC what is the final temperature of the water? (cdiamond =

0.519J

g• oC) (cwater = 4.184

Jg• oC

).

PRACTICE 1: A sample of copper absorbs 35.3 kJ of heat, which increases the temperature by 25oC, determine the mass

(in kg) of the copper sample if the specific heat capacity of copper is 0.385 J

g• oC.

PRACTICE 2: 50.00 g of heated metal ore is placed into an insulated beaker containing 822.5 g of water. Once the metal

heats up the final temperature of the water is 32.08oC. If the metal gains 14.55 kJ of energy, what is the initial temperature

of the water?





19. A 7.55 g sample of aniline (C6H5NH2, molar mass = 93.13 g/mol) was combusted in a bomb calorimeter. If the

temperature rose by 30.0°C, use the information below to determine the heat capacity of the calorimeter. 4 C6H5NH2 (l) + 35 O2 (g) 24 CO2 (g) + 14 H2O (g) + 4 NO2 (g) ΔH°rxn= -1.28 x 104 kJ





CONCEPT: CONSTANT-PRESSURE CALORIMETRY The ______________ of a reaction can be calculated through the use of

a coffee-cup calorimeter.

EXAMPLE: You place 50.0 mL of 0.100 M NaOH into a coffee-cup calorimeter at 50.00oC and carefully add 75.0 mL of 0.100 M HCl,

also at 50.00oC. After stirring, the final temperature is 76.12oC. (Heat capacity and density of water: 4.184 J

g• oC and 1.00

gmL ).

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

a) Calculate qsoln (in J)

b) Calculate the enthalpy, ∆Hrxn (in J/mol), for the formation of water.





CONCEPT: THERMOCHEMISTRY AND STOICHIOMETRY A _______________________ equation is a balanced equation that includes the heat of reaction (∆Hrxn).

EXAMPLE: Iron reacts with oxygen according to the following equation at around 2000oC.

4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s) ∆Ho = + 65.2 kJ/mol

Calculate the amount of grams of iron (III) oxide, Fe2O3, produced from the absorption of 4.82 x 109 J.

PRACTICE: Nitromethane (CH3NO2), sometimes used as a fuel for drag racing, burns according to the following reaction:

4 CH3NO2 (l) + 7 O2 (g) 4 CO2 (g) + 6 H2O (g) + 4 NO2 (g) ∆Ho = – 2441.6 kJ

How much heat is released by burning 125.0 g of nitromethane (MW: 61.044 g/mol)?




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17. Compounds containing boron and hydrogen have been documented for their unusual bonding behaviors. In the following example, diborane (B2H6) forms trihalides even at low temperatures:

B2H6 (g) + 6 Cl2 (g) 2 BCl3 (g) + 6 HCl (g) ∆Hrxn = - 755.4 kJ How many grams of B2H6 (MW: 27.668 g/mol) reacting with excess Cl2 are needed to release -1388.55 kJ of energy?




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CONCEPT: HEAT SUMMATION

Many reactions are difficult, even impossible, to carry out in a single chemical step.

• They may often times require multiple steps to get to the final products.

__________ Law states that the enthalpy change (∆H) of an overall process is the sum of the enthalpy changes of its

individual steps.

EXAMPLE: For the following example calculate the unknown ∆H from the given ∆H values of the other equations.

Calculate the ∆Hrxn for

S(s) + 32

O2 (g) SO3 (g) ∆H = ?

Given the following set of reactions:

12

S (s) + 12

O2 (g) 12

SO2 (g) ∆H1 = – 296.8 kJ

2 SO3 (g) 2 SO2(g) + O2 (g) ∆H2 = 198.4 kJ




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PRACTICE: HEAT SUMMATION

A. Calculate the ∆Hrxn for

CO(g) + NO (g) CO2 (g) + 12

N2 (g) ∆H = ?

Given the following set of reactions:

CO2 (g) CO (g) + 12

O2 (g) ∆H1 = 283.0 kJ

N2 (g) + O2 (g) 2 NO (g) ∆H2 = 180.6 kJ

B. Calculate the ∆Hrxn for

ClF (g) + F2 (g) ClF3 (g) ∆Hrxn = ?

Given the following reactions:

Cl2O (g) + F2O 2 ClF (g) + O2 (g) ∆Hrxn = - 167.4 kJ

4 ClF3 (g) + 4 O2 (g) 2 Cl2O (g) + 6 F2O (g) ∆Hrxn = 682.8 kJ

2 F2 (g) + O2 (g) 2 F2O (g) ∆Hrxn = -181.7 kJ




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CONCEPT: STANDARD HEATS OF FORMATION (∆HRXN)

In a _______________ equation, 1 mole of a compound forms from its elements. The ______________________________

(∆Hof) is the enthalpy change for the chemical equation when all the substances are in their standard states.

C (graphite) + 2 H2 (g) CH4 (g) ΔHfo = −74.9kJ

When calculating ∆Hof remember:

1) An element in its standard state (elemental state) is given an ∆Hof of zero.

Ex: Na (s) P4 (s) Cl2 (g) S8 (g)

2) Most compounds have a negative ∆Hof. 3) To find the ∆Hrxn use the following formula:

ΔHrxno = ΔHf(products)

o −ΔHf(reac tan ts)o

EXAMPLE: The oxidation of ammonia is given by the following reaction:

4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g)

Calculate the ∆Horxn if the ΔHfo

value for NH3 , NO and H2O are – 45.9 kJ/mol, 90.3 kJ/mol and – 241.8 kJ/mol

respectively.

PRACTICE: Ibuprofen is used as an anti-inflammatory agent used to deal with pain and bring down fevers. If it has a

molecular formula of C13H18O2 determine the balanced chemical equation that would give you directly the enthalpy of

formation for ibuprofen.




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22. Choose the reaction that illustrates ΔH fo for Ca(NO3)2.

A) Ca (s) + N2 (g) + 3 O2 (g) Ca(NO3)2 (s)

B) Ca2+ (aq) + 2 NO3- (aq) Ca(NO3)2 (aq)

C) Ca (s) + 2 N (g) + 6 O (g) Ca(NO3)2 (s)

D) Ca(NO3)2 (aq) Ca2+ (aq) + 2 NO3- (aq)

E) Ca(NO3)2 (s) Ca (s) + N2 (g) + 3 O2 (g)





24. Identify a substance that is not in its standard state. a) Na b) P4 c) Cl e) O3 e) Ne





25. An organic compound has a molecular formula of C24H44O3. Answer each of the following questions. a) Write the balanced chemical equation corresponding to its enthalpy of formation.

b) Write the balanced chemical equation corresponding to its enthalpy of combustion.





PRACTICE: STANDARD HEATS OF FORMATION (∆HRXN)

EXAMPLE: Use the following bond strength values (kJ/mol):

C–H 412 C–O 360 C=O 743 C–C 348 H–H 436

C=C 611 C≡C 837 C≡O 1072 O–H 464 O=O 498

Calculate the enthalpy of the reaction shown in the formula below:

H–C≡C–H + H–H + C=O

C

C C

O

HH

H H





4. An unknown gas expands in a container increasing the volume from 8.7 L to 18.9 L at a constant pressure of 1380 mmHg. (a) Calculate the work done (in J) by the gas as it expands. (1 L· atm = 101.3 J). (b) Calculate the internal energy of the system if the system absorbs 235.5 J of energy.

(c) Calculate the internal energy of the system if work was done against a vacuum. (1 L · atm = 101.3 J).





8. Calculate the amount of heat absorbed when 12.0 g of water is heated from 20oC to 100oC. (c = 4.184 J/g· oC).





9. 101.3 g of an unknown metal has an initial temperature of 25oC. If it absorbs 639.1 J of energy to obtain a final

temperature of 32.01oC identify the unknown metal.

Metal Specific Heat Capacity (J/g·oC)

Au 0.129

Fe 0.444

Al 0.900

Hg 0.139





10. Which substance has the highest molar heat capacity?

a) Copper (specific heat Cu (s): 0.39J

g ⋅ oC )

b) Silver (specific heat Ag (s): 0.23J

g ⋅ oC )

c) Iron (specific heat Fe(s): 0.46J

g ⋅ oC )

d) Lead (specific heat Pb (s): 0.13J

g ⋅ oC )





11. 25.00 g of heated metal ore is placed into an insulated beaker containing 615.5 g of water at 42.18oC. If the metal

gains 19.11 kJ of energy, what is the final temperature of the water? (cwater = 4.184 J/g · oC).





12. If 53.2 g Al at 25.0 oC is placed in 110.0 g H2O at 90 oC, what is the final temperature of the mixture? The specific

heat capacities of water and aluminum are 4.184 J/g · oC and 0.897 J/g · oC, respectively.





13. A 20.0 g sample of iron (specific heat Fe (s) = 0.46 J

g ⋅ oC ) has an initial temperature of 30.2 oC. If 0.310 kJ are

applied to the iron sample, calculate its final temperature.





14. A sample of H2O (l) containing 2.50 moles has a final temperature of 45.0 oC. If the sample absorbs 3.00 kJ of

heat, what is the initial temperature of the H2O (l)? The specific heat of H2O (l) is 4.184 J

g ⋅ oC .





Determine the heat released when 80.0 g H2O (l) at 90 oC is cooled to ice at – 10.0 oC. Specific Heat of H2O (l) = 4.184

Jg ⋅ oC . Specific Heat of H2O (s) = 2.09

Jg ⋅ oC . Heat of Fusion of water = 333

Jg .





If 1050 g of aluminum metal with a specific heat capacity of 0.902 J

g ⋅ oC at – 20 oC is placed in liquid water at 0.00 oC,

how many grams of liquid water are frozen by the time that the aluminum metal has warmed to – 10 oC? Heat of Fusion of

water = 333 Jg .





CONCEPT: THERMOCHEMICAL PROCESSES

______________________________ is the branch of physical science concerned with heat and its transformations to and

from other forms of energy.

In terms of a chemical reaction, you will learn that depending on certain conditions they can occur or not:

• A reaction that requires no outside energy source is classified as a natural process and is

______________________ .

• A reaction that requires a continuous energy source to happen is classified as an unnatural process and is

______________________ .

EXAMPLE 1: Which of the following statements is not true?

a) The reverse of a spontaneous reaction is always non-spontaneous.

b) A spontaneous reaction always moves towards equilibrium.

c) A highly spontaneous reaction can occur at a fast or slow rate.

d) It is possible to create a non-spontaneous reaction.


a) The rusting of iron by oxygen is a non-spontaneous reaction.

b) The addition of a catalyst to a reaction increases spontaneity.

c) The movement of heat from a cold object to a hot object is a non-spontaneous reaction.

d) The diffusion of perfume molecules from one side of a room to the other is a non-spontaneous reaction.

e) None of the above.


CH.8 - THERMODYNAMICS: THE SECOND AND THIRD LAWS


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CONCEPT: THERMOCHEMICAL PROCESSES

______________________________ is the branch of physical science concerned with heat and its transformations to and

from other forms of energy.

In terms of a chemical reaction, you will learn that depending on certain conditions they can occur or not:

• A reaction that requires no outside energy source is classified as a natural process and is

______________________ .

• A reaction that requires a continuous energy source to happen is classified as an unnatural process and is

______________________ .

EXAMPLE 1: Which of the following statements is not true?

a) The reverse of a spontaneous reaction is always non-spontaneous.

b) A spontaneous reaction always moves towards equilibrium.

c) A highly spontaneous reaction can occur at a fast or slow rate.

d) It is possible to create a non-spontaneous reaction.


a) The rusting of iron by oxygen is a non-spontaneous reaction.

b) The addition of a catalyst to a reaction increases spontaneity.

c) The movement of heat from a cold object to a hot object is a non-spontaneous reaction.

d) The diffusion of perfume molecules from one side of a room to the other is a non-spontaneous reaction.

e) None of the above.




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CONCEPT: ENTROPY AND SPONTANEOUS REACTIONS

The _______ Law of Thermodynamics states that molecular systems tend to move spontaneously to a state of maximum

randomness or disorder.

This disorder is also called entropy and uses the variable __________.

In general, as we move from a solid liquid gas then entropy will ___________________ and its

sign will be ___________________.

Conversely, if we move from a gas liquid solid then entropy will ___________________ and its

sign will be ___________________.

EXAMPLE 1: Which should have the highest molar entropy at 25oC?

a) Ga (l)

b) Ga (s)

c) Ga (g)

d) All of them have the same molar entropy.

EXAMPLE 2: Which substance has greater molar entropy.

a) CH4 (g) or CCl4 (l)

b) Ne (g) or Xe (g)

c) CH3OH (l) or C6H5OH (l)




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PRACTICE: ENTROPY AND SPONTANEOUS REACTIONS (CALCULATIONS 1)

EXAMPLE 1: Arrange the following substances in the order of increasing entropy at 25oC.

XeF4 (s) HI (g) BaO (s) H2 (g) Hg (l) Br2 (l)

EXAMPLE 2: Containers A and B have two different gases that are allowed to enter Container C. Based on the image of

Container C what is the sign of entropy, ∆So.

B CA

PRACITCE: An ideal gas is allowed to expanded at constant temperature. What are the signs of ∆H, ∆S & ∆G.

BA






EXAMPLE 1: Arrange the following substances in the order of increasing entropy at 25oC.

XeF4 (s) HI (g) BaO (s) H2 (g) Hg (l) Br2 (l)

EXAMPLE 2: Containers A and B have two different gases that are allowed to enter Container C. Based on the image of

Container C what is the sign of entropy, ∆So.

B CA

PRACITCE: An ideal gas is allowed to expanded at constant temperature. What are the signs of ∆H, ∆S & ∆G.

BA






EXAMPLE: Consider the spontaneous fusion of ice at room temperature. For this process what are the signs for ∆H, ∆S,

and ∆G?

∆H ∆S ∆G

a) + + +

b) - + 0

c) - + -

d) + + -

e) - - -

PRACTICE: Consider the freezing of liquid water at 30oC. For this process what are the signs for ∆H, ∆S, and ∆G?

∆H ∆S ∆G

a) + - +

b) - + 0

c) - + -

d) - - +

e) - - -






PRACTICE 1: Predict the sign of ∆S in the system for each of the following processes:

a) Ag+ (aq) + Br – (aq) AgBr (s)

b) CI2 (g) 2 CI – (g)

c) CaCO3 (s) CaO (s) + CO2 (g)

d) Pb (s) at 50oC Pb (s) at 70oC

PRACTICE 2: For each of the following reactions state the signs of ∆H (enthalpy) and ∆S (entropy):

a) Fusion of Ice.




e) Condensation of a water vapor.





CONCEPT: CALCULATING ENTROPY OF A SYSTEM

The 2nd Law of Thermodynamics states in terms of a system the entropy of a system increases spontaneously.

• Besides the system we also have our ___________________ and together they form the total ________________.

Thus, to calculate the total entropy change, ∆STotal, we use the following equation:

ΔSTotal =

So if,

∆STotal > 0, the reaction is ______________________

∆STotal < 0, the reaction is ______________________

∆STotal = 0, the reaction is ______________________

EXAMPLE 1: Calculate the standard entropy (in kJ) of reaction at 25oC for the following reaction:

N2 (g) + 3 H2(g) 2 NH3 (g)

The standard molar entropies of N2, H2 and NH3 are 191.5 , 130.6 and 192.3 respectively.





PRACTICE: CALCULATING ENTROPY OF A SYSTEM (CALCULATIONS 1)

EXAMPLE: The oxidation of iron metal is given by the following reaction:

4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s)

a) Calculate the ∆Ssystem if the standard molar entropies of Fe, O2 and Fe2O3 are 27.3 , 205.0 and 87.4

respectively.

b) Calculate the ∆Ssurroundings if the reaction is spontaneous at 25oC. The standard molar enthalpy of Fe2O3 is – 824.2 .

c) Calculate the ∆STotal and determine if the reaction is spontaneous or non-spontaneous under standard-state conditions?

PRACTICE: Diethyl ether (C4H10O2, MW = 90.1 g/mol) has a boiling point of 35.6oC and heat of vaporization of 26.7 kJ/mol.

What is the change in entropy (in kJ/K) when 3.2 g of diethyl ether at 35.6oC vaporizes at its boiling point?






EXAMPLE: The oxidation of iron metal is given by the following reaction:

4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s)

a) Calculate the ∆Ssystem if the standard molar entropies of Fe, O2 and Fe2O3 are 27.3 , 205.0 and 87.4

respectively.

b) Calculate the ∆Ssurroundings if the reaction is spontaneous at 25oC. The standard molar enthalpy of Fe2O3 is – 824.2 .

c) Calculate the ∆STotal and determine if the reaction is spontaneous or non-spontaneous under standard-state conditions?

PRACTICE: Diethyl ether (C4H10O2, MW = 90.1 g/mol) has a boiling point of 35.6oC and heat of vaporization of 26.7 kJ/mol.

What is the change in entropy (in kJ/K) when 3.2 g of diethyl ether at 35.6oC vaporizes at its boiling point?





CONCEPT: GIBBS FREE ENERGY

Chemists are generally interested in the system (the reaction mixture) rather than the surroundings. In order to define the

free energy of a chemical system they use the following equations:

ΔGo = ΔHo −TΔSo ΔG = ΔGo +RTlnQ

If ∆G < 0, the reaction is _________________________

If ∆G > 0, the reaction is _________________________

If ∆G = 0, the reaction is _________________________

EXAMPLE 1: Which of the following statements is true for the following reaction?

N2O4 (g) 2 NO2 (g) ∆Ho = - 57.1 kJ ∆So = 175.8 kJ

a) The reaction is spontaneous at all temperatures.

b) The reaction is spontaneous at low temperatures.

c) The reaction is spontaneous at high temperatures.

d) The reaction is non-spontaneous at all temperatures.




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PRACTICE: GIBBS FREE ENERGY (CALCULATIONS 1)

EXAMPLE: The reduction of iron(III) oxide with hydrogen produces iron metal and can be written as follows:

Fe2O3 (s) + 3 H2 (g) 2 Fe (s) + 3 H2O (g) ∆Ho = 98.8 kJ ∆So = 141.5 Jk

a) Is this reaction spontaneous under standard-state conditions at 25oC? If not, at what temperature will it become

spontaneous?

PRACTICE 1: If ∆G is small and positive which of the following statements is true?

a) The forward reaction is spontaneous and the system is far from equilibrium.

b) The forward reaction is spontaneous and the system is near equilibrium.

c) The reverse reaction is spontaneous and the system is far from equilibrium.

d) The reverse reaction is spontaneous and the system is near equilibrium.

PRACTICE 2: Nitrogen gas combines with fluorine gas to form nitrogen trifluoride according to the reaction below at 25oC:

N2 (g) + 3 F2 (g) 2 NF3 (g) ∆Ho = -249.0 kJ ∆So = -278 J/K

Calculate ∆Go and state if the reaction favors reactants or products at standard conditions.

a) ∆Go = - 332 kJ; the reaction favors the formation of reactants.

b) ∆Go = - 166 kJ; the reaction favors the formation of products.

c) ∆Go = - 166 kJ; the reaction favors the formation of reactants.

d) ∆Go = - 332 kJ; the reaction favors the formation of products.









spontaneous?


















EXAMPLE: For mercury, ∆Hvap = 58.5 and ∆Svap = 92.9 at 25oC. Does mercury boil at 350oC and 1 atm

pressure?

EXAMPLE: The chemical reaction, 2 NO2Br (g) 2 NO2 (g) + Br2 (g), has a ΔSo = 135 Jmol ⋅k

and ΔHo=

926 kJmol

. Calculate the temperature when Keq = 4.50 x 105.






EXAMPLE 1: Calculate ∆Grxn at 25oC under the conditions shown below for the following reaction.

3 Cl2 (g) 2 Cl3 (g) ∆Go = + 31.6 kJ

The partial pressures of Cl2 and Cl3 are 0.83 atm and 4.9 atm respectively.

EXAMPLE 2: For the reaction: N2 (g) + 2 O2 (g) 2 NO2 (g), ΔGo= 75, 550 J

mol at 175 K and ΔGo

= 41,875

Jmol

at 225 K.

a) Calculate ΔSo and ΔHofor the reaction.






EXAMPLE 1: The normal boiling point of liquid propane is 231 K. What is the enthalpy of vaporization of liquid propane?

EXAMPLE 2: What is the entropy change associated with the expansion of one mole of an ideal gas from 2.5 L to 6.3 L at a

constant pressure of 1.25 atm?





CONCEPT: SOLUTIONS AND INTERMOLECULAR FORCES

Molarity (M) can serve as the connection between the interconversion of ____________ to ____________ and vice versa.

For example, a 5.8 M NaCl solution really means __________________________ per __________________________.

Molarity = Moles of SoluteLiters of Solution

A typical mixture consists of a smaller amount of one substance, the ________________, dissolved in a larger amount of

another substance, the __________________. Together they form a __________________.

According to the theory of “likes” dissolves “likes” compounds with the same ________________________

or ________________ will dissolve into each other.

EXAMPLE: Butane, a nonpolar organic compound, is most likely to dissolve in

a. HCl

b. C6H5OH

c. C8H18

d. AlCl3

e. What the heck is butane?


CH.9 - PHYSICAL EQUILIBRIA


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CONCEPT: EXPRESSING SOLUTION CONCENTRATION

Molarity represents the moles of solute dissolved per liters of solution and _______________, represents the moles of

solute dissolved per kg of solvent.

Molarity = Moles of SoluteLiters of Solution

=Moles of SoluteKg of Solvent

EXAMPLE: If the molality of glucose, C6H12O6, in an aqueous solution is 2.56 what is the molarity? Density of the solution is 1.530 g/mL.

PRACTICE: A solution is prepared by dissolving 43.0 g potassium chlorate, KClO3, in enough water to make 100.0 mL of

solution. If the density of the solution is 1.760 g/mL, what is the molality of KClO3 in the solution? (MW of KClO3 is 122.55 g/mol)





CONCEPT: EXPRESSING MASS AMOUNTS

Mass or weight percent is the percentage of a given element in a compound.

Mass%=Mass of SoluteMass of Solution

x100

=Moles of SoluteMoles of Solution

EXAMPLE: Commercial sulfurous acid, H2SO3, is 90.1% by weight, and its specific gravity is 1.51. Calculate the molarity of

commercial sulfurous acid. (MW of H2SO3 is 82.086 g/mol).

EXAMPLE: If the mole fraction of methanol, CH3OH, in an aqueous solution is 0.060 what is the molality? Density of the solution is 1.39 g/mL.





PRACTICE: EXPRESSING MASS AMOUNTS

PRACTICE 1: What is the weight percent of nitric acid in 3.26 m HNO3 (aq)? (MW of HNO3 is 63.018 g/mol).

PRACTICE 2: Calculate the mole fraction of acetic acid, HC2H3O2, in a 27.13 mass % aqueous solution (d = 0.9883 g/mL). (MW of HC2H3O2 is 60.054 g/mol).

PRACTICE 3: At 25.0oC, a solution is prepared by dissolving 12.7 g NaCl in 95.5 mL of water. What is the ppm of NaCl if

the density of water at this temperature is 0.9983 g/mL.





CONCEPT: SOLUBILITY AS AN EQUILIBRIUM PROCESS

When an ionic solid dissolves, ions leave the solid and become dispersed in the solvent.

In a(n) _________________ solution the maximum amount of dissolved solute is present in the solvent.

In a(n) _________________ solution additional amounts of solute can be further dissolved in the solvent.

In a(n) _________________ solution more than the equilibrium concentration of solute has been dissolved.

EXAMPLE C1: Caffeine is about 10 times as soluble in warm water as in cold water. A student puts a hot-water extract caffeine into an ice bath, and some caffeine crystallizes. What is the identity of the solution before it’s been placed in an ice bath?

a) Saturated

b) Super saturated

c) Unsaturated

d) Not enough information to answer the question.

________________ law explains the relationship between gas pressure and solubility: the solubility of a gas (Sgas) is

directly proportional to the partial pressure of the gas (Pgas) above the solution:

SGas = kH •PGas





PRACTICE: SOLUBILITY AS AN EQUILIBRIUM PROCESS

EXAMPLE: Henry’s Law Constant for nitrogen in water is 1.67 x 10-4 M � atm-1. If a closed canister contains 113 ppb

nitrogen, what would be its pressure in atm?

PRACTICE 1: In general, as the temperature increases, the solubility of gas in a given liquid ________________, and the

solubility of most solids in a given liquid ________________.

a. Increases, increases

b. Increases, decreases

c. Decreases, increases

d. Decreaes, decreases

PRACTICE 2: At a partial pressure of acetylene 1.35 atm, 1.21 moles of it dissolves in 1.05 L of acetonitrile. If the partial pressure of acetylene in acetone is increased to 12.0 atm, then what is its solubility?





CONCEPT: PROPERTIES OF SOLUTIONS

The 4 ______________________ properties help to explain what happens to a pure solvent as we add solute to it.

EXAMPLE: Explain what happens to each of the following properties as solute is added to a pure solvent.

a. Boiling Point ΔTb = i•kb •m

b. Freezing Point ΔTf = i•kf •m

c. Osmotic Pressure Π = i•MRT

d. Vapor Pressure PSolution = XSolvent •PSolvent

o

EXAMPLE: Which of the following compounds will have the highest boiling point? a) 0.10 m sucrose

b) 0.10 m CsBrO3

c) 0.35 m CH3OH

d) 0.15 m SrBr2




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PRACTICE: PROPERTIES OF SOLUTIONS

EXAMPLE 1: Pure water boils at 100oC. What is the expected boiling point of water after the addition of 13.12 g calcium

bromide, CaBr2, to 325 g water. Kb = 0.512 oC/m. (MW of CaBr2 is 199.88 g/mol).

EXAMPLE 2: The vapor pressure of water at 100.0oC is 0.630 atm. Determine the amount (in grams) of aluminum fluoride,

AlF3, (in grams) needed to reduce its vapor pressure to 0.550 atm. (MW of AlF3 is 83.98 g/mol).

PRACTICE 1: Beta-carotene is the most important of the A vitamins. Calculate the molar mass of Beta-carotene if 25.0 mL

of a solution containing 9.88 mg of Beta-carotene has an osmotic pressure of 56.16 mmHg at 30oC.





CONCEPT: THE LIQUID STATE

VAPOR PRESSURE is defined as the partial pressure of vapor molecules above the surface of the liquid under the dynamic equilibrium condition of condensation and evaporation.

EXAMPLE: The vapor pressure of pure liquid A is 550 torr and the vapor pressure of pure liquid B is 320 torr at room temperature. If the vapor pressure of a solution containing A and B is 465 torr, what is the mole fraction of A in the solution?





CONCEPT: THE LIQUID STATE (CALCULATIONS)

EXAMPLE: Determine the vapor pressure lowering associated with 1.32 m C6H12O6 solution (MW: 180.156 g/mol) at 25oC.

PRACTICE: The following boiling points belong to one of the following compounds: 117oC , 78oC, 34.5oC & 23oC

CH3-O-CH3 CH3CH2OH CH3CH2-O-CH2CH3 CH3CH2CH2CH2OH

a) Which boiling point goes with what compound?

b) If each of the following substances were placed in separate sealed clear bottles at room temperature, could you identify one of the substances right away?





CONCEPT: THE EQUILIBRIUM STATE

Most chemical reactions do not go to completion.

• ___________ do not completely convert into ___________ and reactant concentrations do not go down to ______.

• Instead, these reactions reach a state of chemical equilibrium, in which the reaction moves in the forward and

reverse direction.

These reactions are also called ___________________ reactions and are represented by using a double arrow.

PRACTICE: Which one of the following statements does not describe the equilibrium state?

a. While at equilibrium, a dynamic process is still occurring.

b. The concentration of the reactants is equal to the concentration of the products.

c. The concentration of the reactants and products reach a constant level.

d. At equilibrium, the net concentration of all species is not changing.

e. All are true.

AReaction : Bk1

k-1

0"

2"

4"

6"

8"

10"

12"

0" 1" 2" 3" 4" 5" 6" 7" 8"

Time (mins)

Mol

arity

B

A


CH.10 - CHEMICAL EQUILIBRIA


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CONCEPT: THE EQUILIBRIUM CONSTANT

The equilibrium constant, K, is a number equal to the ratio of ____________ to ____________ at a given temperature.

• Its magnitude tells us how far to the left or to the right our chemical equation lies at a particular temperature.

• If K is greater than 1 then ____________ are favored over ____________ and ____________ direction is favored.

• If K is less than 1 then ____________ are favored over ____________ and ____________ direction is favored.

K =

The equilibrium constant, K, takes into account all states of matter except: ________________ and _______________.

EXAMPLE: Write the equilibrium expression for the following reaction.

a) 2 N2O5 (aq) 4 NO2 (aq) + O2 (aq)

b) 2 PbO (s) + O2 (g) 2 PbO2 (s)

c) I2 (s) + 3 XeF2 (s) 2 IF3 (s) + 3 Xe (g)

PRACTICE: State which is greater in amount: reactants or products, based on the given equilibrium constant, K.

a) N2 (g) + O2 (g) 2 NO (g) K = 1.0 x 1020

b) 2 CO (g) + O2 2 CO2 (g) K = 2.2 x 10-22

c) 2 BrCl (g) Br2 (g) + Cl2 (g) K = 1





PRACTICE: THE EQUILIBRIUM CONSTANT (CALCULATIONS)

PRACTICE 1: The decomposition of nitrogen monoxide can be achieved under high temperatures to create the products of nitrogen and oxygen gas.

6 NO (aq) 3 N2 (aq) + 3 O2 (aq) a) What is the equilibrium equation for the reaction above? b) Write the equilibrium expression for the reverse reaction.

PRACTICE 2: The equilibrium constant, K, for the 2 NO (g) + O2 (g) 2 NO2 (g) is 6.9 x 102. What is the [NO] in an

equilibrium mixture of gaseous NO, O2 and NO2 at 500 K that contains 1.5 x 10-2 M O2 and 4.3 x 10-3 M NO2?





CONCEPT: TYPES OF EQUILIBRIUM CONSTANTS

When dealing with gases, we use the equilibrium constant, _______, which uses the partial pressure unit of ________.

When dealing with aqueous solutes, we use the equilibrium constant, ______, which uses the concentration unit of ______.

To relate KP to KC we use the formula:

EXAMPLE: For the following reaction, 2 A (s) + 3 B (g) 2 C(g), Kc = 4.9 x 10-9 at 25oC. Which of the following

statements is true?

a) The reaction is favored in the forward direction.

b) The concentration of the products is greater than the concentration of the reactants.

c) The reaction is favored in the reverse direction.

d) The value of Kp will be larger than the value of Kc.

PRACTICE: Methane (CH4) reacts with hydrogen sulfide to yield hydrogen gas and carbon disulfide, a solvent used in te manufacturing rayon and cellophane. What is the value of KC at 1000 K if the partial pressures in an equilibrium mixture at 1000 K are 0.20 atm methane, 0.15 atm hydrogen sulfide, 0.30 atm carbon disulfide and 0.10 atm hydrogen gas?

CH4 (g) + 2 H2S (g) 4 H2 (g) + CS2 (g)





PRACTICE: TYPES OF EQUILIBRIUM CONSTANTS (CALCULATIONS)

PRACTICE 1: In which of the given reactions is Kp greater than, less than and equal to Kc?

a) SO3 (g) + NO (g) SO2 (g) + NO2 (g)

b) P4 (s) + 5 O2 (g) P4O10 (s)

c) 4 NH3 (g) + 3 O2 (g) 2 N2 (g) + 6 H2O (g)

PRACTICE 2: Given the hypothetical reaction 2 A (s) + ? B (g) 3 C (g), Kp = 0.0105 and Kc = 0.45 at 250

degrees Celsius. What is the value of the coefficient of B?





CONCEPT: HESS’S LAW….KIND OF

We learned to find the total enthalpy change of a reaction by taking into account each individual reaction step, now we will

do it all over again but in finding the rate constant, K.

EXAMPLE 1: The equilibrium constant K for the reaction

CO2 (g) CO (g) + O2 (g)

is 6.83 x 10-12 at 1000 K. Calculate K for the reaction

4 CO (g) + 2 O2 (g) 4 CO2 (g)

EXAMPLE 2: Calculate the rate constant, Kc, for the reaction below:

H (g) + Br (g) HBr (g)

Use the following information to calculate Kc.

H2 (g) 2 H (g) Kc = 11.8

Br2 (g) 2 Br (g) Kc = 1.15 x 10-5

H2 (g) + Br2 (g) 2 HBr (g) Kc = 2.78 x 103





CONCEPT: CALCULATING EQUILIBRIUM CONCENTRATIONS

Sometimes you will be asked to calculate concentrations at equilibrium after being given initial concentrations. To do this we

use our favorite friend the _____________ Chart.

EXAMPLE 1: We have a solution where Ag(CN)2 –(g), CN – (g), and Ag+ (g) have an equilibrium constant, K, equal to 1.8 x

10-19. If the equilibrium concentrations of Ag(CN)2 – and CN – are 0.030 and 0.10 respectively, what is the equilibrium

concentration of Ag+?

Ag(CN)2– (g) 2 CN– (g) + Ag+ (g)

EXAMPLE 2: We place 2.5 mol of CO and 2.5 mol of CO3 in a 10.0 L flask and let the system come to equilibrium. What will

be the final concentration of CO2?

CO (g) + CO3 (g) 2 CO2 (g) K = 0.47

PRACTICE: For the reaction: N2 (g) + 2 O2 (g) 2 NO2 (g), KC = 8.3 x 10-10 at 25oC. What is the concentration of N2 gas at equilibrium when the concentration of NO2 is twice the concentration of O2 gas?




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PRACTICE: CALCULATING EQUILIBRIUM CONCENTRATIONS (CALCULATIONS 1)

EXAMPLE 1: When 0.600 atm of NO2 was allowed to come to equilibrium the total pressure was 0.875 atm. Calculate the Kp of the reaction.

2 NO2 (g) 2 NO (g) + O2 (g)

EXAMPLE 2: An important reaction in the formation of acid rain is,

2 SO2 (g) + O2 (g) 2 SO3 (g)

Initially, 0.023 M SO2 and 0.015 M O2 are mixed and allowed to react in an evacuated flask at 340 oC. When an equilibrium

is established the equilibrium amount of SO3 was found to be 0.00199 M. Calculate the equilibrium constant, Kc, for the

reaction at 340 oC.





PRACTICE: CALCULATING EQUILIBRIUM CONCENTRATIONS (CALCULATIONS 2)

EXAMPLE: If Kc is 32.7 at 300oC for following reaction:

H2 (g) + Br2 (g) 2 HBr (g)

What is the concentration of H2 at equilibrium if a 20.0 L flask contains 5.0 mol HBr initially?

PRACTICE: At a given temperature the gas phase reaction: H2 (g) + O2 (g) 2 NO (g) has an equilibrium constant

of 4.00 x 10-15. What will be the concentration of NO at equilibrium if 2.00 moles of nitrogen and 6.00 moles oxygen are

allowed to come to equilibrium in a 2.0 L flask.





CONCEPT: THE EQULIBRIUM CONSTANT AND THERMODYNAMICS

In Chapters 12 and 13 you learned that ________________ studied the rate at which our reactants changed into products.

In Chapter 14 you will learn that _______________________________ deals with the direction that a chemical reaction at equilibrium will shift.

_____________________________ Principle states that once a system that is at equilibrium is disturbed it will do whatever it can to get back to equilibrium.

EXAMPLE: For the following endothermic reaction Kc = 6.73 x 103. Predict in which direction the reaction will proceed.

4 NH3 (g) + 3 O2 (g) 2 N2 (g) + 6 H2O (g)

a) Addition of a catalyst b) Decreasing the volume

c) Removing H2O (g) d) Increasing the Temperature

e) Addition of NH3 (g) f) Decreasing the pressure

g) Removing H2O (l) h) Addition of a precipitate

i) The addition of an inert gas at constant volume.




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PRACTICE: THE EQULIBRIUM CONSTANT AND THERMODYNAMICS (CALCULATIONS 1)

PRACTICE 1: Consider the reaction below:

CH4 (g) + F2 (g) CF4 (g) + HF (g) ∆H = + 38.1 KJ/mol

The following changes will shift the equilibrium to the left except one. Which one would not cause a shift to the left?

a) Add some CF4.

b) Remove some F2.

c) Decrease the Temperature.

d) Decrease the container volume.

e) Increase the partial pressure of HF.

PRACTICE 2: The following data was collected for the following reaction at equilibrium.

2 A (g) + 3 B (g) C (g)

At 25oC, K is 5.2 x 10-4 and at 50oC K is 1.7 x 10-7. Which of the following statements is true?

a) The reaction is exothermic.

b) The reaction is endothermic.

c) The enthalpy change, ΔH, is equal to zero.

d) Not enough information is given.





PRACTICE: THE EQULIBRIUM CONSTANT AND THERMODYNAMICS (CALCULATIONS 2)

PRACTICE 1: Which direction will the following reaction (in a 10.0 L flask) proceed if a catalyst is added to the system?

CaCO3 (s) CaO (s) + CO2 (g) Kp = 3.2 x 10-28

a) To the right.

b) To the left.

c) The equilibrium position will not change but the rate will increase.

d) The equilibrium position will not change but the concentrations of everything will increase.

PRACTICE 2: Consider the following gas reaction of A2 ( shaded spheres) and B2 ( unshaded spheres)

A2 (g) + B2 (g) 2 AB (g) Which container proceeds more to completion?

A B





CONCEPT: ACID IDENTIFICATION

The most common feature of an acid is that many possess an H+ ion called the _______________________________ .

When it comes to acids there are 2 MAJOR TYPES that exist:

_______________________ are acids where the H+ ion is attached to an electronegative element.

• These types of acids lack the element __________________ and usually possess no __________________ .

• The most common type of these particular acids are the haloacids: _______ , _______ , _______ & _______ .

_______________________ are acids that contain the ________________ , ________________ & ________________.

• They are created by the hydration of nonmetal oxides.

PRACTICE: Which of the following compound(s) cannot be classified as an acid?

a) H2S b) HCN c) H2 d) C6H6 e) All are acids.


CH.11 - ACIDS & BASES


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CONCEPT: BINARY ACID STRENGTH

STRONG ACIDS are considered _________________ Electrolytes so they ionize completely in water.

HCl (aq) H2O

H+ (aq) + Cl – (aq)

WEAK ACIDS are considered __________________ Electrolytes so they don’t completely ionize in water.

HF + H2O F – (aq) + H3O+ (aq)

The strength of a BINARY ACID is based on the _________________________ or ________________ of the nonmetal.

• For elements in the same period then look at their __________________ . The ________, the ________ acidic.

• For elements in the same group then look at their __________________ . The ________, the ________ acidic.

BINARY ACID STRENGTH

PRACTICE 1: Which is the weakest acid from the following?

a) H2S b) H2Se c) H2Te d) All would have the same acid strength.

PRACTICE 2: Which of the following acids would be classified as the strongest?

a) CH4 b) NH3 c) H2O d) HF e) PH3





CONCEPT: OXYACID STRENGTH

The strength of OXYACIDS is based on the number of _____________ or the _____________________ of the nonmetal.

• RULE: If my oxyacid has 2 or More ___________ than ____________ then my oxyacid is a __________ ACID.

HNO3 ___ Oxygens – ___ Hydrogens

C6H5OH ___ Oxygens – ___ Hydrogens

HBrO4 ___ Oxygens – ___ Hydrogens

When comparing the strengths of different oxyacids remember:

• If they have different number of oxygens then the _________ oxygen the ___________ acidic

• If they have the same number of oxygens then the _________ electronegative the nonmetal the ________ acidic.

ElectronegativityH2C2O4 ___ Oxygens – ___ Hydrogens

HSO4 –

___ Oxygens – ___ Hydrogens

Two Exceptions

PRACTICE: Rank the following oxyacids in terms of increasing acidity.

a) HClO3 b) HBrO4 c) HBrO3 d) HClO4





CONCEPT: BASE STRENGTHS

STRONG BASES are considered _________________ Electrolytes so they ionize completely in water.

NaOH (aq) H2O

Na+ (aq) + OH – (aq)

WEAK BASES are considered __________________ Electrolytes so they don’t completely ionize in water.

NH3 + H2O NH4+ (aq) + OH – (aq)

Bases possess THREE major features: __________________ or __________________ or __________________ .

Group ________:

• Any Group ______ metal when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.

Group ________:

• Any Group ______ metal, from _____ to _____ , when combined with OH –, H –, O2– or NH2 – makes a STRONG

BASE.

_____________:

• ____________________________________ are considered WEAK BASES.

Ex:

• ____________________________________ are considered WEAK ACIDS.

Ex:





PRACTICE: ACID & BASE IDENTIFICATION

EXAMPLE: Classify each of the following as a strong acid, weak acid, strong base or weak base.

a) HCHO2 c) H2NNH2

b) (CH3CH2)3NH+ d) HBrO3

PRACTICE 1: Classify each of the following as a strong acid, weak acid, strong base or weak base.

a) KOCH3 b) CH3OH


a) HOCN b) H5IO6


a) NaN3 b) SrH2





CONCEPT: ARRHENIUS ACIDS & BASES

The most general definition for acids and bases was developed by Svante Arrhenius near the end of the 19th century.

• According to him, the _______ cation and the _________ anion are fundamental to the concept of acids and bases.

• His definition however failed to describe acidic and basic behavior in nonaqueous media.

The Arrhenius definition states an acid is a compound that increases _______________ when dissolved in a solvent.

The Arrhenius definition states a base is a compound that increases _______________ when dissolved in a solvent.

PRACTICE 1: Which ions are formed from the dissociation of the following compound?

a) Sr(OH)2 (s) Dissolves in H2O


a) H2SO4 (l) Dissolves in H2O


a) HBO32-

Dissolves in H2O




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CONCEPT: BRONSTED LOWRY ACIDS & BASES

In 1923, Johannes Brønsted and Thomas Lowry developed a new set of definitions for acids and bases.

According to the Bronsted-Lowry definition, acids are considered _____________________________ and bases are

considered _____________________________.

• Unlike Arrhenius acids and bases, they are not limited to aqueous solutions.

• Every Arrhenius acid is a Brønsted-Lowry acid (and likewise for the bases).

• Brønsted-Lowry acids and bases always occur in pairs called _____________________________________ .

EXAMPLE 1: Write the formula of the conjugate base for the following compound:

HSO4 –

EXAMPLE 2: Write the formula of the conjugate acid for the following compound:

V2O52-

PRACTICE 1: Write the formula of the conjugate base for the following compound:

H2Se

PRACTICE 2: Write the formula of the conjugate for the following compound:

NH2NH2




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PRACTICE: BRONSTED LOWRY ACIDS & BASES (CALCULATIONS)

EXAMPLE 1: Identify the acid, base, conjugate acid and conjugate base in the following reactions:

a) HF (aq) + H2O (aq) F – (aq) + H3O+ (aq)

EXAMPLE 2: Identify the acid, base, conjugate acid and conjugate base in the following reactions:

a) CN – (aq) + H2O (aq) HCN (aq) + OH – (aq)

PRACTICE 1: Which of the following is a Bronsted-Lowry acid? a) CH4 b) HCN c) NH3 d) Br2

PRACTICE 2: Determine the chemical equation that would result when carbonate, CO32-, reacts with water.




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CONCEPT: AMPHOTERIC SPECIES

An amphoteric, or _________________________, is a species that can act as a(n) ACID or BASE.

• Water is prime example of an amphoteric species.

Partially dissociated conjugate bases of polyprotic acids are also amphoteric.

• These compounds possess _________________ and a __________________________.

Ex:

PRACTICE: Which of the following species is/are amphoteric?

a) CO32– b) HF c) NH4

+ d) HPO32- e) CH3O –




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CONCEPT: LEWIS…THE FINAL TYPE OF ACID & BASE

In the 1920s, Gilbert Lewis proposed a new set of definitions for acids and bases.

A Lewis acid is a(n) ______________________________________.

• ________ acts as a Lewis acid when connected to an electronegative element: ___ , ___ , ___ , ___ , or ________

• _____________________ charged hydrogen or metals.

• If your central element has _________________ 8 valence electrons.

A Lewis base is a(n) ______________________________________.

• Compounds with _________________________ .

NH3 H2O CH3OH CH3OCH3

• Compounds with a _________________________ .

CN – OH – CH3O – N3–




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PRACTICE: LEWIS….THE FINAL TYPE OF ACID & BASE (CALCULATIONS)

EXAMPLE: Identify each of the compounds in the following chemical equation.

H3CH2C O

CH2H3C

Al

Br

Br Br

H3CH2C O

CH2H3C

Al

Br

Br

Br

PRACTICE 1: Identify the Lewis acids and bases in the following reactions.

a) H+ + OH – H2O

b) Cl – + BCl3 BCl4–

c) SO3 + H2O H2SO4

PRACTICE 2: Identify each of the following compounds as either a Lewis acid, a Lewis base or neither.

a) ZnCl2 b) CN –

c) NH4+ d) Co3+




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CONCEPT: pH and pOH

To deal with incredibly small concentration values of [H+] and [OH-] we can use the pH scale.

• Under normal conditions, the pH scale operates within the range of ______ to ______ .

By taking the – log of [H+] and [OH-] we can find pH and pOH.

pH = − log[H+ ] pOH = − log[OH− ] p = − log

By recognizing the relationship between [H+] and [OH-] with pH and pOH we can create new formula relationships.

pH = − log[H+ ] pOH = − log[OH− ]

A species with a pH greater than 7 is classified as _____________ and the [H+] is ___________________ than the [OH-].

• The ______________ the base then the ______________ the pH.

A species with a pH less than 7 is classified as _______________ and the [H+] is ____________________ than the [OH-].

• The ______________ the acid then the ______________ the pH.

A species with a pH equal to 7 is classified as ______________ and the [H+] is _____________________ than the [OH-].

By using – log with the equilibrium expression for water a relationship between pH and pOH can be created.

pH+ pOH =14




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PRACTICE: pH and pOH (CALCULATIONS 1)

EXAMPLE: What is the hydroxide ion and hydronium ion concentration of an aqueous solution that has a pH equal to 6.12?

PRACTICE 1: Which of the following solutions will have the lowest concentration of hydronium ions?

a) 0.100 moles C6H5NH2

b) 0.100 moles Be(OH)2

c) 0.100 moles SrH2

d) 0.100 moles (CH3)2NH

PRACTICE 2: Which of the following statements about aqueous solutions is/are true?

a) For an basic solution the concentration of H3O+ is greater than the concentration of OH –.

b) The pH of a neutral aqueous solution is 7.00 at all temperatures.

c) An acidic solution under normal conditions has a pH value less than 7.00.

d) If the concentration of H3O+ decreases then the concentration of OH – will also decrease.

e) The pH of aqueous solutions is less than 7.





PRACTICE: pH and pOH (CALCULATIONS 2)

EXAMPLE: A solution is prepared by dissolving 0.235 mol Sr(OH)2 in water to produce a solution with a volume of 750 mL.

a) What is the [OH-]?

b) What is the [H+]?

PRACTICE: What is the Kw of pure water at 20.0°C, if the pH is 7.083?

a) 8.26 × 10-8 b) 6.82 × 10-15 c) 7.23 × 10-14 d) 1.00 × 10-14





CONCEPT: AUTO IONIZATION OF WATER

Water can react with itself in a reaction called self–ionization where ______________ and ______________ are produced.

H2O (l) + H2O (l)

This reaction is usually written more simply as:

H2O (l) + H2O (l)

The equilibrium equation for water is called the ________________________ (KW) for water and is given by the following:

KW = [H+ ][OH− ]

At 25oC, KW = ___________________, but remember KW, like all other constants K, is temperature dependent.

• Increasing the temperature will ______________ KW.

Constant

0oC

10 oC

50 oC

100 oC

KW

1.14 x 10-14

2.93 x 10-14

5.476 x 10-14

5.13 x 10-13

EXAMPLE: Determine the concentration of hydronium ions for a neutral solution at 25oC and at 50oC.





CONCEPT: CALCULATING pH and pOH OF STRONG SPECIES

STRONG ACIDS & BASES are considered _________________ Electrolytes so they ionize completely in water.

HCl (aq) H2O

H+ (aq) + Cl – (aq)

NaOH (aq) H2O


EXAMPLE 1: Calculate the pH of a 0.0782 M solution of CaH2.

EXAMPLE 2: Calculate the pH of a 0.000550 M HBr solution to the correct number of significant figures.

a) 3.3

b) 3.26

c) 3.260

d) 3.2596

e) All are correct

PRACTICE: Calculate the pH of 50.00 mL of 4.3 x 10-7 M H2SO4.





CONCEPT: CALCULATING pH and pOH OF WEAK SPECIES

WEAK ACIDS & BASES are considered __________________ Electrolytes so they don’t completely ionize in water.

HF + H2O F – (aq) + H3O+ (aq)

NH3 + H2O NH4+ (aq) + OH – (aq)

EXAMPLE 1: Pryridine, an organic molecule, is a very common weak base.

C5H5N (aq) + H2O (l) C5H5NH+ (aq) + OH- (g)

Assume you have a 0.0225 M aqueous solution of pyridine, C5H5N, determine its pH. The Kb value for the

compound is 1.5 x 10-9.





PRACTICE: CALCULATING pH and pOH OF WEAK SPECIES (CALCULATIONS 1)

EXAMPLE: An unknown weak base has an initial concentration of 0.750 M with a pH of 8.03. Calculate its equilibrium base constant.

PRACTICE: Determine the pH of a solution made by dissolving 6.1 g of sodium cyanide, NaCN, in enough water to make a

500.0 mL of solution. (MW of NaCN = 49.01 gmol

). The Ka value of HCN is 4.9 x 10-10.





CONCEPT: ACID & BASE CONSTANTS As you might already realize, there are relatively few strong acids. The great majority of acids are weak acids.

Consider a weak monoprotic acid, HA, and its ionization in water:

HA (aq) + H2O (l) A – (aq) + H3O+ (aq)

The equilibrium expression for this ionization would be:

Ka =

ProductsReac tan ts

=

Where Ka represents the _________________________________________ and it measures the strength of weak acids.

When looking at weak bases we don’t use Ka, but instead _______, which represents the __________________________.

• The relationship between Ka and Kb can be expressed with the following equation:

In general, the __________________ the Ka the stronger the acid and the __________________ the concentration of H+.

In general, the __________________ the pKa the stronger the acid and the __________________ the concentration of H+.

PRACTICE: If the Kb of NH3 is 1.76 x 10-5, determine the acid dissociation constant of its conjugate acid.

KW =Ka ⋅Kb




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PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 1)

EXAMPLE 1: Knowing that HF has a higher Ka value than CH3COOH, determine, if possible, in which direction the following equilibrium lies.

HF (aq) + CH3COO – (aq) F – (aq) + CH3COOH (aq)

a) Equilibrium lies to the left.

b) Equilibrium lies to the right.

c) Equilibrium is equal and balanced.

d) Not enough information given.

EXAMPLE 2: What is the equilibrium constant for the following reaction and determine if reactants or products are favored.

HCN (aq) + ClO2 – (aq) CN – (aq) + HClO2 (aq)

The acid dissociation constant of HCN is 4.9 x 10-10 and the acid dissociation of HClO2 is 1.1 x 10-2.

HCN (aq) + H2O (aq) CN – (aq) + H3O+ (aq)

HClO2 (aq) + H2O (aq) ClO2 – (aq) + H3O+ (aq)





PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 2)

EXAMPLE: Which of the following solutions will have the lowest pH?

a) 0.25 M HC2F3O2

b) 0.25 M HIO3

c) 0.25 M HC3H5O3

d) 0.25 M H2CO3

e) 0.25 M HSeO4 –

PRACTICE 1: Which Bronsted-Lowry base has the greatest concentration of hydroxide ions?

a) C2H8N2 (Kb = 8.3 x 10-5)

b) C5H5N (Kb = 1.7 x 10-9)

c) (CH3)3N (Kb = 1.0 x 10-6)

d) C3H7NH2 (Kb = 3.5 x 10-4)

e) C6H5NH2 (Kb = 3.9 x 10-10)

PRACTICE 2: Which Bronsted-Lowry acid has the weakest conjugate base?

a) HCNO (Ka = 2.0 x 10-4)

b) HF (Ka = 3.5 x 10-4)

c) HN3 (Ka = 2.5 x 10-5)

d) H2CO3 (Ka = 4.3 x 10-7)





CONCEPT: ACID & BASE NEUTRALIZATION

When an acid neutralizes a base an ionic compound called a _______________ is formed.

• These solutions can be neutral, acidic or basic, based on acid-base properties of the cations and anions formed.

RULES FOR IDENTIFYING YOUR IONS

CATIONS (POSITIVE IONS)

1) Transition Metals: If your transition metal has a charge of +2 or higher it is acidic. If the charge is less than +2 then it is

neutral.

EX:

2) Main-Group Metals: If your main-group metal has a charge of +3 or higher it is acidic. If the charge is less than +3 then

it is neutral.

EX:

3) Positive Amines are acidic.

EX:

ANIONS (NEGATIVE IONS)

1) NEGATIVE ION: If you have a negative ion then add an H+ to it. If you create a weak acid then your negative ion is

basic.

EX:




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PRACTICE: ACID & BASE NEUTRALIZATION 1

EXAMPLE: Determine if each of the following compounds will create an acidic, basic or neutral solution.

a) NaOCl b) PbCl4

PRACTICE 1: Determine if each of the following compounds will create an acidic, basic or neutral solution.

a) LiC2H3O2 b) C6H5NH3Br


a) Co(HSO4)2 b) Sr(HSO3)2


a) C3H7NH3F






EXAMPLE 1: Determine whether each compound will become more soluble in an acidic solution.

a) NaBr b) LiCl c) KIO

EXAMPLE 2: Determine the pH of a 0.50 M NH4Cl solution. The Kb of NH3 is 1.75 x 10-5.

PRACTICE: Determine the pH of a 0.55 M NaCN solution. The Ka of hydrocyanic acid, HCN, is 4.9 x 10-10.





CONCEPT: ACID IDENTIFICATION

The most common feature of an acid is that many possess an H+ ion called the _______________________________ .

When it comes to acids there are 2 MAJOR TYPES that exist:

_______________________ are acids where the H+ ion is attached to an electronegative element.

• These types of acids lack the element __________________ and usually possess no __________________ .

• The most common type of these particular acids are the haloacids: _______ , _______ , _______ & _______ .

_______________________ are acids that contain the ________________ , ________________ & ________________.

• They are created by the hydration of nonmetal oxides.

PRACTICE: Which of the following compound(s) cannot be classified as an acid?

a) H2S b) HCN c) H2 d) C6H6 e) All are acids.


CH.12 - AQUEOUS EQUILIBRIA


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CONCEPT: BINARY ACID STRENGTH

STRONG ACIDS are considered _________________ Electrolytes so they ionize completely in water.

HCl (aq) H2O

H+ (aq) + Cl – (aq)

WEAK ACIDS are considered __________________ Electrolytes so they don’t completely ionize in water.

HF + H2O F – (aq) + H3O+ (aq)

The strength of a BINARY ACID is based on the _________________________ or ________________ of the nonmetal.

• For elements in the same period then look at their __________________ . The ________, the ________ acidic.

• For elements in the same group then look at their __________________ . The ________, the ________ acidic.

BINARY ACID STRENGTH

PRACTICE 1: Which is the weakest acid from the following?

a) H2S b) H2Se c) H2Te d) All would have the same acid strength.

PRACTICE 2: Which of the following acids would be classified as the strongest?

a) CH4 b) NH3 c) H2O d) HF e) PH3





CONCEPT: OXYACID STRENGTH

The strength of OXYACIDS is based on the number of _____________ or the _____________________ of the nonmetal.

• RULE: If my oxyacid has 2 or More ___________ than ____________ then my oxyacid is a __________ ACID.

HNO3 ___ Oxygens – ___ Hydrogens

C6H5OH ___ Oxygens – ___ Hydrogens

HBrO4 ___ Oxygens – ___ Hydrogens

When comparing the strengths of different oxyacids remember:

• If they have different number of oxygens then the _________ oxygen the ___________ acidic

• If they have the same number of oxygens then the _________ electronegative the nonmetal the ________ acidic.

Electronegativity

H2CO4 ___ Oxygens – ___ Hydrogens

HSO4 –

___ Oxygens – ___ Hydrogens

Two Exceptions

PRACTICE: Rank the following oxyacids in terms of increasing acidity.

a) HClO3 b) HBrO4 c) HBrO3 d) HClO4





CONCEPT: BASE STRENGTHS

STRONG BASES are considered _________________ Electrolytes so they ionize completely in water.

NaOH (aq) H2O


WEAK BASES are considered __________________ Electrolytes so they don’t completely ionize in water.

NH3 + H2O NH4+ (aq) + OH – (aq)

Bases possess THREE major features: __________________ or __________________ or __________________ .

Group ________:

• Any Group ______ metal when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.

Group ________:

• Any Group ______ metal, from _____ to _____ , when combined with OH –, H –, O2– or NH2 – makes a STRONG

BASE.

_____________:

• ____________________________________ are considered WEAK BASES.

Ex:

• ____________________________________ are considered WEAK ACIDS.

Ex:





CONCEPT: CLASSIFICATION AND IDENTIFICATION OF BUFFERS

Solutions which contain a _________________ acid and its __________________ base are called buffer solutions because they resist drastic changes in pH.

• They resist drastic changes in pH by keeping ___________ and ___________ constant.

• Adding a small amount of STRONG BASE and, the pH ____________, but not by much because the

________________ neutralizes the STRONG BASE added.

• Adding a small amount of STRONG ACID and, the pH ____________, but not by much because the

________________ neutralizes the STRONG ACID added.

PRACTICE 1: Which one of the following combinations does not create a buffer?

a) HC2H3O2 and K C2H3O2

b) H2SO3 and NaHSO3

c) H3PO4 and NaH2PO4

d) HNO3 and KNO3

e) NH4Cl and NH3

PRACTICE 2: Which of the following combinations can result in the formation of a buffer?

a) HF and HI

b) HC2H3O2 and NH3

c) CH3CH2NH2 and CH3CH2NH3+

d) NaCl and NaOH




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CONCEPT: CLASSIFICATION AND IDENTIFICATION OF BUFFERS

Solutions which contain a _________________ acid and its __________________ base are called buffer solutions because they resist drastic changes in pH.

• They resist drastic changes in pH by keeping ___________ and ___________ constant.

• Adding a small amount of STRONG BASE and, the pH ____________, but not by much because the

________________ neutralizes the STRONG BASE added.

• Adding a small amount of STRONG ACID and, the pH ____________, but not by much because the

________________ neutralizes the STRONG ACID added.

PRACTICE 1: Which one of the following combinations does not create a buffer?

a) HC2H3O2 and K C2H3O2

b) H2SO3 and NaHSO3

c) H3PO4 and NaH2PO4

d) HNO3 and KNO3

e) NH4Cl and NH3


a) HF and HI

b) HC2H3O2 and NH3

c) CH3CH2NH2 and CH3CH2NH3+

d) NaCl and NaOH





CONCEPT: CREATING A BUFFER

There are 3 ways to form a buffer:

1) Mixing a ______________ acid and its ______________ base.

2) Mixing a ______________ acid and a ______________ base.

3) Mixing a ______________ acid and a ______________ base.





PRACTICE: CREATING A BUFFER

EXAMPLE: Which of the following combinations can result in the formation of a buffer?

a) 0.01 moles HClO (hypochlorous acid) and 0.05 moles of NaOH.

b) 0.01 moles HClO (hypochlorous acid) and 0.05 moles of HCl.

c) 0.01 moles HClO (hypochlorous acid) and 0.05 moles of NH3.

d) 0.01 moles HClO (hypochlorous acid) and 0.001 moles of NaOH


a) 50 mL of 0.10 M HF with 50 mL of 0.10 M NaOH.

b) 50 mL of 0.10 M HNO2 with 25 mL of 0.10 M Ca(OH)2.

c) 50 mL of 0.10 M CH3CO2H with 60 mL of 0.10 M NaOH.

d) 50 mL of 0.10 M HF with 30 mL of 0.10 M NaOH.

PRACTICE 2: A buffer solution is comprised of 50.0 mL of a 0.100 M HC2H3O2 and 60.0 mL of a 0.100 M NaC2H3O2. Which

of the following actions would completely destroy the buffer?

a) Adding 0.003 mol HC2H3O2

b) Adding 0.007 mol Ca(C2H3O2)2

c) Adding 0.005 mol NaOH

d) Adding 0.004 mol HCl

e) Adding 0.001 mol HCl





CONCEPT: CALCULATING THE pH OF BUFFERS

We learned that whenever we had a(n) ____________ acid or base we were supposed to use our favorite friend the

_____________ Chart in order to calculate the pH or pOH.

Now, whenever we have a buffer solution we can skip it and use the ___________________________________ Equation.

Buffer Equation:

pH = pKa + log (conjugate base)(weak acid)

EXAMPLE 1: What is the pH of a solution consisting of 2.75 M sodium phenolate (C6H5ONa) and 3.0 M phenol (C6H5OH).

The Ka of phenol is 1.0 x 10-10.

PRACTICE: Calculate the pH of a solution formed by mixing 200 mL of a 0.400 M C2H5NH2 solution with 350 mL of a 0.450

M C2H5NH3+ solution. (Kb of C2H5NH2 is 5.6 x 10-4).





PRACTICE: CALCULATING THE pH OF BUFFERS (PART 1)

EXAMPLE 1: What is the buffer component concentration ratio, Pr–

HPr, of a buffer that has a pH of 5.11. (The Ka of HPr is

1.30 x 10-5).

EXAMPLE 2: Over what pH range will an oxalic acid (H2C2O4) / sodium oxalate (NaHC2O4) solution work most effectively?

The acid dissociation constant of oxalic acid is 6.0 x 10-2.

a) 0.22 – 2.22 b) 1.00 – 3.00 c) 0.22 – 1.22 d) 2.0 – 4.0

PRACTICE: Determine how many grams of sodium acetate, NaCH3CO2 (MW: 82.05 g/mol), you would mix into enough

0.065 M acetic acid CH3CO2H (MW: 60.05 g/mol) to prepare 3.2 L of a buffer with a pH of 4.58. The Ka is 1.8 x 10-5.





PRACTICE: CALCULATING THE pH OF BUFFERS (PART 2)

EXAMPLE: Which weak acid-conjugate base combination would be ideal to form a buffer with a pH of 4.74.

a) Cyanic acid and Potassium cynate (Ka = 4.9 x 10-10)

b) Benzoic acid and Lithium benzoate (Ka = 6.3 x 10-5)

c) Acetic acid and Sodium acetate (Ka = 1.7 x 10-5)

d) Ammonium chloride and Ammonia (Ka = 5.56 x 10-10)

e) Formic acid and Cesium formate (Ka = 1.7 x 10-4)

PRACTICE: A buffer solution is made by combining a weak acid with its conjugate salt. What will happen to the pH if the solution is diluted to one-fourth of its original concentration?

a) The pH will increase.

b) The pH will decrease.

c) The pH will remain constant.

d) The solution will become more neutral.





CONCEPT: pH TITRATION CURVES

The shape of a pH titration curve makes it possible to identify the equivalence point, the point at which _________________

of acid and base are mixed together.

At the equivalence point, the pH of a strong acid and strong base is ___________ 7.

At the equivalence point, the pH of a strong acid and weak base is ___________ 7.

At the equivalence point, the pH of a weak acid and strong base is ___________ 7.




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PRACTICE: pH TITRATION CURVES (CALCULATIONS 1)

EXAMPLE: The following questions refer to the titration curve given below. a) The titration curve shows the titration of a strong acid a weak acid a strong base a weak base with a strong base with a strong base with a strong acid with a strong acid b) Which point on the titration curve represents a region where a buffer solution has formed? point A point B point C point D c) Which point on the titration curve represents the equivalence point? point A point B point C point D d) Which of the following would be the best indicator to use in the titration? erythrosin B methyl red bromthymol blue o-cresonphthalein pKa = 2.9 pKa = 5.4 pKa = 6.8 pKa = 9.0






EXAMPLE 1: The acid form of an indicator is red and its anion is blue. The Ka value for this indicator is 10-9. What will be

the approximate pH range over which this indicator changes color?

a) 3-5 b) 4-6 c) 5-7 d) 8-10 e) 9-11

PRACTICE : What will be the color of the indicator in the above question in a solution that has a pH of 6?

EXAMPLE 2: Consider the titration of 100.0 mL of 0.016 M H2SO4 with 0.400 M NaOH at the equivalence point. How many

many milliters of 0.400 M NaOH are required to reach the equivalence point?






EXAMPLE: Consider the titration of 40.0 mL of 0.0800 M HCl with 0.0160 M Ca(NH2)2.

a) How many milliliters of 0.0160 M Ca(NH2)2 are required to reach the equivalence point?

b) What is the pH of this solution?

PRACTICE: Consider the titration of 60.0 mL of 0.200 M a H3PO3 solution with 0.350 M potassium hydroxide, KOH solution.

How many milliliters of base would be required to reach each of its equivalence points?





CONCEPT: WEAK ACID & BASE TITRATIONS

In the past, we reacted WEAK acids or bases with __________________ and used a(n) _________________ Chart.

Remember in this case the units in this chart are in _______________.

Now we will react WEAK acids and bases with ______________ acids and bases.

When you react a WEAK species with a ____________ species you have to use a(n) _________________ Chart.

In this case the units must be in _______________.

EXAMPLE: Consider the titration of 75.0 mL of 0.0300 M H3C3O3 (Ka = 4.1 X 10-3) with 12.0 mL of 0.0450 M KOH.

Calculate the pH.

PRACTICE: Calculate the pH of the solution resulting from the mixing of 75.0 mL of 0.100 M NaC2H3O2 and 75.0 mL of 0.60

M HC2H3O2 with 0.0040 moles of HBr.





PRACTICE: WEAK ACID & BASE TITRATIONS (CALCULATIONS 1)

EXAMPLE: A buffer contains 167.2 mL of 0.25 M propanoic acid, CH3CH2COOH, with 138.7 mL of 0.42 M sodium

propanoate, CH3CH2COONa. Find the pH after the addition of 150.2 mL of 0.56 M HCl. (The Ka of CH3CH2COOH is 1.3 x

10-5).

PRACTICE: In order to create a buffer 7.510 g of sodium cyanide is mix with 100.0 mL of 0.250 M hydrocyanic acid, HCN.

What is the pH of the buffer solution after the addition of 175.0 mL of 0.300 M NaNH2?





EXAMPLE: Consider the titration of 75.0 mL of 0.60 M HNO2 with 0.100 M NaOH at the equivalence point. What would be

the pH of the solution at the equivalence point? The Ka of HNO2 is 4.6 x 10-4.





CONCEPT: STRONG ACID & BASE TITRATIONS

Whenever we had a STRONG ACID or STRONG BASE we ____________ use an ICE CHART.

Now, whenever you titrate a STRONG ACID with STRONG BASE you ____________ use an ICF CHART.

EXAMPLE: Calculate the pH of the solution resulting from the titration of 75.0 mL of 0.100 M HBrO4 with 55.0 mL of 0.100

M NaNH2.

PRACTICE: Calculate the pH of the solution resulting from the mixing of 175.0 mL of 0.250 M HNO3 with 75.0 mL of 0.200

M Ba(OH)2.





CONCEPT: CALCULATING SOLUBILITY FROM Ksp

__________________________ was the maximum amount of solute that could successfully dissolved in a solvent.

In this section we learn that associated with any solid is a ___________ value, which stands for the solubility

product constant.

The larger the solubility product constant then the _________ soluble the solid is in a solvent.

The smaller the solubility product constant then the _________ soluble the solid is in a solvent.

EXAMPLE: Consider the following compounds. Which has the highest concentration of OH- ions?

a) Co(OH)2 Ksp = 1.3 x 10-15

b) Cu(OH)2 Ksp = 2.2 x 10-20

c) Ni(OH)2 Ksp = 6.0 x 10-16

d) Fe(OH)2 Ksp = 4.1 x 10-15

e) Zn(OH)2 Ksp = 3.0 x 10-16

EXAMPLE: The solubility of silver sulfate, Ag2SO4, is 0.025 M at 25oC. Calculate its solubility product constant, Ksp.




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PRACTICE: CALCULATING SOLUBILITY FROM Ksp (CALCULATIONS 1)

EXAMPLE: Find the solubility of CoCl3 (Ksp = 2.8 x 10-13) in

a) pure water

b) 0.20 M NaCl.




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CONCEPT: OXIDATION-REDUCTION REACTIONS

Chemists use some important terminology to describe the movement of electrons.

• In ______________ reactions we have the movement of electrons from one reactant to another.

L

E

O

G

E

R

Agent Agent

Rules for Assigning an Oxidation Number (O.N.)

A. General Rules

1. For an atom in its elemental form (Na, O2, S8, etc.): O.N. = 0

2. For an ion the O.N. equals the charge: Na+ , Ca2+ , NO3 –

B. Specific Rules

1. Group 1A: O.N. = +1

2. Group 2A: O.N. = +2

3. For hydrogen: O.N. = +1 with nonmetals

O.N. = -1 with metals and boron

4. For Fluorine: O.N. = -1

5. For oxygen: O.N. = -1 in peroxides (X2O2 , X = Group 1(A) element)

O.N. = − 12

in superoxides (XO2 , X = Group 1(A) element)

O.N. = - 2 in all other compounds

6. Group 7A O.N. = -1 (except when connected to O)


CH.13 - ELECTROCHEMISTRY



CONCEPT: OXIDATION-REDUCTION REACTIONS (PRACTICE)

EXAMPLE: In the following reaction identify the oxidizing agent and the reducing agent:

a. 2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (g)

PRACTICE: What is the oxidation number of each underlined element?

a. P4 b. BO33-

c. AsO42- d. HSO4

–

PRACTICE: In the following reaction identify the oxidizing agent and the reducing agent:

a. Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O





Consider the reaction of 3.0 M dichromate: Cr2O72– Cr3+ (Acidic Medium). What would be

the coefficient on water in the balanced equation?

a. 7

b. 5

c. 6

d. 4




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Consider the equation: Cr2O72– + H+ + I– Cr3+ (aq) + H2O + I3–

Which coefficient would be needed to balance I–?

a) 9 b) 7 c) 5 d) 14





What is the coefficient of the oxidizing agent when it is balanced within a basic solution?

Cl2(g) + S2O32-(aq) Cl-(aq) + SO42-(aq)

a) 1 b) 2 c) 3 d) 4





The purpose of a galvanic cell is to:

a. Transduce chemical energy to electrical energy.

b. Purify solids.

c. Allow for oxidation without reduction.

d. To consume electricity.




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During the process for electrolysis of water a current is passed through water and produces hydrogen gas

and oxygen gas. Which of the following statements is true?

a. O2 gas is produced at the anode.

b. H2 gas is produced at the cathode.

c. In the reaction, H2 moles are twice the O2 moles.

d. All of the following are correct.





Which statement is false?

a. Reduction occurs at the cathode.

b. A reducing agent will lose electrons.

c. Cations migrate to the cathode in both electrolytic and electrochemical cells.

d. Li (s) is the strongest oxidizing agent; F2 is the strongest reducing agent.





Which of the following reactions may occur at the anode?

a. Ga3+ (aq) + 3 e– Ga (s)

b. Cu2+ (aq) + 2 e– Cu (s)

c. 2 Cl – (aq) Cl2 (g) + 3 e–

d. Co (s) + e– Co+ (aq)





Define a salt bridge.

A) A pathway, composed of salt water, that ions pass through.

B) A pathway in which no ions flow.

C) A pathway between the cathode and anode in which ions are reduced.

D) A pathway between the cathode and anode in which ions are oxidized.

E) A pathway by which counterions can flow between the half-cells without the solutions in the half-cell

totally mixing.





What statement is NOT true about standard electrode potentials?

A) E°cell is positive for spontaneous reactions.

B) Electrons will flow from more negative electrode to more positive electrode.

C) The electrode potential of the standard hydrogen electrode is exactly zero.

D) E°cell is the difference in voltage between the anode and the cathode.

E) The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative

to the standard hydrogen electrode and therefore has a positive E°.





Use the standard reduction potentials below to determine which element or ion is the best reducing agent.

Pd2+ (aq) + 2 e – Pd (s) E° = + 0.90 V

2 H+ (aq) + 2 e – H2 (g) E° = 0.00 V

Mn2+ (aq) + 2 e – Mn (s) E° = – 1.18 V

a) Pd (s) b) H+ (aq) c) Mn2+ (aq) d) H2 (g)





Consider an electrochemical cell where the following reaction takes place:

Na2O (aq) + Ba (s) 2 Na (s) + BaO (aq)

What is the cell notation for this cell?

What is the ratio of oxidizing agent to reducing agent?





CONCEPT: GALVANIC VS. ELECTROLYTIC CELLS

With REDOX reactions we will now deal with a new variable: a reaction’s cell potential, which uses the variable _________ .

• The greater this variable then the more likely ______________________ will occur.

• The smaller this variable then the more likely ______________________ will occur.

Galvanic/Voltaic Cell: ____________ or ____________ electricity so it’s a ________________ .

Galvanic/Voltaic Cell

__________________ ( – ) __________________ ( + )

Cd2+ (aq) + 2 e – Cd (s) Eo = - 0.40 V

Ni2+ (aq) + 2 e – Ni (s) Eo = - 0.25 V

For all redox reactions under standard conditions it is possible to determine if a reaction is spontaneous or not.

∆SUniverse ∆Go K Eo Reaction Under Standard State Conditions

Cell Type

Positive Negative >1 Positive

Negative Positive <1 Negative

0 0 1 0

Anode ___________

Producing ñ Voltage

[Anode] ______

[Cathode] ______

Ionization Energy __________

Cathode __________

Electron Affinity __________





CONCEPT: RATES OF CHEMICAL REACTIONS

____________________________ is the study of reaction rates, and tells us the change in concentrations of reactants or

products over a period of time.

Although a chemical equation can help us calculate the theoretical yield from reactants, it can’t tell us how fast it goes.

• Looking at a chemical reaction in the simplest way can be seen as ___________ breaking down to form

______________.

A BReaction :

0 Seconds 30 Seconds 60 Seconds 90 Seconds


CH.14 - CHEMICAL KINETICS


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CONCEPT: FACTORS INFLUENCING REACTION RATES

VS.

1. Concentration: Molecules must _________________ to react.

• Increasing the number of molecules in a container, increases their _______________ and thereby causes the rate

to increase.

A BReaction : A B+

2. Surface Area: The frequency of collisions increases with ____________________ surface area.

CH3CH2CH2CH3 VS.

H2C

H2C CH2

CH2

3. Temperature: Increasing the temperature increases the reaction rate by increasing the ___________ and

___________ of collisions.

4. Catalyst: A catalyst increases the rate of a reaction by ______________________ the energy of activation.




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CONCEPT: GENERAL RATE

The General Rate of a chemical reaction is the change in some variable over a period of time.

Rate = Δ [A]Δ time

A = ____________ or ____________ .

The use of brackets. [ ], means ____________________ .

EXAMPLE 1: The following equation shows the production of NO and H2O by oxidation of ammonia.

4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (l)

a. What is the average rate of each compound in the balanced

equation?

b. What is the rate of NH3 in the reaction between 2 and 6 minutes at 40oC?

c. Determine the instantaneous rate of the following reaction.




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PRACTICE: STOICHIOMETRIC RATES

EXAMPLE: The decomposition of dinitrogen pentoxide is described by the chemical equation

2 N2O5 (g) 4 NO2 (g) + O2 (g)

If the rate of disappearance or decomposition of O2 is equal to 2.20 M/min at a particular moment, what is the rate of appearance or formation of N2O5 at that moment?

PRACTICE: The formation of alumina, Al2O3, can be illustrated by the reaction below:

4 Al (s) + 3 O2 (g) 2 Al2O3 (s)

At 750 K it takes 267 seconds for the initial concentration of Al2O3 to increase from 6.18 x 10-5 M to 5.11 x 10-4 M. What is

the rate of Al?




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CONCEPT: RATE LAW

Although, chemical reactions can be reversible we will only look at the ___________ reaction for chemical reactions.

• By ignoring ___________ reactions then the rate depends only on __________ concentrations and ___________.

4 NO (g) + O2 (g) 2 N2O3 (g)

Rate Law = k [NO]x [O2]y k = ___________________________________

x & y = ___________________________________

Unless the reaction is classified as a ______________ step, a rate-determining step, then x & y must be calculated

experimentally.

EXAMPLE: For the following reaction, use the given rate law to determine the best answer for the reaction with respect to each reactant and the overall order.

H2O2 (aq) + 3 I – (aq) + 2 H+ (aq) I3 – + 2 H2O (l) Rate = k [H2O2]2 [I –]

a) H2O2 is 1st order, I – is 1st order, 2nd order overall.

b) H2O2 is 2nd order, I – is 1st order, 3nd order overall.

c) H2O2 is 0th order, I – is 1st order, H+ is 1st order, 3rd order overall.

d) H2O2 is 2nd order, I – is 1st order, H+ is 0th order, 3rd order overall.




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PRACTICE: REACTION ORDERS

Answer each of the following question based on the following chemical reaction:

3 A (g) + B (g) + 2 C (g) D (g)

Experiment Initial [A] Initial [B] Initial [C] Initial Rate

1 0.0500 M 0.0500 M 0.0100 M 6.25 x 10-3

2 0.1000 M 0.0500 M 0.0100 M 1.25 x 10-2

3 0.1000 M 0.1000 M 0.0100 M 5.00 x 10-2

4 0.0500 M 0.0500 M 0.0200 M 6.25 x 10-3

EXAMPLE 1: Calculate the reaction order for reactant A.

PRACTICE: Calculate the reaction orders for reactants B and C.

EXAMPLE 2: Calculate the rate constant and the new rate for the given reaction if the the initial concentrations of [A] =

0.300 M, [B] = 0.150 M and [C] = 0.150 M.





PRACTICE: RATE CONSTANT & RATE

EXAMPLE 1: A certain chemical reaction has the given rate law:

Rate = k [A]3 [B]2 [C]-1

What are the units of the rate constant for the given reaction?

a) M-2 · s-1 b) M2 · s-1 c) M-4 · s-1 d) M-3 · s-1 e) M3 · s-1

EXAMPLE 2: The reaction of 3 A + B 2 C + D, was found to be: Rate = k [A]2 [B]3. How much would the rate

increase by if A were tripled while B were increased by half?

a) 0.50 b) 30.38 c) 0.75 d) 20.25 e) 1.125

PRACTICE: If the rate law for the following reaction is found to be the following: Rate = [Cl2 ] ⋅[HCCl3] . What are

the units for the rate constant, K?

Cl2 (g) + HCCl3 (g) HCl (g) + CCl4 (g)

a) M3/2

s b) Ms c)

1M ⋅s d)

1M3/2 ⋅s e) M ⋅s





CONCEPT: TEMPERATURE AND RATE

Temperature can have a huge effect on the reaction rate. It affects the rate by affecting the rate constant, k.

• Increasing the temperature or ______________ amount of catalyst will ______________ the rate constant, k.

• The ______________ Equation relates both the temperature and the rate constant.

For a reaction to be successful the molecules must collide with sufficient ____________________ and the correct

____________________.

Conversion of the equation into its logarithmic form can help us find Ea when two rate constants or temperatures are given.

ln k2k1=−EaR

1T2−1T1

"

#$

%

&'

A plot-wise approach can also be used to convert the equation in order to find Ea.

ln A

slope = ΔyΔx

= −EaRln k

1T

k =Ae−Ea

RT

A =

Ea =

R =

T =

lnk = − EaR

"

#$

%

&'1T"

#$

%

&'+ lnA





PRACTICE: TEMPERATURE AND RATE (CALCULATIONS)

EXAMPLE: The reaction 2 HCl (g) H2 (g) + Cl2 (g) has Ea of 1.77 x 104 kJ/mol and a rate constant of 1.32 x 10-1

at 700 K. What is the rate constant at 685 K?

PRACTICE 1: If a first order reaction has a frequency factor of 3.98 x 1013 s-1 and Ea of 160 kJ, then calculate the rate

constant at 25oC.

PRACTICE 2: Generally, the slower the rate of the reaction then the ___________ the energy of activation (Ea) and the higher the temperature, the ____________ the value of rate constant, k.

a) larger, larger b) smaller, smaller c) smaller, larger d) larger, smaller





CONCEPT: 1ST, 2ND or 0TH: WHAT’S YOUR ORDER?

If you were super observant you might have noticed that although we were talking about rate, the rate law didn’t include

___________ as a variable.

A BReaction :

The good thing is that the _______________ Rate Laws help to answer an important question in kinetics:

• “How long will it take x moles per liters of A to be consumed?”

Rate Laws Zero-Order First-Order Second-Order

Rate Equations [A]t = −kt +[A]0 ln[A]t = −kt + ln[A]0 1[A]t

= kt + 1[A]0

[A]

Time

[A]0

Slope = - k

ln[A

]

Time

ln[A]0

Slope = - k

1/[A

]

Time

1/[A]0

Slope = k

Half-Life t 12=[A]02k t 1

2=ln2k t 1

2=

1k[A]0




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PRACTICE: INTEGRATED RATE LAWS (CALCULATIONS)

EXAMPLE 1: The oxidation of ethane follows a first order mechanism, with a very high rate constant of 32 s-1, to form H2O

and CO2 as products. If the initial [C2H6] is 4.12 M, what is the concentration after 1.12 x 10-3 minutes?

EXAMPLE 2: Iodine-123 is used to study thyroid gland function. This radioactive isotope breaks down in a first order

process with a half-life of 8.50 hours at 800 K. How long will it take for the concentration of iodine-123 to be 74.1%

complete?

PRACTICE 1: At 25oC, 2 NOBr (g) 2 NO (g) + Br2 (g). The rate of the reaction is found to be: rate = k [NOBr]2.

The constant at 25oC is 7.80 x 10-4 M-1 · s-1. If 0.550 moles of HBr (g) is placed in a 5.0 L container, how long will take for

the concentration to reach 0.063 moles of HBr (g)?

PRACTICE 2: In a typical chemical reaction, nitrogen trioxide, NO3, reacts to produce nitrogen dioxide, NO2, and oxygen

gas, O. 2 NO3 (g) 2 NO2 (g) + 2 O (g)

A plot of [NO3] vs. time is linear and the slope is equal to 0.183. If the initial concentration of NO3 is 0.930 M, how long will it

take for the final concentration to reach 0.400 M?





PRACTICE: INTEGRATED RATE LAWS (CALCULATIONS 2)

EXAMPLE: Given the following graph for a second order reaction:

a) Calculate the frequency factor.

b) Calculate the energy of activation in (J/mol):

PRACTICE: The three plots were done based on a chemical reaction.

a. What is the rate constant of the reaction if it takes 21.2 minutes for the reaction to be 38.0% complete?





CONCEPT: THE TRANSITION METALS The transition metals represent elements found in the _____ – block of the periodic table.

Whereas the main group elements show similar chemical behavior because of their valence electrons, transition

metal similarities are treated differently.

• Transition metals show great chemical similarities in both their horizontal __________________ and their

vertical __________________.

• In the gradual addition of electrons to transition metals new electrons are added to the inner core electrons,

which do not participate in chemical bonding.

• For transition metals each additional electron is added to the _____ – block orbitals, while for lanthanides

and actinides they are added to the _____ – block orbitals.


CH.16 - THE ELEMENTS: THE D-BLOCK


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CONCEPT: PROPERTIES OF TRANSITION METALS Like most main group metals the transition metals possess similar physical properties such as: luster , high

densities , good electrical and thermal conduction , _________ melting points and hardness.

Conduction

• _________ possesses the greatest electrical conduction with _________ coming in second.

Melting Point

• _________ possesses the highest melting point at 3400 oC, while _________ is the only metal that is a liquid

at room temperature.

Hardness

• Metals such as _________ and _________ are strong, but other metals such as _________ , _________

and _________ are considered soft.

Oxidation States

• Remember that transition metals possess variable charges and so the use of

__________________________ is necessary to identify the correct charge of the element.





CONCEPT: ELECTRON CONFIGURATION OF TRANSITION METALS Electron configurations are a representation of how electrons are distributed into orbitals and furthermore how those orbitals fit into different energy levels.

Recall there are exceptions that exist with the electron configuration of transition metals that end with _______ or _______.

• These exceptions are normally observed with only the _____________________ transition metals.

EXAMPLE: Write the condensed configuration for the following element:

Chromium (24 electrons)

1s2s$$$$$$2p3s$$$$$$3p$$$$$$3d4s$$$$$$4p$$$$$$4d$$$$$4f5s$$$$$$5p$$$$$$5d$$$$$5f$$$$$5g6s$$$$$$6p$$$$$$6d$$$$$6f$$$$$6g$$$$6h$7s$$$$$$7p$$$$$$7d$$$$$7f$$$$$7g$$$$7h





PRACTICE: ELECTRON CONFIGURATION OF TRANSITION METALS

EXAMPLE: Write the condensed electron configuration of the following element.

W

EXAMPLE: Write the condensed electron configuration of the following ion.

Mn4+

EXAMPLE: Write the condensed electron configuration of the following ion.

Fe6+





CONCEPT: TRENDS IN ATOMIC SIZE

Recall it ____________ going from left to right across a period and ______________ going down a group.

• For transition metals, when moving across a period the size for the most part remains constant.

Ex:

• When moving from the 3d transition metals to the 4d transition metals we see a (n) _____________ in atomic radius, but moving from 4d to 5d the pattern is quite different and stays relatively constant.

• This phenomenon is referred to as the lanthanide contraction.

e-

e-

e- e-

e-

e- e-

e-

e-

e-




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CONCEPT: TRENDS IN IONIZATION ENERGY

Recall it ____________ going from left to right across a period and ______________ going down a group.

• However when moving down a group we see that the ionization energy of the third row is higher than the first and second row.

• This opposite trend is the result of the element’s _______________________________________.

e-

e-

e- e-

e-

e- e-

e-

e-

e-





CONCEPT: TRENDS IN OXIDATION STATES

As we’ve stated in the past oxidation involves _______________ electrons. • One of the most common features of transition metals is that they possess multiple oxidation states.

Ox. State

3B (3) 4B (4) 5B (5) 6B (6) 7B (7) 8B (8) 8B (9) 8B (10) 1B (11) 2B (12)

Sc Ti V Cr Mn Fe Co Ni Cu Zn 0 ! ! ! ! ! ! ! ! ! !

+1 ! ! ! ! ! ! ! +2 ! ! ! ! ! ! ! ! ! +3 ! ! ! ! ! ! ! ! !

+4 ! ! ! ! ! ! !

+5 ! ! ! !

+6 ! ! !

+7 !





PRACTICE: TRENDS IN OXIDATION STATES

EXAMPLE 1: Determine the oxidation state of the underlined element: [Ni(H2O)6]Cl2.

EXAMPLE 2: Determine the oxidation state of the underlined element: [Co(NH3)4(H2O)Br]Br2.

PRACTICE 1: in which compound does Ti exhibit greater metallic behavior: TiF2 or TiF6? PRACTICE 2: Which oxide, CrO3 or CrO, forms a more acidic aqueous solution?





CONCEPT: COORDINATION COMPOUNDS The most prevalent aspect of transition metal chemistry is the formation of coordination compounds or complexes.

• Within a coordination complex there is at least one _______________________, a species that is made of a metal cation that is connected to molecules and/or anions called _______________.

• In order to maintain the overall neutrality of the compound, _______________________ are used.

[Ni(NH3)4]Cl2

_______________________ _______________________ _______________________

COORDINATION NUMBERS The coordination number is the number of ligand atoms bonded to the central metal cation.

• The coordination number is based on the _________ of metal cation and its _________________________.

• The most common coordination number is _____ , however 2 and 4 are also common.

GEOMETRIES The types of geometries allowed are based on the coordination number of the central metal ion.

Coordination Number Shape

Linear

Tetrahedral

Square Planar




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PRACTICE: COORDINATION COMPOUNDS EXAMPLE 1: Determine the geometry for the following complex ion: [Zn(NH3)4]2+. EXAMPLE 2: Determine the geometry for the following complex ion: [AuBr2] –. PRACTICE: Determine the geometry for the following complex ion: [Cr(H2O)4Cl2]2+.





CONCEPT: LIGANDS A ligand can be thought of as a ___________________ because it bonds to a central metal cation in a complex ion by using its lone pair.

Ligands are characterized by the number of elements in the molecule that can donate a lone pair.

• These compounds use this lone pairs to grab onto these metal cations and are referred to as ____________________ agents.

Ligands that possess only _____ element(s) able to donate a lone pair are referred to as monodentate ligands.

Monodentate ("One-toothed")

H2O X C N HO

NH3 S C N

or

O N O

or

Ligands that possess ________________ element(s) able to donate a lone pair are referred to as bidentate ligands.

Bidentate ("Two-toothed")

CO

OC

O

O

2 –

H2C CH2H2N NH2





Ligands that possess ______________ element(s) able to donate a lone pair are referred to as polydentate ligands.

Polydentate ("Many-toothed")

CH2 CH2

NH NH2

H2C

H2CN N

CH2

CH2

H2C

H2C

C

C

C

C

O

O

O

O

O

O

O

O

4 –

Ethylenediaminetetraacetate ion (EDTA4 –)

H2CH2C

H2N

Diethylenetriamine

P

O

O O

O

P

O

O

O

P

O

O

O

5 –

Triphosphate ion





CONCEPT: NAMING COORDINATION COMPOUNDS In the early days of coordination compounds they were named after the person who first prepared them or for their brilliant colors.

• Today, the naming of coordination compounds is similar to naming ionic compounds.

Rules for Naming: a. The metal cation is written before the nonmetal anion.

• If the metal is a transition metal we must use a _________________________ to describe its positive charge. • The only space in the name should appear between the ___________ and the ___________ .

b. For the complex ion portion, neutral ligands are written before anionic ligands, and the formula for the whole ion is placed inside of brackets.

• Within the complex ion, the ligands are named in alphabetical order before the metal ion.

• Anionic ligands drop the – ide and add – o after the root name.

Neutral Ligands Anionic Ligands

Formula

Name \

Formula

Name

H2O Aqua Fluoro F – NH3 Ammine Chloro Cl – CO Carbonyl Bromo Br – NO Nitrosyl Iodo I – Hydroxo OH – Cyano CN –




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c. A numerical prefix is used to determine the number of a particular ligand.

d. If the complex ion is an anion, we replace the ending of the metal name and add – ate.

Complex Anions

Metal

Name in Anion

Iron (Fe) Ferrate Copper (Cu) Cuprate Lead (Pb) Plumbate Silver (Ag) Argentate Gold (Au) Aurate Tin (Sn) Stannate

[Ti(NH3)4Br2]Br ________________________________________

Na[Pt(NH3)3Cl5] ________________________________________





PRACTICE: NAMING COORDINATION COMPOUNDS EXAMPLE 1: Give the systematic name for the following formula: Na[Ag(CN)2]. EXAMPLE 2: Give the formula based on the given name: Tetraamminezinc carbonate PRACTICE 1: Give the systematic name for the following formula: [Co(NH3)4(H2O)Cl]Cl2. PRACTICE 2: Give the formula based on the given name: Lithium bis(thiosulfato)argentate (I)





CONCEPT: STRUCTURE AND ISOMERISM Isomers are compounds that possess the same molecular formula, which means they have the same atoms, but they differ in their location of each atom.

ISOMERS_______ formula but ____________ structures

Different Connections

CONSTITUTIONAL ISOMERS

Coordination Isomers

Ligand and/or counterion swapping

Linkage Isomers

Difference in Donor Atom

Without Stereocenters

GEOMETRIC ISOMERS

Cis/Trans Isomers (Diastereomers)

Different spatial arrangements

With Stereocenters

OPTICAL ISOMERS

Enantiomers

Nonsuperimposable mirror images

Same Connections

STEREOISOMERS




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CONCEPT: NUCLEAR REACTIONS

Nuclear Reactions deal with chemical processes in _______________ nuclei atoms.

Unlike normal chemical reactions where the identity of the elements stay the same, nuclear reactions often

result in elements changing into _______________ elements.

Early studies of radioactive nuclei by the British physicist Ernest Rutherford in 1897 shows that there are three

common types of radiation and nuclear reactions:

• _______________________________

• _______________________________

• _______________________________


CH.17 - NUCLEAR CHEMISTRY


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CONCEPT: ALPHA (α) DECAY

Alpha decay occurs when an unstable nucleus emits a particle composed of _________ protons and _________

neutrons.

An alpha particle can be represented by ______________ or ______________.

In terms of the size of radioactive particles, alpha particles are the ___________________.

• It is the most damaging to biological cells because it has the ____________ ionizing power.

• Has the ____________ penetrating power and can be stopped by clothing and the air of our

environment.

EXAMPLE: Write balanced nuclear equations for each of the following alpha emissions. a) Curium (Cm) – 248

b) Bismuth (Bi) – 207





CONCEPT: BETA (β) DECAY

Beta (β) decay occurs when an unstable nucleus emits a(n) _______________________.

• A beta particle can be represented by ______________ Beta decay can be represented by the following reaction: In terms of the size of radioactive particles, beta particles are ________________ than alpha particles.

• It is the not as damaging to biological cells and so has a ____________ ionizing power.

• ____________ penetrating power and can only be stopped by a sheet of metal or a large block of wood.

EXAMPLE 1: Write balanced nuclear equations for each of the following beta emissions. a) Magnesium (Mg) – 25

b) Ruthenium (Ru) – 102

EXAMPLE 2: Pb – 208 is formed from Th -232. How many alpha and beta decays have occurred?

a. 6, 2 b. 6, 6 c. 6, 4 d. 4, 6 e. 8, 2





CONCEPT: GAMMA RAY EMISSION

Gamma radiation is related to the electromagnetic spectrum. Gamma rays have the highest energy and therefore they have _________________ wavelength and _________________ frequency. A gamma particle can be represented by ______________.

• It causes no change in the atomic mass or atomic number and usually happens with alpha or beta decay.

• Gamma particles have the _____________ ionizing power.

• Gamma particles have the _____________ penetrating power so thick layers of lead shielding are needed.

EXAMPLE: Which of the following represents an element that has experienced a gamma emission?

a) Cl: 1s22s22p63s23p5

b) Be: 1s22s2

c) Na: 1s22s22p63p1

d) N: 1s22s22p3





CONCEPT: ELECTRON CAPTURE Electron capture involves the absorption of an electron by an unstable nucleus and is represented by the following reaction:

EXAMPLE: Write balanced nuclear equations for each of the following elements after undergoing electron capture. a) Rutherfordium (Rf) – 263

b) Nobelium (No) – 260

c) Lead (Pb) – 207





CONCEPT: POSITRON EMISSION Positron emission occurs when an unstable nucleus emits a positron. The positron is the antiparticle of the electron,

with the same mass as an electron, but with the opposite sign.

A positron particle can be represented by ______________ and in the overall reaction a proton is converted into a

neutron and emits a positron.

EXAMPLE 1: Write balanced nuclear equations for each of the following positron emissions. a) Uranium (U) – 235

b) Radon (Rn) – 222

EXAMPLE 2: A nuclide of Th – 225 undergoes 3 alpha decays, 4 beta decays and a gamma emission. What is the product?

a. Radium b. Radon c. Actinium d. Cadmium e. Antimony





CONCEPT: RADIOACTIVE DECAY RATES Radioactive decay is kinetically a first – order process, whose rate is proportional to the number of radioactive nuclei

N, where k is a first – order rate constant called the _____________ constant.

Decay Rate = k x N

Radioactive decay rates follow a _________ order process so we can reuse the _________ order Integrated Rate Law

from the Chapter on ___________________________.

EXAMPLE: A sample of radon-222 has an initial α particle activity (A0) of 8.5 x104 dps (disintegrations per

second). After 7.3 days, its activity (A) is 3.7 x 104 dps. What is the half-life of radon-222?




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CONCEPT: RADIOACTIVE DECAY RATES (CALCULATIONS)

EXAMPLE 1: Gallium citrate, containing the radioactive isotope gallium – 67, is used medically as a tumor seeking

agent. It has a half – life of 78.2 hours. How long will it take for a sample of gallium citrate to decay to 20.0% of its

original activity?

EXAMPLE 2: What percentage of carbon – 14 ( t1/2 = 5715 years) remains in a sample estimated to be 16,230 years old?




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TOPIC: ORGANIC CHEMISTRY

Organic Chemistry is the study of carbon and the other common nonmetals it is connected to:

____ , ____ , ____ & ____ .

Some organic molecules are made of just carbons and hydrogens and are aptly named _____________________.

C

H

H

H

C

H

H

C

H

H

H

A B

CC

CC

O

H

H

H H H

H

H

H

C

H

H

H

N H

H

C D

H – Cl


CH.20 - ORGANIC CHEMISTRYCH.18 - ORGANIC CHEMISTRY I: THE HYDROCARBONS


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TOPIC: STRUCTURAL FORMULAS

Your typical organic molecule can be drawn in a few different ways.

For example if you take a look at pentane, C5H12.

C

H

H

H

C C C C

H

H

H

H

H

H

H

H

H

Structural Formula Condensed Formula

CH3–CH2–CH2–CH2–CH3

Skeletal Formula

EXAMPLE 1: Determine the molecular formula for the following organic compound.

EXAMPLE 2: Convert the following structural formula into its condensed formula.

C

H

H

H

C C C C

H

H

CH3

H

H

H

C

H

H

H




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PRACTICE: STRUCTURAL FORMULAS 1



EXAMPLE 3: Convert the following structural formula into the carbon skeleton formula, otherwise known as

a kekulé structure.

C

O

HO C O C C

H

H

C

H

N

H H

H

H

H

H

H




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TOPIC: CHIRALITY

A molecule is chiral when it possesses a carbon that has ________ unique groups attached to it.

This molecule is said to be non-superimposable.

The mirror image of any chiral molecule is called an _____________________ and together they

are referred to as ____________ isomers or _____________ isomers.

Br

C

HOCH3

H

mirror

Br

C

OHH3C

H

EXAMPLE 1: Identify the compound that possesses an asymmetric center.

Cl C

H

H

H

A B C D

C C

CH3

C

CH3

HO–CH2 H C C

F

H

H

H

Cl

H

H3C

Cl




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PRACTICE: CHIRALITY 1

EXAMPLE 1: From the previous question draw the mirror image of the chiral molecule.

EXAMPLE 2: Draw the mirror images for the following molecule.

Mirror Method Inversion Method

OHCl

PRACTICE: Draw the mirror images for the following molecule.

Mirror Method Inversion Method




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TOPIC: OPTICAL ACTIVITY

We recently learned that a chiral molecule is one that possesses a carbon connected to four distinct groups.

One key feature of chiral molecules is that they rotate plane polarized light.

Clockwise rotation = dextrorotatory (d) or (+) Counterclockwise rotation = levorotatory (l) or (–)

Chiral Solution

Chiral Solution

The names and degrees of rotation have nothing to do with the chirality of the compounds.

EXAMPLE: Which of the following compounds would be optically active?

A B C D

Cl

Br

OHF




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TOPIC: HYDROCARBONS

The term of “hydrocarbon” refers to compounds that contain only carbons and hydrogens. These compounds may

possess single, double or triple bonds.

Carbons are ___________________ meaning that when they are neutral they must make 4 bonds.

C

H

H

H

C

H

H

H

Alkane

C

H

H

C

H

H

Alkene

CH C H

Alkyne

C

CC

C

CC

H

H

H

H

H

H

Benzene

Alkane Prefixes The name of alkanes is based on the number of carbons in the compound. These alkane “prefixes” must be

memorized in order to name more complex structures later on in the chapter.

# of Carbons

Alkane Prefix # of

Carbons

Alkane Prefix

1

6

2

7

3

8

4

9

5

10




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Alkanes Alkanes are hydrocarbons that contain only single bonds. They are sometimes referred to simply as “ saturated”

hydrocarbons.

Since all of the carbons are connected to 4 electron groups they all have an _______ hybridization.

Alkanes have a generic formula of _____________.

C

H

H

H

C

H

H

C

Propane

C

H

H

H

C

H

H

H

Ethane

H

H

H

EXAMPLE 1: Determine the formula and name of a hydrocarbon that contains only single bonds and 8 carbons. There is a direct relationship between an alkane’s weight and it’s measured boiling point.

Generally, the _____________ the molecular weight of an alkane the higher its boiling point.

Number of Carbons Physical State Usage

Up to 4

Gas

Cooking and Heating

5 to 7

Volatile Liquids

Gasoline and solvents

6 to 18

Liquids

Gasoline

12 to 24

Liquids

Jet Fuel

16 to 50

High Boiling Liquids

Diesel, heating oils and grease

Over 50

Solids

Petroleum jelly, waxes





Alkenes Alkenes are hydrocarbons that contain at least one double bond. They are sometimes referred to simply as

“unsaturated” hydrocarbons.

The double bonded carbons are connected to 3 electron groups and so have an _______ hybridization.

Alkenes and cycloalkanes have a generic formula of _____________.

CH

H

C

H

C

Propene

C

H

H

C

H

H

Ethene

H

H

H

Cyclohexane

EXAMPLE 2: Determine the formula and simplest name of a hydrocarbon that contains one double bond and 6 carbons.





Alkynes Alkynes are hydrocarbons that contain at least one triple bond. They are sometimes referred to simply as

“unsaturated” hydrocarbons just like the alkenes.

The triple bonded carbons are connected to 2 electron groups and so have an _______ hybridization.

Alkynes have a generic formula of _____________.

CH C C

Propyne

CH C H

Ethyne

H

H

H

EXAMPLE 3: Determine the formula and simplest name of a hydrocarbon that contains one triple bond and 9 carbons.





Aromatics Aromatic compounds have a benzene ring (C6H6) as their major defining feature.

C

CC

C

CC H

H

H

H

H

H

C

CC

C

CC H

H

H

H

H

H

Resonance Structures

C

CC

C

CC H

H

H

H

H

H

Resonance Hybrid

Benzene rings can be found in many everyday compounds from fossil fuels to even some vital medications.





TOPIC: ALKYL GROUPS

The removal of a hydrogen atom from an alkane compound creates an alkyl group. Remembering these alkyl groups

are the first step in learning how to name organic compounds.

H C

H

H

H

Methane

C

H

H

H

Methyl

When you have a two carbon chain then the alkyl name will be ______________.

CH3 – CH3

________________

CH2CH3

________________




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When you have a three carbon chain then the alkyl name will be __________________ or __________________.

________________

CH3–CH2–CH3

________________

CH2CH2CH3 CH3CHCH3

________________

When you have four carbons then the alkyl names will be ____________________ , ____________________ ,

_____________________ or _____________________ .

________________

CH3–CH2–CH2–CH3

________________

CH2CH2CH2CH3 CH3CHCH2CH3

________________





________________

CH3–CH–CH3

________________ ________________

CH3

CH2–CH–CH3

CH3 CH3–C

CH3

CH3

EXAMPLE: Based on your knowledge of chemical structures draw the three isomers of C5H12 (pentane).





TOPIC: NAMING ALKANES When naming any alkane compound there is a set of rules you should follow to get the correct answer.

1. Find the longest carbon chain and assign a root name accordingly.

If there is a tie between longest chains, choose the chain that gives _____________ substituents.

Substituents are the groups that branch off the main chain and weren’t counted as part of the main chain.

2. Number the chain from the end closest to a substituent.

3. Substituents will be named alphabetically and a number will accompany the location of their attachment.

4. If more than one of a similar alkyl group is present then we must use numerical prefixes to describe the

number of them.

__________ – 2 __________ – 3 __________ – 4 __________ – 5




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PRACTICE: NAMING ALKANES 1

EXAMPLE 1: Naming the following alkane compound.

a) EXAMPLE 2: Naming the following alkane compound.

b) EXAMPLE 3: Naming the following alkane compound.

c )





PRACTICE: NAMING ALKANES 2 EXAMPLE 4: Determine the structure for the following alkane compound.

3, 4, 5-trimethyldecane EXAMPLE 5: Determine the structure for the following alkane compound.

4-tert-butyloctane EXAMPLE 6: Determine the structure for the following alkane compound.

1-bromo-2,3-dichlorocyloheptane





TOPIC: NAMING ALKENES Naming an alkene is similar to naming an alkane with a few differences.

1. Find the longest carbon chain and change the “ane” ending to “ene”.


2. Number the chain from the end closest to the double bond and provide a number for its location.



number of them.

__________ – 2 __________ – 3 __________ – 4 __________ – 5

EXAMPLE 1: Naming the following alkene compound.




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PRACTICE: NAMING ALKENES 1 EXAMPLE 2: Naming the following alkene compound.

EXAMPLE 3: Determine the structure for the following alkene compound.

5-bromo-4,4-dimethyl-2-heptene EXAMPLE 4: Determine the structure for the following alkene compound.

2, 5–octadiene EXAMPLE 5: Determine the structure for the following alkene compound.

Trans–3–hexene




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TOPIC: NAMING ALKYNES Naming an alkyne is similar to naming an alkane with a few differences.

1. Find the longest carbon chain and change the “ane” ending to “yne”.


2. Number the chain from the end closest to the triple bond and provide a number for its location.



number of them.

__________ – 2 __________ – 3 __________ – 4 __________ – 5

EXAMPLE 1: Naming the following alkyne compound.

EXAMPLE 2: Provide the structural formula from the following alkyne name. 4–phenyl–2–nonyne




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TOPIC: ALKANE REACTIONS Alkanes are also referred to as _______________, which is derived from Latin meaning “little affinity”.

This explains their very low reactivity.

Alkanes in general undergo only two main types of reactions.

a) _______________________________

b) _______________________________




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TOPIC: FREE RADICAL CHAIN REACTION

As previously stated, alkanes are almost completely unreactive.

And one of the two reactions it undergoes is free-radical halogenation.

Under this reaction a ________________ on a carbon is replaced by a ________________.

The Radical Chain Reaction

Alkanes will react with diatomic halogens in the presence of __________________________________.

Initiation: (A stable compound undergoes homolytic cleavage: 0 free radicals 2 free radicals)

Propagation: (The radical reacts with a stable compound to create a new radical and new stable compound.)

Termination: (Two radicals join to create a stable compound: 2 free radicals 0 free radicals)





PRACTICE: FREE RADICAL CHAIN REACTION 1

EXAMPLE 1: Determine the major product(s) of the following reaction.

CH3CH3 + Br2

Heat or Light

EXAMPLE 2: Determine the major product(s) of the following reaction.

CH3CH2CH3 + Cl2Heat or Light





TOPIC: HALOGENATION Alkenes and alkynes undergo addition reactions in which elements add across their

_______ bonds to create new ___________ bonds.

• Under Halogenation, ______ halogens are added to ______ pi bond.

EXAMPLE 1: Determine the major product from the following Halogenation reaction.

CH3CH=CH2 + Cl2

EXAMPLE 2: Determine the major product from the following Halogenation reaction.

+ 2 moles Br2




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TOPIC: HYDROGENATION Hydrogenation can be seen as a reduction reaction in which _____ hydrogens are added to _____ pi bond.

• In General Chemistry reduction refers to gaining _________ , but in Organic Chemistry reduction is thought of as gaining ________________.

EXAMPLE 1: Determine the major product from the following Hydrogenation reaction.

+ H2Catalyst

EXAMPLE 2: How many moles of hydrogen would be need to complete reduce the following

compound?

C CH

CH3

C

CH3





TOPIC: HYDROHALOGENATION Under Hydrohalogenation, a _________________ and a _________________ are added across a pi bond.

1) Follows Markovnikov’s Rule

a) The _____________________ goes to the double bonded carbon with _______ hydrogens.

b) The _____________________ goes to the double bonded carbon with _______ hydrogens.

EXAMPLE 1: Determine the major product from the following Hydrohalogenation reaction.

C CH

CH3

CH2CH2CH3

CH2CH3+ HCl

EXAMPLE 2: Determine the major product from the following Hydrohalogenation reaction.

+ 2 moles HBrH3C C CH





TOPIC: AROMATIC REACTIONS Although benzene rings have pi bonds like alkenes and alkynes, they cannot undergo addition reactions because they are much too stable.

• We often refer to benzene as an aromatic compound in order to describe its high level of stability.

Instead of doing addition reactions benzene rings will under ________________ reactions in order to

maintain their aromaticity. The two major types of Aromatic reactions we will investigate will be:

1. _______________________________ – where a halogen is placed on benzene.

H

H

H

H

H

H

FeBr3+ Br2

H

H

H

H

H

H

FeCl3+ Cl2




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2. _______________________________ – where an alkyl group is placed on benzene.

H

H

H

H

H

H

AlCl3+ CH3Cl

EXAMPLE 1: Determine the major product from the following Aromatic reaction.

H

H

H

H

H

H

AlBr3+Br




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TOPIC: NOMENCLATURE OF FUNCTIONAL GROUPS Naming a compound with a functional group is similar to previous examples of naming we’ve done.

1. Find the longest carbon chain and change the “ane” ending based on the functional group present.

• If there is a tie between longest chains, choose the chain that gives _____________ substituents.

2. Number the chain from the end closest to the functional group present and provide a number for its location.

• Some functional groups don’t need their location numbered because they are terminal.


4. If more than one of a similar alkyl group is present then we must use numerical prefixes to describe the number

of them. __________ – 2 __________ – 3 __________ – 4 __________ – 5

Class Suffix Name Carboxylic Acid

-oic acid

Ester

-oate

Aldehyde

-al

Ketone

-one

Alcohol

-ol

Amine

-amine

Alkene

-ene

Alkyne

-yne

Alkane

-ane

Ether

ether

Alkyl Halide

---




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PRACTICE: NOMENCLATURE OF FUNCTIONAL GROUPS 1 EXAMPLE 1: Provide the name of the following alcohols and alkyl halides and state if they are primary, secondary or tertiary.

a) BrCH2CH2CH2OH b)

CH3CH2CH2CH2CCH3

CH2CH3

OH

c) CH3CH2CH2CHCH3

Cl

EXAMPLE 2: Provide the name of the following functionalized organic compounds.

a) CH3COH

O

b) CH3CH2CH2CH2CH

O

c) CH3CH2CH2CH2CCH2CH3

O





PRACTICE: NOMENCLATURE OF FUNCTIONAL GROUPS 2

EXAMPLE 3: Provide the name of the following functionalized organic compound.

EXAMPLE 4: Provide the name of the following amines and state if they are primary, secondary or tertiary.

a) CH3CH2CH2CHCH2CH3

NH2

b) CH3CH2CH2NHCH2CH3

EXAMPLE 5: Provide the name of the following ethers

a) CH3CH2CH2OCH3 b) CH3CH2OCH2CH3





PRACTICE: NOMENCLATURE OF FUNCTIONAL GROUPS 3

EXAMPLE 6: Provide the structure of the following aromatic hydrocarbons. p-ethyltoluene EXAMPLE 7: Provide the structure of the following aromatic hydrocarbons. m-bromobenzoic acid EXAMPLE 8: Provide the structure of the following aromatic hydrocarbons. 2,4,6-trimethylnitrobenzene





TOPIC: ALCOHOL REACTIONS Alcohols can be easily identified by the presence of the hydroxyl group ________ .

• Because of the presence of this highly polarized group alcohols can be involved in many organic reactions.

1. Substitution Reactions – under this reaction alcohols react with the hydrohalic acids of ______ or ______ to become alkyl halides.

CH3CH2CHCH2CH2CH3

OH

+ HBr

2. Elimination Reactions – under this reaction under the presence of concentrated acids such as __________ or

__________ an alcohol undergoes dehydration to form an alkene.

CH3CCH3

OH

+ H2SO4

CH3




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3. Oxidation Reactions – An alcohol can be seen as an alkane that has undergone an initial oxidation. The

alcohol can then continue to be oxidized to form either a(n) _________________ or _________________.

• This is done with the strong oxidizing agent sodium dichromate, Na2Cr2O7 and the strong acid sulfuric

acid, H2SO4.

• There is also an inverse relationship that can change our products back into their reactants.

CH3CCH3

OH

CH3 H2SO4

Na2CrO7

CH3CH2CHCH3

OH

H2SO4

Na2CrO7

CH3CH2CH2 OH H2SO4

Na2CrO7




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TOPIC: CARBOXYLIC ACID DERIVATIVE REACTIONS Esters and amines are involved in reactions that deal with carboxylic acids.

1. Esterification – Under this reaction a carboxylic acid reacts with an alcohol and undergoes a dehydration reaction

in order to form an ester.

+ CH3CH2CH2OHHeatC

O

OH

2. Amide Formation – Under this reaction a carboxylic acid reacts with an amine and undergoes a dehydration to

reaction in order to form an amide.

+ CH3CH2NHCH3Heat

CH3CH2C

O

OH




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