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Dr. Chirie’s Stoichiometry-chemistry Notes 2-5-19 Name:_____________________ Period: __ (Notes Due: 3/1/19) Pts:___/100

UNIT: Stoichiometry Stoichiometry describes the proportions between atoms in molecules and reactants/products in chemical reactions. You can use this information to balance chemical equations. Topics:

1. Types of chemical reactions 2. Balancing chemical equations 3. Mole concept: Mole and Molarity

a. Number of particles and mass connection b. Molar mass (formula mass) of compounds and atoms c. Molarity calculations

4. Chemical formula vs. Empirical formula of a substance a. Formula weight (mass) = molar mass b. Empirical formula weight (mass)

5. Percent composition compounds 6. Mole relationships in Balanced equations 7. Mass relationships in Balanced equations 8. Limiting reactants and theoretical yield

____________________________________________________________________________________________ BASIC INFO ON CHEMICAL EQUATIONS

• Reactants: Substances that exist BEFORE the chemical reaction • Products: Substances produced AFTER the chemical reaction • Arrow: points to the DIRECTION of products. Arrow with a triangle = product are heated • Reversible reactions: have arrows pointing to both directions. • Plus signs used to separate different compounds • Superscripts with charges are used to state charges (oxidation states) of ions • Subscripts are used to state the number of atoms present • Coefficients placed before a formula of a compound indicated the number of MOLECULES or MOLES present.

The only part you change when balancing a chemical equation.

TOPIC-1: TYPES OF CHEMICAL REACTIONS (5 types)

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1. Combination (Synthesis) Reactions: Two or more reactants form one product in a combination reaction. An example of a combination reaction is the formation of sulfur dioxide when sulfur is burned in air:

S (s) + O2 (g) → SO2 (g) *See figure 1b

2. Decomposition Reactions: In a decomposition reaction, a compound breaks down into two or more substances. Decomposition usually results from electrolysis or heating. An example of a decomposition reaction is the breakdown of mercury(II) oxide into its component elements.

2HgO (s) + heat → 2Hg (l) + O2 (g) *See figure 1b

3. Single Displacement Reactions: A single displacement reaction is characterized by an atom or ion of a single compound replacing an atom of another element. An example of a single displacement reaction is the displacement of copper ions in a copper sulfate solution by zinc metal, forming zinc sulfate:

Zn (s) + CuSO4 (aq) → Cu (s) + ZnSO4 (aq) *See figure 1b

Single displacement reactions are often subdivided into more specific categories (e.g., redox reactions).

4. Double Displacement Reactions: In this type of reaction, elements from two compounds displace each other to form new compounds. Double displacement reactions may occur when one product is removed from the solution as a gas or precipitate or when two species combine to form a weak electrolyte that remains undissociated in solution. An example of a double displacement reaction occurs when solutions of calcium chloride and silver nitrate are reacted to form insoluble silver chloride in a solution of calcium nitrate.

CaCl2 (aq) + 2 AgNO3 (aq) → Ca(NO3)2 (aq) + 2 AgCl (s) *See figure 1b

A neutralization reaction is a specific type of double displacement reaction that occurs when an acid reacts with a base, producing a solution of salt and water. An example of a neutralization reaction is the reaction of hydrochloric acid and sodium hydroxide to form sodium chloride and water:

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

5. Combustion (burning) Reactions: In this type of reaction, an element or compound combines with O2 gas to produce Oxides. When compounds with H and C are burned, CO2 and H2O are ALWYS produced.

2Mg + O2 → 2MgO CH4 + 4O2 → CO2 + 2H2O _______________________________________________________________________________________________ TOPIC- 2: Balancing Chemical Equations (3-steps):

1. Write the unbalanced equation. • Chemical formulas of reactants are listed on the left-hand side of the equation. • Products are listed on the righthand side of the equation. • Reactants and products are separated by and between them to show the direction of the reaction. Arrow

points at products. (Reactions at equilibrium will have arrows facing both directions.) • Use the one- and two-letter element symbols to identify elements. • When writing a compound symbol, the cation in the compound (positive charge) is listed first, before the

anion (negative charge). For example, table salt is written as NaCl and not ClNa

2. Balance the equation. • Apply the Law of Mass Conservation to get the same number of atoms of every element on each side of

the equation. Tip: Start by balancing an element that appears in only one reactant and product • Once one element is balanced, proceed to balance another, and another until all elements are balanced • Balance chemical formulas by placing coefficients in front of them. Do not add subscripts, because this will

change the formulas

3. Indicate the physical states of matter of the reactants and products. • Use (g) for gaseous substances • Use (s) for solids • Use (l) for liquids • Use (aq) for species soluble in water • Write the state of matter immediately following the formula of the substance it describes

Balancing Equation: Worked Example Problem : QUESTION: Tin(IV) oxide is heated with hydrogen gas (H2) to form tin metal and water vapor. Write the balanced equation that describes this reaction. Step-1. Write the unbalanced equation:

SnO2 + H2 → Sn + H2O __________ equation-A

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Stoichiometry- chemistry notes Name:____________________________ Period: ___ Date: 2-5-2019

Step-2. Balance the equation: Look at the equation-A and see which elements are not balanced. In this case, there are two oxygen atoms on the left side and only one on the right side of the equation. Correct this by this by putting a coefficient of 2 in front of water:

SnO2 + H2 → Sn + 2H2O__________ equation-B This puts the hydrogen atoms out of balance: now there are two H-atoms on the left and four on the right. To get four hydrogen atoms on the right, add a coefficient of 2 for the hydrogen gas. Remember, coefficients are multipliers, so if we write 2H2O it means 2 x 2=4 hydrogen atoms and 2 x 1=2 Oxygen atoms.

SnO2 + 2 H2 → Sn + 2 H2O __________ equation-C The equation is now balanced. Be sure to double-check your math! Each side of the equation has 1 atom of Sn, 2 atoms of O, and 4 atoms of H.

**Refer to Table-5 and Table-6 in Nomenclature notes if you have trouble writing the chemical formulas of the products and reactants.

Step-3. Indicate the physical states of the reactants and products: Oxides are solids, hydrogen forms a diatomic gas, tin is a solid, and the term “water vapor” indicates that water is in the gas phase:

SnO2(s) + 2 H2(g) → Sn(s) + 2 H2O(g) ------ Final Balanced Completed equation

This is the balanced equation for the reaction. Be sure to check your work! Remember Conservation of Mass requires the equation to have the same number of atoms of each element on both sides of the equation. Multiply the coefficient (number in front) times the subscript (number below an element symbol) for each atom. For this equation, both sides of the equation contain:

1) 1 Sn atom 2) 2 O atoms 3) 4 H atoms

Home Work: HW-4 (Due 2-8-19) Name:__________________________ Period___ Score:______/100

Balancing equations Practice problems

Balance the reaction below Type of chemical reaction: 1. __Ca + _ O₂ → __ CaO Combination reaction

2. __ Mg + _NaOH → __ Mg(OH)2 + __ Na

3. __ Al(OH)3 + _HBr → __AlBr3 + __ H2O

4. __ MgO → _ Mg + __ O2

5. __ HNO3 + _Ca(OH)2 → __ Ca(NO3)2 + __ H2O

6. __ KNO₃ + _ H₂CO₃ → __ K₂CO₃ + __ HNO₃

7. __ AgCl + _ Na₂S → __ Ag₂S + __ NaCl

8. __ Ba₃N₂ + _ H₂O → __ Ba(OH)₂ + __ NH₃

9. __ FeS + _ O₂ → __ Fe₂O₃ + __ SO₂

10. __ Na₃PO₄ + _ HCl → __ NaCl + __ H₃PO₄

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Fill in the missing compounds and balance the equations: Type of chemical reaction: 11. _______ + ______→ __ TiO₂ + __ HCl

12. __ C₂H₆O + __ O₂ → ______ + __ H₂O

13. __ CaCl₂ + __ Na₃PO₄ → _________+ _________

14. ___CH4 + __ O2 → ________ + _________

15. __ AgBr + __ Li₂S → _______ + _________

________________________________________________________________________________________________ TOPIC- 3: THE MOLE CONCEPT: Mole and Molarity The Italian Physicist, Amedeo Avogadro (1776–1856) proposed that equal volumes of gases under the same conditions contain the same number of molecules, a hypothesis that proved useful in determining atomic and molecular weights and which led to the concept of the mole.

What is a MOLE?

• A mole is a counting-unit in chemistry. (a common counting unit = dozen = 12) • I mole has a value of = 6.02x1023 (this number is also called Avogadro’s number)

1 dozen = 12 of anything (ex: 12 eggs) similarly,

1 mole = 6.02x1023 of anything (ex: 6.02x1023 =602000000000000000000000 eggs) • The SI symbol for the unit is mol (moles) • In chemistry, 1 mole is defined as the number of carbon-12 atoms in 12g of carbon. • The relationship between mass and moles: Atomic mass in the periodic table is the mass of 1 mole of the atoms

listed. Example: o 1 mole of Hydrogen has 6.02 × 1023 atoms and has a mass of 1.0 grams o 1 mole of Carbon has 6.02 × 1023 atoms and has a mass of 12.0 grams. o 1 mole of Oxygen has 6.02 × 1023 atoms and has a mass of 15.99 grams. o 1 mol of NaCl has 6.02 × 1023 particles and has a mass of (Na=22.99g) + (Cl=35.45g) = 58.44 grams.

• 1 mole of a compound has 6.02 × 1023 particles (Ions/ atoms/molecules) and has a mass equal to the Formula mass of a compound

• The concept of the mole helps to put quantitative information about what happens in a chemical equation on a macroscopic level. For example, in the chemical reaction, 2H2O → O2 + 2H2, 2 moles of water are decomposed into 2 moles of hydrogen gas and 1 mole of oxygen gas. The mole can be used to determine the simplest formula of a compound and to calculate the quantities involved in chemical reactions.

What is MOLARITY?

• When dealing with reactions that take place in solutions, the related concept of molarity is useful. Molarity (M) is defined as the number of moles of a solute in 1 liter (1L) of any solution.

• SI unit = M (“molar”) • Use this formula for problems:

o Molarity = Moles of compound Liters

Examples: o A 1 molar (1M) solution of NaCl has the 1 mol (58.44g) of NaCl dissolved in a final volume of 1 liter of water. o A 1 molar (1M) solution of NaCl has 6.02 × 1023 particles of NaCl dissolved in a final volume of 1 liter of water.

o Q1: What is the molarity of a solution containing 116.88g of NaCl in 1L of water?

MOLAR MASS: Mass of one mole of any substance. This value is equal to the molecular formula of a given substance. (See “Formula mass’- topic below) Molar mass = Mass in grams moles

• Molar mass units = grams/ mole or g/mol

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Stoichiometry- chemistry notes (contd.) Name:____________________________ Period: __ Date: 2-12-2019

GASES: At a Temperature of 273K (0oC) and pressure of 1 atm, or Standard Temperature and Pressure (STP) one mole of any gas occupies 22.4 L of volume. (atm= atmosphere units = pressure unit)

PRACTICE PROBLEMS on MOLE and MOLARITY CONCEPT:

Q1: Sulfuric acid has the chemical formula of H2SO4.

i. What is the formula weight(mass) of sulfuric acid?_____________________________________________

ii. How may moles of any compound is found in a mass equal to its formula mass?____________________

iii. What is molar mass of sulfuric acid?_________________________________________________________

iv. Moles of sulfuric acid in 392g?___________________________________

v. If you add 392g of sulfuric acid to a final volume of 1 liter, what is the molarity of this solution?___________

vi. If you add 392g of sulfuric acid to a final volume of 2 liters, what is the molarity of this solution?__________

vii. Molarity value __________ when the number of moles of a solute in solution is increased.

PRACTICE PROBLEMS on GASES, VOLUME and MOLES CONCEPTs At STP,

1. the volume of 1 mole of H2 gas: _____________________________________________________

2. the volume of 1 mole of water vapor: __________________________________

3. Number of water molecules in 44.8L water vapor:__________

4. Mass of 44.8L water vapor:________________

5. Moles in 18g of water (H2O):_____________

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TOPIC-4: CHEMICAL FORMULA AND EMPIRICAL FORMULA OF COMPOUNDS What is a Molecular/Chemical Formula? This is the ________ Formula __________ of the compound. It tells you the actual number of each type of atom (element) present in the formula of a chemical. A. Formula Mass (weight) : • The formula mass of a compound is equal to the _________mass of _______________ in its molecular formula. • SI unit = grams (g). • In chemical bottles, formula Weight is abbreviated as, FW. • The formula mass is also equal to the Molar mass (mass of __________ of a compound). • The mass of 6.02x1023 particles of a chemical substance = formula weight = ____________

Ex. Calculate the formula mass (mass of 1 mol) of the following: 1) H2SO4 = (H x 2)+(Sx1)+(Ox4) = (1g x 2)+(32g x1)+(16g x4) = 2 + 32 + 64 = 98 g

2) N2 = (N x 2) = 14.0g x2 =28g/mol

3) H2O = (H x 2)+(O x1) = (1.0g x 2) +(15.99g x1) = 17.99g = 18g

4) H2O2 = (H x 2) + (O x2) = (1.0g x 2) +(16g x2) = 2 + 32 = 34g

5) NH4HCO3 = ______________________________________________________________________________

6) Al2(SO4)3 = ______________________________________________________________________________

7) C3H6O3 = ______________________________________________________________________________

8) C6H12O6 = ______________________________________________________________________________

9) C12H24O12 =______________________________________________________________________________

B. Empirical formula and Empirical formula Mass (weight): • The Empirical formula that gives the simplest whole-number ratio of atoms in a compound. • You __________ use the empirical formula to determine the identity of a chemical as some chemicals

can have the same empirical formula. • Ex. Compare the empirical formulas of C12H24O12 and C6H12O6. They both have the same empirical

formula of CH2O. • For some molecules or compounds, _______ the molecular formula and empirical formula are the same.

Ex. H2O, MgCl2, Na2CO3

Complete the table below: Molecular formula

Molar mass/formula mass Empirical formula

Empirical formula mass

Ca(OH)2 H2O H2O2 NH4HCO3 Ca3(PO4)2 C3H6O3 C6H12O6 C12H24O12

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Q1: Circle the correct answer. Which of the following tells you about the exact number and types of different atoms in a given compound or molecule? a. Atomic mass b. empirical formula c. molar mass d. molecular formula

Q2: Which of the following does NOT have the same empirical formula?

a. CH3CH2COOH b. CH2O c. C7H14O7 d. CH3COOH Q3: Find the number of moles in:

CONCEPTS: • 1mole = 6.02x 1023 particles. • At STP, the volume of 1mole of any gas is 22.4L

A) 6.02 X 1023 molecules of NO2 gas:__________________ B) 224 liters of NO2 gas at S.T.P.: __________________ C) 2.30 grams of NO2 gas: __________________ D) 3.01 X 1023 grains of rice: __________________ E) 56.0 liters of H2S gas at S.T.P. __________________ F) 136.4 grams of H2S gas__________________ Stoichiometry- chemistry notes (contd.) Name:____________________________ Period: __ Date: 2-12-2019

TOPIC- 5: PERCENT COMPOSITION:

Percent by mass of each element in a compound is called percent composition. Percent mass of an element in a compound = Total mass of all the atoms of that element x 100 Molar mass = formula mass of compound Example: Find the percent mass of Carbon, oxygen and hydrogen in C3H6O3

1. Percent mass of Carbon = (C x3) x 100 (C x 3)+(H x 6)+(O x 3)

= (12.0g x3) x 100 (12.0g x3)+(1.0g x6)+(15.99g x3)

= (12.0g)x3 x 100 (36.0g)+(6.0g)+(47.97g)

= 12.0 X3g x 100 89.97g

= 40 %

2. Percent mass of Oxygen = (O x3) x 100 (C x 3)+(H x 6)+(O x 3)

= (15.99 x3) x 100 (12.0g x3)+(1.0g x6)+(15.99g x3)

= (47.97g) x 100 (36.0g)+(6.0g)+(47.97g)

= 47.97g x 100 89.97g

= 0.533 x100 = 53.3 %

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Stoichiometry- chemistry notes (contd.) Name:____________________________ Period: __ Date: 2-12-2019

PRACTICE PROBLEMS ON Percent Composition CONCEPT: Q1: What is the percent composition of each element (type of atom) in H2SO4?

i. What is the formula mass of this compound?________________________________________________

ii. What is the percent composition of H?______________________________________________________

iii. What is the percent composition of S?______________________________________________________

iv. What is the percent composition of O?______________________________________________________

Q2: What is the percent composition of each element in (NH4)2CO3?

i. What is the formula mass of this compound?__________________________________________

ii. What is the percent composition of H?______________________________________________________

iii. What is the percent composition of N?______________________________________________________

iv. What is the percent composition of C?_____________________________________________________

v. What is the percent composition of O?______________________________________________________

Q3: What is the percent composition of each element in C6H12O6?

Formula mass =

H% =

C% =

O% =

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Stoichiometry- chemistry notes (contd.) Name:____________________________ Period: __ DUE Date: 2-13-2019

LAB_11: COMBUSTION LAB _______/50 Form Pts. Pre-Lab questions (Group work):

Recall our first lab (LAB1 we did the “Candle in the water” experiment. Below is the structure of a molecule of paraffine/candlewax. What do you think happens when you light a candle made up of paraffine and a wick? Discuss with your groups for 5 minutes.

Chemical structure of paraffine

a. Molecular formula of paraffine = C__H__

b. What type of chemical bond are indicated by the lines connecting atoms?_____________________

c. If you light up this candle what would be the TYPPE of chemical reaction that you will observe? ____________________

d. Write the SKELETON REACTION for the reactants and products of this reaction. (TIP: when matter is burned, OXYGEN is

ALWAYS one of the reactants) ______________________________________________________________

e. Calculate the formula mass of all the reactants and products for this chemical reactants: a. Paraffine: ___________________________________________

b. O2 : ___________________________________________

c. CO2 : ___________________________________________

d. H2O : ___________________________________________

f. What is the percent composition of H and C in Paraffine?

g. Write the balanced equation for this chemical reaction:_____________________________________________________

h. Watch Dr. Chirie’s Demonstration of this experiment in class. Draw the experimental set up below and label all parts:

_____________________________________ _________________________________

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Experiment after reaction Control experiment

i. Carefully compare the experimental-reaction and the Control, then write down 4 observations of what happened to the experimental reaction. Now match your observations with its relationship to the balanced chemical reaction in (d) in the table below:

Experimental OSERVATION Relationship to the balanced chemical reaction

1 Candle becomes smaller compared to the control

Paraffine is consumed (used up) in the reaction

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All the O2 gas inside the beaker is consumed in the reaction

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Water/H2O is produced

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CO2 is produced

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Energy is released into the environment when chemical bonds rearrange

Work to be done back in your assigned seat, alone and answer the questions below:

Today’s Chemical reaction: C_ H_ + O2 -> H2O + CO2

1. What are the reactants in this reaction:__________________________________________________________________

2. What are the products in this reaction:___________________________________________________________________

3. What is the relationship between molar mass and formula mass of a compound?________________________________

4. Calculate the molar mass of paraffine:___________________________________________________________________

5. Calculate the molar mass of Oxygen gas:_________________________________________________________________

6. Calculate the molar mass of Carbon dioxide gas:___________________________________________________________

7. Calculate the molar mass of Water:_____________________________________________________________________

8. What is the relationship between the COEFFICIENTS of a balanced equation and MOLES?__________________________

9. ______ mol C_ H_ combines with _______ mol of O2 gas and produces _____ mol of H2O and ________ mol of CO2

10. How may molecules of Paraffine are found in 3 moles?______________________________________________________

11. What is the relationship between the moles of a substance and its mass?_______________________________________

12. THINK TIME: If you needed to find the moles of paraffine USED UP in this reaction, what measurements will you take from

the candle?___________________________________________________________________________________