so what is chemistry? the study of matter matter pure substances elements na, h, o compounds...

So What is Chemistry?

The Study of Matter

Can be separated chemically

Can be separated physically

UNIT ILAB SKILLS and MEASUREMENT

UNIT IPerformance Indicators: You will:• will be able to respond appropriately to lab safety situations. • will be able to infer/select the correct unit to be used in a measurement. • will be able to distinguish between precision and accuracy. • will be able to calculate percentage error given laboratory and reference material data. • will be able to use factor label/dimensional analysis in problem solving. • will be able to transform quantities from standard notation to scientific notation.• will be able to use significant digits to express final calculations. Essential Questions1. What does a “safe” lab look like?2. How does one know how precise or accurate his/her results are?3. What are the most common measurements used in lab and how can one go from one to

another?

Measurements: SI system• The first standardized system of measurement, based on the decimal was

proposed in France about 1670. However, it was not until 1791 that such a system was developed.

• It was called the "metric" system, based on the French word for measure. The driving force was the growing importance of weights in the sciences, especially chemistry. At that time, every country had their own system of weights and measures. England had three different systems just within its own borders!!

• On May 20, 1875, delegates of 17 countries signed the Meter Convention. It was amended in 1921 and today 48 countries are signatories.

• The modern metric system has been renamed Systeme International d'Unites (International System of Units) and is denoted by the letters SI. SI was established in 1960, at the 11th General Conference on Weights and Measures. It was then that units, definitions, and symbols were revised and simplified.

There are three major parts to the metric system: :

• the seven base units • the prefixes and units built up from the base

units.

Here is a list of the base units which make up the metric system:

Physical Quantity Name of SI unit Symbol for SI unit

length metre (meter) m

mass kilogram kg

time second s

electric current ampere A

temperature Kelvin K

amount of substance mole mol

luminous intensity candela cd

FACTOR ...or in full ... or inwords

SI PREFIX

SI SYMB

OL

1,0E+241,0E+211,0E+181,0E+151,0E+121,0E+91,0E+61,0E+31,0E+21,0E+11,0E-11,0E-21,0E-31,0E-61,0E-9

1,0E-121,0E-151,0E-18 1,0E-211,0E-24

1 000 000 000 000 000 000 000 0001 000 000 000 000 000 000 0001 000 000 000 000 000 0001 000 000 000 000 0001 000 000 000 0001 000 000 000 1 000 0001 000100100,10,010,0010,000 0010,000 000 0010,000 000 000 0010,000 000 000 000 0010,000 000 000 000 000 0010,000 000 000 000 000 000 0010,000 000 000 000 000 000 000 001

septillionsextillion

quintillionquadrillion

trillionbillionmillion

thousandhundred

tententh

hundredththousandthmillionth billionth trillionth

quadrillionthquintillionthsextillionthseptillionth

yotta-zetta-exa-peta-tera-giga-

mega-kilo-

hecto-deca-deci-centi-milli-

micro-nano-pico-

femto-atto- zepto-yocto-

YZEPTGMkh

dadcmµnpfa zy

• In 1958, the International Committee on Weights and Measures added Mega-, Giga-, and Tera- to the multipliers and micro-, nano-, and pico- to the fractions. In 1960, at the 11th General Conference on Weights and Measures, everything was officially adopted.

• Since that time, additional prefixes have been added as the need arose. Typically, as scientific instruments get better and better, smaller and smaller quantities can be detected. So, new fractional prefixes need to be added. When they are, new multipliers are added also, to keep the system symmetrical.

Non-SI Units Commonly Used

• Liter: symbol = L. The SI unit for volume is m3 (cubic meter). One dm3 (cubic decimeter) equals one L. A cubic decimeter is a cube 0.1 m on a side.

• cubic centimeter: symbol = cm3. Often used for measuring the volume of solids, one cm3 equals one milliliter (mL).

Converting from One Metric Unit to Another

Skills you need to do this include: 1. Memorize the metric prefixes names and

symbols2. Determine which of two prefixes represents

a larger amount3. Determine the exponential "distance"

between two prefixes4. Significant figure rules5. Scientific notation

The Factor-Label Method

• The key skill in solving these problems is to construct a conversion factor.

• This conversion factor will make the old unit go away and create the new unit in its place.

• Along with this change, there will be a change in the value of the number.

What is a unit factor?

• It is a fraction that allows us to go from one unit to another unit using multiplication and division

• Unit factors are interchangeable.• 100 pennies = 1 dollar• 1 dollar = 100 pennies 100 pennies = 1 dollar____ 1 dollar 100 pennies

Factor label method

Steps:1. Identify starting & ending units.

2. Line up conversion factors so units cancel.

3. Multiply all top numbers & divide by each bottom number.

4. Check units & answer.

Multiply across the top and divide the answer

How many milliliters are in 1.00 quart of milk?

1.00 qt 1 L

1.057 qt= 946 mL

qt mL

1000 mL

1 L

Now.. Scientific Notation!!

Dealing with very small numbers: the mass of an electron= 0.00000000000000000000000000000091kg

Dealing with very large numbers

The distance to the sun= 93,000,000 miles

Scientific notation

Science must be able to manage very large and very small numbers.

Scientific notation is just a shorthand method to accomplish this task

Scientific notation

Numbers in scientific notation are made up of three parts: the coefficient, the base and the exponent.

5.67 x 105

Scientific notation

Numbers in scientific notation are made up of three parts: the coefficient, the base and the exponent.

5.67 x 105CoefficientExponent

Base

In order for a number to be in correct scientific notation, the following conditions must be true:

The coefficient must be: greater than or equal to 1 and less than 10.

The base must be 10.

The exponent must show the number of decimal places that the decimal needs to be moved to change the number to standard notation.

A negative exponent means that the decimal is moved to

the left when changing to standard notation. Positive to the right

FACTOR ...or in full ...1.0 x10+24

1.0 x10+21

1.0x10+18

1.0x10+15

1.0x10+12

1.0x10+9

1.0x10+6

1.0x10+3

1.0x10+2

1.0x10+1

1.0x10-1

1.0x10-2

1.0x10-3

1.0x10-6

1.0x10-9

1.0x10-12

1.0x10-15

1.0x10-18 1.0x10-21

1.0x10-24

1 000 000 000 000 000 000 000 0001 000 000 000 000 000 000 0001 000 000 000 000 000 0001 000 000 000 000 0001 000 000 000 0001 000 000 000 1 000 0001 000100100,10,010,0010,000 0010,000 000 0010,000 000 000 0010,000 000 000 000 0010,000 000 000 000 000 0010,000 000 000 000 000 000 0010,000 000 000 000 000 000 000 001

Scientific notation

You will be required to change numbers into scientific notation.

You will be required to change numbers back to standard notation.

You will be required to calculate with numbers in scientific notation.

Scientific notation

5.67 x 105

567000.

Between 1 and 9.9Indicates decimal

placement

Scientific notation

What about converting small numbers?0.0000025

Make this number larger than 1 and less than ten by moving the decimal.

0.0000025

2.5 x 10-6

Move to right is negative

Calculations

Calculators are very helpful tools, but unless you can do these calculations without them, you can never check to see if your answers make sense.

Any calculation should be checked using your logic, so don't just assume an answer is correct

Multiplication

Rule for Multiplication - When you multiply numbers with scientific notation, multiply the coefficients together and add the exponents.

The base will remain 10.

Ex 1. Multiply: (3.45 x 107) x (6.25 x 105)

First rewrite the problem as: (3.45 x 6.25) x (107 x 105)• Then multiply the coefficients and add the

exponents: 21.5625 x 1012 • Then change to correct scientific notation and

round to correct significant digits: 2.16 x 1013

NOTE - we add one to the exponent because we moved the decimal one place to the left.

Ex. 2. Multiply (2.33 x 10-6) x (8.19 x 103)

Rewrite the problem as: (2.33 x 8.19) x (10-6 x 103)

Then multiply the coefficients and add the exponents:

19.0827 x 10-3

Then change to correct scientific notation and round to correct significant digits 1.91 x 10-2

Division

Rule for Division - When dividing with scientific notation, divide the coefficients and subtract the exponents.

The base will remain 10.

Ex. 1 Divide 3.5 x 108 by 6.6 x 104

rewrite the problem as: 3.5 x 108 --------- 6.6 x 104

Divide the coefficients and subtract the exponents to get: 0.530303 x 104

Change to correct scientific notation and round to correct significant digits to get: 5.3 x 103

Note - We subtract one from the exponent because we moved the decimal one place to the right.

Addition/Subtraction

Rule for Addition and Subtraction - when adding or subtracting in scientific notation, you must express the numbers as the same power of 10.

This will often involve changing the decimal place of the coefficient.

Ex. 1 Add 3.76 x 104 and 5.5 x 102

Move the decimal to change 5.5 x 102 to 0.055 x 104

add the coefficients and leave the base and exponent the same:

3.76 + 0.055 = 3.815 x 104 following the rules for rounding, our final

answer is 3.815 x 104

Ex. 2 Subtract (4.8 x 105) - (9.7 x 104)

Move the decimal to change 9.7 x 104 to 0.97 x 105

subtract the coefficients and leave the base and exponent the same:

4.8 - 0.97 = 3.83 x 105 round to correct number of significant digits:

3.83 x 105

SIGNIFICANT FIGURESThe meaningful digits in a measured or

calculated quantityExcept when all the numbers involved are

integers (e.g. counting the number of students in class) it is often impossible to obtain the exact value of the quantity under investigation

When significant figures are counted the last digit is understood to be uncertain

•

SIGNIFICANT FIGURES

The amount of uncertainty depends on the measuring device used

Keeping track of significant figures in a measurement insures that calculations involving data will correctly represent the precision of the measurement

Precision represents agreement between several measurements of the same quantity

RULES FOR DETERMINING SIGNIFICANT FIGURES:

Insert picture document 9/2810• # #sig figs• • Any digit that is not zero is significant 845m 3• • • • Zeros between nonzero digits are 40,501 J 5• significant• • • Zeros to the left of the first nonzero digit 0.00035g 2• are not significant-indicate the placement • of the decimal point• • • If a number is greater than 1 then all zeros 3.040m 4• to the right of the decimal point • are significant • • • If a number is less than one then only zeros 0.090g 2• at the end of the number are significant- and 0.3005m 4 • zeros in between nonzero digits

DENSITY

Depends on:MassVolume

D = m/v (g/cm3)Mass usually expressed in gramsVolume usually expressed in cm3 or liters, etc.

What would take up more space??? A kilogram of feathers…..or a kilogram of steel??

OR

• The proximity of like atoms or molecules• More than just the “heaviness” of a

substance, density includes how much space an object takes up!!

• All substances have density including liquids, solids, and gases

Density is the measure of the “compactness” of a material

Gases• Real life application…..

• Low pressure weather system means warmer air tends to rise,

• High pressure systems indicate a colder more dense air mass that will…….

• SINK!!!

Balloon and liquid nitrogen

http://hyperphysics.phy-astr.gsu.edu/hbase/thermo/balloon.html#c1

http://paer.rutgers.edu/pt3/movies/TVrhoandFb.mov

www.dkimages.com



LIQUIDS• The more dissolved solids in a solution, the

more dense (such as ocean water)• Cold water in lakes tend to sink (this creates a

constant mixing of water, nutrients, and other substances)

– Kinetic energy again!!

Denser layers to less dense layers…..

What would happen????

• Mercury density = 13600kg/m3

• Lead density = 11340kg/m3

Lead floats on liquid mercury!

SOLIDS

Ice is less dense than water (which is why lakes and ponds have a thin layer of ice covering in winter, with water underneath)

Various rocks, woods, metals have a characteristic density specific to that substance

Wouldn’t you like to have a bunch of THIS dense material?

Solving for density

This image will help you in figuring out how to solve density problems:

What am I talking about?•It is the ability to perform

work•It may be transformed

•Some of you have very little•Some of you have too much

Energy!

• How is energy classified?

• How many types of energy can you name?

Potential vs. Kinetic EnergyEnergy associated with an objects position or distance

Also referred to as stored energy

The energy of motion

More motion equals more kinetic energy

Examples:A coiled springA boulder on top of a cliffChemical bonds

Examples:The uncoiling of a springThe boulder falling off of the cliffChemical bonds breaking

Potential Energy• Chemical Energy is energy stored in the bonds of atoms and molecules. It is the

energy that holds these particles together. Biomass, petroleum, natural gas, and propane are examples of stored chemical energy.

• Stored Mechanical Energy is energy stored in objects by the application of a force. Compressed springs and stretched rubber bands are examples of stored

mechanical energy.

• Nuclear Energy is energy stored in the nucleus of an atom––the energy that holds the nucleus together. The energy can be released when the nuclei are combined or split apart. Nuclear power plants split the nuclei of uranium atoms in a process called fission. The sun combines the nuclei of hydrogen atoms in a process called

fusion. Scientists are working on creating fusion energy on earth, so that someday there might be fusion power plants.

• Gravitational Energy is the energy of position or place. A rock resting at the top of a hill contains gravitational potential energy. Hydropower, such as water in a

reservoir behind a dam, is an example of gravitational potential energy.

Kinetic Energy• Electrical Energy is the movement of electrical charges. Everything is made of tiny particles

called atoms. Atoms are made of even smaller particles called electrons, protons, and neutrons.

• Radiant Energy is electromagnetic energy that travels in transverse waves. Radiant energy includes visible light, x-rays, gamma rays and radio waves. Light is one type of radiant energy.

Solar energy is an example of radiant energy.• Thermal Energy, or heat, is the internal energy in substances––the vibration and movement

of the atoms and molecules within substances. Geothermal energy is an example of thermal energy.

• Motion Energy is the movement of objects and substances from one place to another. Objects and substances move when a force is applied according to Newton’s Laws of Motion.

Wind is an example of motion energy.

• Sound is the movement of energy through substances in longitudinal (compression/rarefaction) waves. Sound is produced when a force causes an object or

substance to vibrate––the energy is transferred through the substance in a wave

Conservation of Energy

• It is never created or destroyed• It is transformed during a chemical reaction• Exothermic reactions (combustion) release

energy • Endothermic reactions absorb energy

What is the difference between Heat and Temperature?

• Think about a drinking glass full of water and a swimming pool full of water. They are both at 25 degrees C.

• Which one has more heat?• Formula time!!!

Calorimetry• The amount of heat released or absorbed during

a chemical reaction• The preferred SI unit for energy is the joule• We can also use the calorie• A calorie is the amount of heat required to raise

the temp of 1 gram of water 1 degree Celsius.• Also called specific heat or heat capacity• The specific heat of water is 1 cal/g*C• The food calorie is actually a Kilocalorie• There are 4.184 joules in a calorie

Calorimetry problems

• They allow us to track heat flow from one system or substance to another

• Here is the formula:• Heat ( H)= mass (g) x C (specific heat)x T (oC)• The means change in……

By the end of today..You will solve these

• How many joules are needed to warm 25.5 grams of water from 14oC to 22.5oC?

• Calculate the heat, in calories, needed to warm 225 grams of water from 88.0oC to its boiling point, 100.0oC.

Kinetic Theory

• All particles of matter are in constant, random motion.

• Sometimes particles collide and release heat• The kinetic energy of a system will determine

its phase or state

Phases of Matter (also state)

• Solid = definite volume and definite shape, Low kinetic energy

Phases of Matter

• Liquid = definite volume but can change shape by flowing, higher kinetic energy

Phases of Matter

• Gaseous or gas = no definite volume of shape, doesn’t really respond to gravity, highest kinetic energy

• If confined in a closed container it will fill the container but escape through any opening

Phase changes

• Matter can change phase depending on whether or not the system is losing or gaining kinetic energy

• Temperature does not change during a phase change

• The whole system has to change phase before temperature will change

• Heating curves or phase change diagrams measure the changes in temp. over time

Phase change diagram

Pressure!

• Force per unit area (shoes on a floor,syringe)• All phases exert a pressure• Solid = a block of ice on top of your head• Liquid= hundreds of liters of water on your

head• Gas = kilometers of atmospheric gas on your

head (mountains, sea level,airplanes)

Pressure!

Measuring Pressure

• Atmospheric pressure is the pressure exerted on your body by a column of gas above your head

• When you increase altitude you decrease the pressure (the column shrinks)

• Our bodies establish a pressure equilibrium• Sometimes our ears pop• Weather is often discussed in terms of

pressure systems

Measuring Pressure

• Units for pressure:• You’ve probably heard of pounds per square

inch or psi. Not needed in chem class• We use atm or atmospheric pressure• We use mmHg or millimeters of mercury• We use the Pascal or kiloPascal• 1.00 atm = 760mmHg = 101.325kPa

Graph it….Pressure (atm) Volume (mL) Constant = P x V

0.20 500

0.40 250

0.60 167

0.80 125

1.00 100

2.00 50.0

3.00 33.3

4.00 25.0

5.00 20.0

6.00 16.7

7.00 14.3

8.00 12.5

9.00 11.1

Graphical Relationships

• Indirect or inverse proportion= one when variable increases the other decreases

• Increase by a factor of 2 the other variable will decrease by a factor of 1/2

Graphical Relationships

• Direct relationship or proportion= when one variable increases the other variable increases

• One factor increases by a factor of 2 the other increases by a factor of 2

Robert Boyle and his law 1627-1691

Boyles Law

• He used a manometer to study the affect of pressure on the volume of a gas

Formula time!

• Four variables• You will be given three out of the four P1V1=P2V2

Pressure= atm or mmHg or KPa1 atm= 760mmHg= 101.325kPaVolume in liters or mL

Practice….

A sample of gas under a pressure of 720mmHg has a volume of 300mL. The pressure is increased to 800mmHg. What volume will the gas then occupy? Temp. is constant

P1=

V1=

P2=

V2=

What happens to the volume of a tire when it is very cold outside?

• Charles' Law states that at constant pressure, the volume occupied by a fixed amount of gas is directly proportional to its absolute temperature (K).

• Boyle’s Law= the temperature was held constant

• Charles’ law = the pressure is held constant• What about the temperature?

The Kelvin scale..it’s absolute

• If you decrease the temperature on a sample of gas the volume will decrease

• On the Celsius scale the volume will reach 0mL at -273o C

• The Kelvin scale changes this value to 0 • It is said as zero Kelvin not zero degrees Kelvin• So at 0 K the volume of a gas will be zero• This is called absolute zero..and it is very cold

Converting Temps.• K= oC + 273• oC= K – 273Standard temp. is 273KPractice:How many K is 100 o C ? ans= 373 KHow many oC are in 273 K?Ans = O o CAlways use Kelvin temperature for Charles’ Law

Why Kelvin?

• The relationship is not a direct proportion when Celsius is used

• With the Kelvin scale the line on the graph intercepts the y axis at the origin

Formula time!..Charles’ Law

Practice

• A gas is collected and found to fill 2.85 L at 25.0°C. What will be its volume at standard temperature?

Answer:

(298)(x)= (2.85)(273)X= 2.61 L

Practice…

• 4.40 L of a gas is collected at 50.0°C. What will be its volume upon cooling to 25.0°C?

Answer:

(323)(x)= (4.40)(298) x= 4.06L

Let’s review

• The volume of a gas is directly proportional to its temperature

• That means when the temperature of a gas decreases the volume of the gas will also decrease

• This relationship is know as Charles’ Law• Always use Kelvin temperature when solving

problems

Kinetic Theory and the Gas Laws

• These statements are made only for what is called an ideal gas.

• They cannot all be rigorously applied (i.e. mathematically) to real gases, but can be used to explain their observed behavior qualitatively.

• An ideal gas is one whose pressure-volume-temperature behavior can be explained by the ideal gas equation.


• All matter is composed of tiny, discrete particles (molecules or atoms).

• Ideal gases consist of small particles (molecules or atoms) that are far apart in comparison to their own size.

• The molecules of a gas are very small compared to the distances between them.


• These particles are in rapid, random, constant straight line motion. This motion can be described by well-defined and established laws of motion.

• There are no attractive forces between gas molecules or between molecules and the sides of the container with which they collide.


• Molecules collide with one another and the sides of the container.

• Energy can be transferred in collisions among molecules.

• The average kinetic energy of all the molecules is proportional to the absolute temperature.

This creates an Ideal Situation

• PV=nRT• What does that MEAN???!!!!• Pressure must be in atm• Volume must be in Liters• Temperature must be in Kelvin

Avogadro…Yo!

• Equal volumes of different gases, at the same temperature and pressure contain the same number of molecules

• Known as Avogadro’s Law• He coined the term the mole (mol)• Molar volume of an ideal gas is 22.4L• That means if you have 22.4 liters of gas you

have one mole of that gas• We will deal with Avogadro’s number later!!!!

So what about the ideal gas equation?

• PV=nRT• n is equal to the number of moles of gas• R is called the proportionality constant and • ya it never changes..it’s a constant!!!!• R= 0.0821 L*atm / K * mol• How many sig. figs.?• You will have to solve for three of the four

variables…so now…you must…

Well Helloooo!!!

You Practice Now!!!!!

How many moles of gas are contained in 890.0 mL at 21.0 °C and 750.0 mm Hg pressure?

0.374mol

Calculate the volume 3.00 moles of a gas will occupy at 24.0 °C and 762.4 mm Hg.

Buckyballs!!!

Atomic Structure

The Greeks..Democritus etc..

• 440 B.C.• All matter is made up of tiny invisible particles• Suggested that these particles could not be

divided• They named these particles “atomos”• That means indivisible in Greek

The evidence to support the existence of atoms

• The law of conservation of matter• Lavoisier 1770’s• Matter can neither be created or destroyed

during a chemical reaction• The mass of reactants is equal to the mass of

the products (Before and after)


• 1799 • French Chemist Joseph Proust• Law of definite proportions• The proportion by mass of the elements in a

given compound is always the same• Water always has an 8:1 ratio of Oxygen to

Hydrogen • No matter where you get the water from


• 1803 John Dalton working with carbon and hydrogen

• Law of multiple proportions• Two elements can combine to form more than

one compound (H20 and H202)

• The mass of one element and the mass of the other element are in small whole number ratios

Dalton’s Atomic Theory• 1803• All elements are composed of atoms• All atoms of the same element are exactly alike

(same mass)• Atoms of different elements are different

(different masses)• Compounds are formed by the joining of atoms

of two or more elements• In any compound the atoms are joined in whole

number ratios

Laviosier,Proust,Dalton,Democritus

Not so fast Johnny!!!..Modern Atomic Theory

• Atoms have a detailed structure that is temporarily changed during a chemical reaction

• Atoms can be changed from one element to another..In lab

• Atoms of the same element are not necessarily alike

• Atoms of the same element can and do have different masses

So… what do atoms look like?What are they composed of?

• 1897 English physicist William Crookes and his tube

• High voltage passed through in a vacuum and a light appeared

So… what do atoms look like?What are they composed of?

• Crookes suggested that the light rays were composed of electrically charged particles

• These ray were named cathode rays• 20 years later (that would be 1917 about)• J.J. Thomson proved, using a magnet that

these rays were deflected in the tube• He also determined that the rays were

actually particles in motion

J.J. Thomson

• He was a smart guy• He proved that these rays were not only

particles….• He proved that they have mass (very small)• He proved that they have a negative charge• So charged particles, negative charge?????• You people are so smart…• The Electron!!!

J.J. Thomson

• His picture

J.J. Thomson

20th century

• It was generally accepted that the atom was composed of a dense center called a nucleus

• This was yet to be confirmed or proved• Several models of the atom began to emerge• Before we look at the models let’s talk about

the sub-atomic particles

20th century

• Ok so Thomson discovered the electron (-)• Using Hydrogen it was determined that atoms

have some positive charge possibly located in the nucleus

• Henry Mosely R.I.P. determined that different elements have different numbers of protons

• He came up with the idea of atomic number• Atomic number = number of protons..more on

that later

20th century

• 1932 James Chadwick• He discovered that in addition to a positive

particle that has mass…• There is another particle that has almost equal

mass but no electric charge…it’s neutral????• The neutron!!!!• Usually elements have the same number of

protons and neutrons..When they don’t???• Isotopes!!!

20th century

• It is very difficult in nature to achieve sameness

• One out of 6000 Hydrogen atoms is an isotope• Isotopes can be used for a variety of things

Isotopes..what are they good for?

• Smoke Detectors and Americium-241• Agricultural Applications – radioactive tracers•

Isotopes..what are they good for?• Medical UsesBone imaging is an extremely important use of radioactive

properties. Supposed a runner is experiencing severe pain in both shins. The doctor decides to check to see if either tibia has a stress fracture. The runner is given an injection containing 99Tcm. This radioisotope is a gamma ray producer with a half-life of 6 hours.

After a several hour wait, the patient undergoes bone imaging. At this point, any area of the body that is undergoing unusually high bone growth will show up as a stronger image on the screen. Therefore if the runner has a stress fracture, it will show up on the bone imaging scan.

Do atoms have a mass?

• Of course..you need a really small scale• No…they are determined experimentally• Usually compared to Hydrogen• All you need to do is take the number of

protons (atomic number) and the number of neutrons and you get atomic mass

• All elements have a known atomic mass

Symbols…and numbers

• All elements on the periodic table have a symbol

• It is a capital letter sometimes followed by a lower case letter

• Is not necessarily the first letter of the element

• For example..find the element sodium• What is the symbol?

116

Atomic Mass on the Periodic Table

11Na

22.99

Atomic Number

Symbol

Atomic Mass

You practice now…..

Element

Number of Protons

Number of Neutrons

Number of Electrons

Atomic Mass Atomic Number

lithium

carbon

chlorine

silver

lead

What about models?

• J.J. Thomson and plum pudding• He said that the atom is a sea of positivecharge with pieces of negative charged distibuted throughout

What about models?

• Lord Ernest Rutherford 1909• Gold foil experiment• Most people thought that the atom was

mostly empty space• So Rutherford figured that if he shot a piece of

gold foil with positively charge alpha particles • The particles would pass through the foil and

illuminate a piece of film on the other side

Not so fast…

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf



Rutherford’s model

• Was wrong• He did come up with the name proton to

describe the positive center of the atom• Different elements have different numbers of

protons so they have different amounts of positive charge.

• His model suggests a circular orbit for electrons and…circular orbits don’t work

Rutherford’s model

The Bohr Model

• The planetary model• More accurate thanks to quantum mechanics• Light as a particle and so on..• Electrons can only be predicted to occupy a

certain space at any given time..they exist in orbitals that are not circular

The Bohr Model

You laugh now!!

Two atoms are walking down the street.One atom says to the other, “Hey! I think I lost

an electron!”The other says, “Are you sure??”“Yes, I’m positive!”

You laugh now!!!

A neutron walks into a restaurant and orders a couple

of drinks. As she is about to leave, she asks the waiter

how much she owes. The waiter replies, “For you,

No Charge!!!”

ElectronsElectrons (-) charge no mass located outside the nucleus

ProtonsProtons (+) charge 1 amu located inside the nucleus

NeutronsNeutrons no charge 1 amu located inside the nucleus

Particles in the Atom

Modern View

• The atom is mostly empty space• Two regions

– Nucleus • protons and neutrons

– Electron cloud• region where you might find an electron

What happens when atoms lose or gain electrons?

• Some elements cant hold onto their electrons• Some elements are good at pulling electrons

away from other elements• When an electron is gained or lost….• Ions= charged particles• (+) ions = cations• (- ) ions = anions

How many electrons?

11

Na22.99

Ions

• Positive ions have lost electrons • Negative ions have gained electrons• Symbols for ions are written with a charge as a

superscript in the upper right• A +1 charge can be written as just +• Same thing for a -1

Ions

• Na+1 Lost or gained?• How many electrons?• How many protons?• Cl-1 Lost or gained?• How many electrons?• How many protons?

Enough about ions..Where do the electrons go? Electron Configuration• Understand that electrons spin around the

nucleus• The path they take is called an orbital and it has

shape• All electrons are attracted to the nucleus• Only a limited number of electrons can get front

row seats• The farther away from the nucleus the more

energy these electrons have

Where do the electrons go?

• Ok so electrons have to occupy what are called principal energy levels

• Like rows in a theatre• These energy levels are numbered 1,2,3,4 etc..• They are called principal quantum numbers (n)• The higher the number the farther away the

electron is from the nucleus


• So the principal quantum number is equal to the number of sublevels

• What??? • n=1 …that means that there is one sublevel• n=2?• Right…2 sublevels• So what the heck is a sublevel?


• Sublevels are like sections in a theatre..• They are regions within an energy level where

you find electrons• There are four sublevels in atoms and they are

labeled with letters : s,p,d,f• Why letters?...Well they refer to the

shapes..let’s leave it at that.• Think of these sublevels as different kinds of

seats in a theatre..big and cumfy or small

Tweeter Center

Let’s Review

Sublevel # orbitals # electrons

s 1 2

p

3 6

d 5 10

f 7 14

Now we’re ready!..orbital notation..configuration and dots

• Orbital notation..all about the spin…Some rules

• Boxes represent orbitals• Arrows represent spin

Some Rules

• Pauli Exclusion Principle

– Each orbital can hold TWO electrons with

opposite spins.

Some Rules

Aufbau Principle

• Electrons fill the lowest energy orbitals first.

• “Lazy Tenant Rule”

Some Rules

• Hund’s Rule– Within a sublevel, place one e- per orbital before

pairing them.

– “Empty Bus Seat Rule”

Incorrect Correct

Orbital Filling

You go now!!!

• Use your table and complete the orbital notation for the first 5 elements

Bond…ionic bond

Bonding

Chemical Bonding is two or more atoms combining by sharing electrons so that a new substance is produced that has different physical and chemical properties than its component elements

Not so fast..

• What is a valence electron?• Electrons in the outer most energy level or

shell or orbit• O.K. So we can use what are called electron

dot diagrams to show these valence electrons• These dot diagrams are also called Lewis

structures

Dots

• We will use a generic symbol “X” to represent any element

• The valence shell has an s orbital and possibly 3 p orbitals

• Why will a valence shell never have a d orbital• The s orbit will hold 2 electrons • The p orbitals will hold a total of 6• So elements can have a maximum of eight

valence electrons

The Octet Rule• Atoms themselves can gain or lose electrons so that they have a

full valence shell of 8 (or 2) electrons. . .

• -- when atoms gain or lose electrons, they are called ions

• Depending on whether or not the atom gains or loses an electron, it can be positively-charged (cation) or negatively-charged (anion)

• This need to be stable and have 8 electrons is called the octet rule

• In addition to forming ions, atoms can also “share” their valence electrons with other atoms, giving each atom 8 valence electrons

Back to dots..I’ll draw them

Element # valence e’s

X

2

X 4 x 6 X 8

Now try this…

• Write the e dot configuration for Mg and Cl?

Ionic Bonding..electron transfer

• Electrons are lost or gained• Metals tend to lose valence electrons to

satisfy the octet rule• Non-metals tend to gain electrons to satisfy

the octet• The number gained or lost has to do with the

closest noble gas

Ionic Bonding..Naming ionic compounds

• Use electron dot to determine if electrons are lost or gained

• The metal’s or anions name does not change• The nonmetal you need remove the end of

the name and add –ide• Usually the last two or three letters are

removed• The cation is always first in the name

Ionic Bonding..Naming ionic compounds

So once again…• Ionic bonding is the combination of two or more ions to form an

electrically-neutral compound

How Ionic Bonding Works

1) The giving atom loses a valence electron (or 2 or 3) so that it has a full valence shell, but a positive charge

2) The gaining atom gains a valence electron (or 2 or 3) so that it has a full valence shell, but a negative charge

3) The negative and positively-charged ions are attracted to each other (like a magnet) based on their opposite charge

If you look at your periodic table, you will see that ionic bonding usually occurs in compounds formed between one metal and one nonmetal

And again..naming ionics• Ionic compounds all have two-word names

• The first word in the name is the same as the name of the first cation (for example, sodium, ammonium, potassium, etc)

The second word in the name is either:

1. If the second ion is polyatomic, it is just the name of the polyatomic ion

2. If the second ion is an element, the end of the element’s name changes to –ide

Example: chlorine chloride

oxygen oxide

You're getting sleepy

Characteristics of ionics

Moving on….characteristics of ionics

• Well..they are all called “salts” and compounds

• Crystal Pattern – every ion is attracted to all other ions with the opposite charge

• this results in a repeating 3-dimensional crystal pattern


• High Melting Point – the attraction in the crystal pattern leads to very strong bonds, making it hard to break apart ionic compounds

-- ionic compounds melt at high temperatures

• Conductivity – when dissolved in water, ionic compounds conduct electricity

-- in water, the bonds dissociate (fall apart), leaving lots of ions to carry charge

• Solid ionic compounds do not conduct electricity very well

-- melted ionic compounds do conduct electricity fairly well

Polyatomic Ions• There are some ions that are made up of more

than one type of atom, these are called polyatomic ions

• For example, the polyatomic ion known as ammonium NH4

+ has 4 atoms of hydrogen and one atom of nitrogen, HOWEVER, the whole ion has an overall charge of +1

• -- you will be given the charges of any polyatomic ions

The polyatomics…Name Formula/Charge Name Formula/Charge

acetate CH3COO-1 dichromate Cr2O7-2

acetate C2H3O2-1 formate HCOO-1

cyanide CN-1 permanganate MnO4-1

carbonate CO3-2 ammonium NH4

+1

bicarbonate HCO3-1 nitrite NO2

-1

bromate BrO3-1 nitrate NO3

-1

oxalate C2O4-2 hydroxide OH-1

hypochlorite ClO-1 phosphate PO4-3

chlorite ClO2-1 sulfite SO3

-2

chlorate ClO3-1 sulfate SO4

-2

perchlorate ClO4-1 thiosulfate S2O3

-2

How do we deal with them?

• Treat them as more than one atom together acting as a single unit.

• And they have a name that is usually different than the individual atoms

• They also have an overall charge.• For example:• NH4

+ This is ammonium, it is a cation

• It consists of one Nitrogen and 4 hydrogen atoms


• How ‘bout OH- • Well It’s called hydroxide and it is negative• This polyatomic consists of one oxygen and

one hydrogen• Ok so what happens when we combine

ammonium and hydroxide?• We get ammonium hydroxide• What is the formula?


• NH4OH

• The charges cancel each other out!!• What about ammonium phosphate?• Well we know NH4

+

• And phosphate is PO4-3

• We need to use parentheses• So what do you get?


• (NH4)3PO4

• Ammonium phosphate..there you have it..• The parentheses go around the whole ion!!• How many Hydrogen atoms in this compound• Now we need to deal with this roman numeral

business..• Copper(II)sulfate for example• What the heck is the roman numeral all about?

Variable Oxidation states

• The oxidation number is equal to the number of electrons an atom either gains or loses

• The number is equal to the charge• So Na has an oxidation number of +1• Technically these numbers are the basis for a group

of reactions called oxidation reduction reactions.• If an atoms oxidation number is decreased that

element is reduced and if the number increases the element is oxidized.


• Enough about that nonsense…that’s later• In polyatomics the sum of the oxidation numbers

of all the elements in the ion must equal the net charge of the ion.

• What??? Ok so NH4+1

• Nitrogen is -3 and there are 4 Hydrogen at +1 each so…

• -3 + (4x1)= +1• You try Hydroxide


• Well some elements (the ones in the middle) can lose two, three, four electrons depending on who they are bonded to..they vary!!!

• So when naming ionic compounds that have elements from the transition metals (middle) you need to indentify the charge and use a roman numeral to indicate which variable oxidation number is being used


• So give it a shot..• (Cu)2(SO4)3

• Name it• Copper(III) sulfate• You have to put parentheses around the

roman numeral

Now here’s the trickCriss/Cross

• When you now the charge..• Simply criss/cross the charges to write the

formula• Superscripts become subscripts• The anion charge is now the subscript for the

cation and vice/ versa• Do it for Mg and Cl• You need parenthesis' for polyatomics

You try…

• I’ll give you the examples

What about molar mass?

• By definition it is the mass of a compound that equals one mole of the compound

• You take the the individual mass of each element and add them together

• For example:

Let’s Review

Ionic Bonds

Metals and Nonmetals

Complete transfer of Electrons

Metals lose and nonmetals gain

Metals are on the left before the

staircase

They form crystals

High melting point

They are good conductors of

electricity

Covalent bonding You need to know this table

Prefix Number

Mono One

Di Two

Tri Three

Tetra Four

Penta Five

Hexa Six

Hepta Seven

Octa Eight

NonaDeca

NineTen

You need to know the diatomics

• The diatomics are a group of elements that are never found alone in nature.

• They come is pairs of two..unless they are ions• BrINClHOF• Hydrogen • Nitrogen• Oxygen• Fluorine• Bromine• Chlorine• Iodine

Covalent Bonding..Sharing

• Involves nonmetals and nonmetals• They share valence electrons• Sometimes equally and sometimes not• No ions are formed and the end result is called a

molecule• Covalent molecules have low boiling and melting

points and are poor conductors of electricity• They are usually liquid or gas at room temp.


• All biological macromolecules are covalent• Most of the food you eat are covalent

molecules..Carbs, Lipids(oils),Proteins,DNA• They are also solvents like alcohol and acetone• Including water…But why don’t water and oil

mix?• It turns out that they are different kinds of

covalent bonds…but let’s name them first…


• Yes we have to name them• We need to use prefixes to indicate the number of

atoms in a covalent molecule• For example: CO • Carbon Monoxide• Notice that there is not a prefix used for the first

element when there is only one atom• What about H20?

• Dihydrogen Monoxide

Covalent Bonding cont..

• When atoms don’t share equally• Electronegativity is the ability to pull electrons away

from other atoms• It is a numerical scale from 0-4• 0 is the least electronegative which means these

atoms cannot hold onto electrons and are very reactive

• 4.0 is the highest they are very strong pullers of electrons and are also very reactive

• Which atoms on the table are the most reactive?

Electronegativity Table

Types of Sharing

Non-Polar Covalent

Polar Covalent

Ionic

Equal sharingElectronegativity difference is 0.0-.03

Unequal sharingElectronegativity difference is 0.3-1.7

No sharing=transferElectronegativity difference is 1.7-4.0

All of the diatomsLike Cl2

Water All salts

Polarity

• When a molecule is polar it has a region of partial positive charge around the lower electronegative atom and partial negative charge around the higher electro. value

• H2O is an example of a polar molecule

Lewis Structures

Lewis structures are a way to write chemical compounds where all the atoms and electrons are shown. Sometimes, people have a lot of trouble learning how to do this. However, using the methods on this page, you should have very little trouble.

The first method given allows you to draw Lewis structures for molecules with no charged atoms, while the second allows you to do it for charged molecules (such as polyatomic ions).

Lewis Structures• How to draw Lewis structures for molecules that contain no charged atoms• 1) Count the total valence electrons for the molecule: To do this, find the number of valence electrons for each atom in the

molecule, and add them up. • 2) Figure out how many octet electrons the molecule should have, using the octet rule: The octet rule tells us that all

atoms want eight valence electrons (except for hydrogen, which wants only two), so they can be like the nearest noble gas. Use the octet rule to figure out how many electrons each atom in the molecule should have, and add them up. The only weird element is boron - it wants six electrons.

• 3) Subtract the valence electrons from octet electrons: Or, in other words, subtract the number you found in #1 above from the number you found in #2 above. The answer you get will be equal to the number of bonding electrons in the molecule.

• 4) Divide the number of bonding electrons by two: Remember, because every bond has two electrons, the number of bonds in the molecule will be equal to the number of bonding electrons divided by two.

• 5) Draw an arrangement of the atoms for the molecule that contains the number of bonds you found in #4 above: Some handy rules to remember are these:

– Hydrogen and the halogens bond once. – The family oxygen is in bonds twice. – The family nitrogen is in bonds three times. So does boron. – The family carbon is in bonds four times.

• A good thing to do is to bond all the atoms together by single bonds, and then add the multiple bonds until the rules above are followed.

• 6) Find the number of lone pair (nonbonding) electrons by subtracting the bonding electrons (#3 above) from the valence electrons (#1 above). Arrange these around the atoms until all of them satisfy the octet rule: Remember, ALL elements EXCEPT hydrogen want eight electrons around them, total. Hydrogen only wants two electrons.

Lewis Structures• Let's do an example: CO2

Note: Each of the numbers below correspond to the same numbered step above.• 1) The number of valence electrons is 16. (Carbon has four electrons, and each of the oxygens have

six, for a total of 4 + 12 = 16 electrons). • 2) The number of octet electrons is equal to 24. (Carbon wants eight electrons, and each of the

oxygens want eight electrons, for a total of 8+16 = 24 electrons). • 3) The number of bonding electrons is equal to the octet electrons minus the valence electrons, or 8. • 4) The number of bonds is equal to the number of bonding electrons divided by two, because there

are two electrons per bond. As a result, in CO2, the number of bonds is equal to 4. (Because 8/2 is 4). • 5) If we arrange the molecule so that the atoms are held together by four bonds, we find that the

only way to do it so that we get the following pattern: O=C=O, where carbon is double-bonded to both oxygen atoms.

• 6) The number of nonbonding electrons is equal to the number of valence electrons (from #1) minus the number of bonding electrons (from #3), which in our case equals 16 - 8, or 8. Looking at our structure, we see that carbon already has eight electrons around it. Each oxygen, though, only has four electrons around it. To complete the picture, each oxygen needs to have two sets of nonbonding electrons, as in this Lewis structure:

The Mole….

How long would it take to spend a mole of $1 coins if they were being spent at a

rate of 1 billion per second?

Background: atomic masses

• Look at the “atomic masses” on the periodic table. What do these represent?

• E.g. the atomic mass of C is 12 (atomic # is 6)

• We know there are 6 protons and 6 neutrons• Protons and neutrons have roughly the

same mass. So, C weighs 12 u (atomic mass units).

• What is the actual mass of a C atom?

Background: atomic masses

• Answer: approx. 2 x 10-23 grams (protons and neutrons each weigh about 1.7 x10-24 grams)

Two problems1. Atomic masses do not convert easily to

grams2. They can’t be weighed (they are too

small)

Background: Molecular weight

• The molecular weight of a substance is the weight in atomic mass units of all the atoms in a given formula.

• The molecular weight of a substance is needed to tell us how many grams are in one mole of that substance.

• The mole is the standard method in chemistry for communicating how much of a substance is present.


• You should have a periodic table for looking up atomic weights and a calculator.

• Point #1 - You need to know how many atoms of each element are in a substance in order to calculate its molecular weight.


• For example H2O has two atoms of hydrogen and one atom of oxygen. H2O2 has two atoms each of oxygen and hydrogen. Mg(OH)2 has one atom of magnesiun and two each of oxygen and hydrogen.

• If a subscript follows an atom with no parenthesis, that number tells you how many of that atom are present. If parentheses are involved, you must multiply each subscript inside by the one which is outside.


• How many of each element are in the following examples:

• KCl • Fe2O3

• Al(NO3)3

• NH4NO3


• Point#2 - You need to know the atomic weight of each element in order to calculate the molecular weight of the substance.

• That would be the number that is a decimal• You can round…….

How to calculate the molecular weight of a substance

• Multiply each element's atomic weight by how many atoms are present in the formula, then add the answers.

• Example #1 - Al2(SO4)3 There are: • two atoms of aluminum and the atomic weight of Al is 26.98 amu. • three atoms of sulfur and the atomic weight of S is 32.06 amu. • twelve atoms of oxygen and the atomic weight of O is 16.00 amu. First multiply: • 2 x 26.98 = 53.96 total weight of all Al in formula • 3 x 32.06 = 96.18 total weight of all S in formula • 12 x 16.00 = 192.00 total weight of all O in formula Then add: 53.96 + 96.18 + 192.00 = 342.14 amu. • This answer, 342.14 amu, represents the molecular weight of Al2(SO4)3

Review

• Step One: Determine how many atoms of each different element are in the formula.

• Step Two: Look up the atomic weight of each element in a periodic table.

• Step Three: Multiply step one times step two for each element.

• Step Four: Add the results of step three together and round off as necessary.

The mole

• Atomic mass and molecular mass are only useful if they are measured in grams

• The mole is the standard method in chemistry for communicating how much of a substance is present.

• In one mole, there are 6.022 x 1023 atoms. Here's another way: there are 6.022 x 1023 atoms of carbon in 12 grams of carbon.

The mole

• Let's say that real clearly: one mole of ANYTHING contains 6.022 x 1023 entities.

• One mole of donuts contains 6.022 x 1023 donuts

• One mole of H2O contains 6.022 x 1023 molecules

• One mole of nails contains 6.022 x 1023 nails • One mole of Fe contains 6.022 x 1023 atoms

The mole

• 6.022 x 1023 is so important in chemistry that it has a name.

• It is called Avogadro's Number and has the symbol N.

• It is so named in honor of Amedeo Avogadro, an Italian chemist, who, in 1811, made a critical contribution (recognized only in 1860 after his death) which helped greatly with the measurement of atomic weights.

The mole

• Counting atoms or molecules is very difficult since they are so small. However, we can "count" atoms or molecules by weighing large amounts of them on a balance.

• When we weigh one mole of a substance on a balance, this is called a "molar mass" and has the units g/mol (grams per mole). This idea is very critical because it is used all the time.

The mole

• A molar mass is the weight in grams of one mole.

• One mole contains 6.022 x 1023 entities. • Therefore, a molar mass is the mass in grams

of 6.022 x 1023 entities.• OK. How does one calculate a molar mass? You

already know how to calculate a molar mass.• The molar mass of a substance is the molecular

weight in grams. Just use the unit “g/mol”

The mole

• You try!!!• Calculate the molar mass of Al(NO3)3

• 1 x (26.98) + 3 x (14.007) + 9 x (16.00) = 213.00 g/mol

• 213.00 grams is the mass of one mole of aluminum nitrate.

• 213.00 grams of aluminum nitrate contains 6.022 x 1023 entities of Al(NO3)3

The mole

• How long would it take to spend a mole of $1 coins if they were being spent at a rate of 1 billion per second?

• $ 6.02 x 1023 / $1 000 000 000 = 6.02 x 1014 payments = 6.02 x 1014 seconds

6.02 x 1014 seconds / 60 = 1.003 x 1013 minutes1.003 x 1013 minutes / 60 = 1.672 x 1011 hours1.672 x 1011 hours / 24 = 6.968 x 109 days6.968 x 109 days / 365.25 = 1.908 x 107 years

A: It would take 19 million years

Mole ConversionsGiven grams, convert to moles

• In chemistry, the mole is the standard measurement of amount. However, balances DO NOT give readings in moles. Balances give readings in grams.

• So the problem is that, while we compare amounts of one substance to another using moles, we must also use grams, since this is the information we get from balances.


There are three steps to converting grams of a substance to moles.

• Determine how many grams are given in the problem.

• Calculate the molar mass of the substance.• Use the molar mass as a unit factor and

multiply (cancel unwanted units)


• Remember unit factors are quantities that equal one..like 100 pennies = $1.00

• So molar mass is the same• For example: H20 has a molar mass of:

18.01 grams 1 molYou can also write it as 1mol 18.01grams


• So let’s try one• Convert 25.0 grams of KMnO4 to moles.

• Step One: The problem will tell you how many grams are present. Look for the unit of grams. The number immediately preceeding it will be how many grams. Common abbreviations for grams include g (just the letter) and gm.

• The problem gives us 26.0 g


• Step Two: You need to know the molar mass of the substance.

• The molar mass of KMnO4 is 158.034 grams/mole.

• Step Three: use the molar mass as a unit factor and multiply

25.0g x 1mole = 0.158mol 158.034g


• Now you• 2.00 grams of H2O

• 75.57 grams of KBr• 100. grams of KClO4

• 0.111mol• 0.6350mol• 0.722mol

Given Moles, Convert to grams

• Ok again use the molar mass as a unit factor and follow the steps.

• How many grams are there in 0.5 mol of KMnO4 ?

• Step #1: identify what you have been given and what you want…give moles, want grams

• Step #2: find the molar mass: 158.034 g/mol.

Given Moles, Convert to grams

• Step # 3: use the factor-label method and multiply..cancel units

0.5mol x 158.034g = 79.02g 1molYou try: 0.25mol of H2O

0.125mol of KBr4.50g H2O, 14.75g KBr

Use Avagadro’s number as a unit factor to count molecules and

atoms• Refer to calculating with numbers in scientific

notation of page 641• Addition and subtraction numbers must be in

same power of ten• Multiplication you add the exponents • Division you subtract the exponents

Use Avagadro’s number as a unit factor to count molecules and

atoms

• The unit factor is 6.022x 1023 molecules• 1mole• So when you are given grams you must first

convert to moles and then to molecules• How many molecules are there in 20 grams of

NaCl?

Converting to molecules

• So we know the molar mass is 58 grams per mol

• given 20g x 1mole = 0.34 x6.022x1023molecules

58 g 1mol = 0.34 x 6.022= 2.05 x 1023 molecules There is no change in the power of ten!!!!

Converting to grams from molecules

• Solve for moles and then grams• Check your power of ten!!

Chemical Equations

• A condensed statement of facts about a chemical reaction

• A chemical reaction is any action that brings about a chemical change

• Atoms are rearranged to form something different

Chemical Equations

• There are two ways to express a chemical equation• With words: hydrogen gas reacts with oxygen gas to

produce/yield water• With symbols: H2 + O2 H2O

• All substances on the left side of the arrow are called reactants and they exist before the reacation

• All substances on the right side of the arrow are called products and they exist as a result of the reaction. They are formed from reactants

Chemical Equations

• The world needs balance• Matter is not created or destroyed in a

chemical reaction• Niether is mass• They are rearranged but the number of

reactant atoms always equals the number of product atoms. They have to balance

Chemical Equations

• We can use large numbers called coefficients placed in front of molecules or single elements to change the numbers or atoms on both sides.

• We cannot break chemical bonds and place coefficients between atoms in one molecule

Chemical Equations

• So our example:H2 + O2 H2O Not balanced

2H2 + O2 2 H2O

Balanced!!!Two molecules of hydrogen gas reacts with one

molecule of oxygen gas to produce two molecules of water.

Chemical Equations

• 1. ____ Sb + ____ Cl2 -----> ____ SbCl3

• 2. ____ Mg + ____ O2 -----> ____ MgO

• 3. ____ CaCl2 -----> ____ Ca + ____ Cl2

• 4. ____ NaClO3 -----> ____ NaCl + ____ O2

% composition• Compounds are made up of two or more

elements• The law of definite proportion states that the

proportion, by mass of the elements in a given compound is always the same

• All samples of water are 11% hydrogen and 89% oxygen by mass

• So % comp is the percentage by mass of each element in a molecule

% composition

• A sample of a compound containing carbon and oxygen had a mass of 88g. Experimental evidence showed that 24g of this sample was carbon and the remaining 64g was oxygen. What is the percentage composition of this compound?

Empirical Formula

• the formula of a compound expressed as the smallest possible whole-number ratio of subscripts of the elements in the formula

• Molecular formula C6H12O6

• Empirical formula CH2O

• Molecular formula=the formula of a compound in which the subscripts give the actual number of each element in the formula

Empirical Formula

• Notice two things:• 1. The molecular formula and the empirical

formula can be identical.2. You scale up from the empirical formula to the molecular formula by a whole number factor.

Empirical Formula

• Percent to mass

• Mass to mole

• Divide by small

• Multiply 'til whole

Empirical Formula

• Here's an example of how it works.• A compound consists of 72.2% magnesium

and 27.8% nitrogen by mass. • What is the empirical formula?

Empirical Formula

• Percent to mass: • Assume 100 g of the substance, then 72.2 g

magnesium and 27.8 g nitrogen. • Mass to moles: for Mg: 72.2 g Mg x (1 mol Mg/24.3 g Mg) = 2.97

mol Mgfor N: 27.8 g N x (1 mol N/14.0 g N) = 1.99 mol N

Empirical Formula

Divide by small: • for Mg: 2.97 mol / l.99 mol = 1.49

for N: 1.99 mol / l.99 mol = 1.00 Multiply 'til whole: • for Mg: 2 x 1.49 = 2.98 (i.e., 3)

for N: 2 x 1.00 = 2.00 and the formula of the compound is Mg3N2.

Chemical reactions

• We can name• We can identify types of bonds• We can deal with moles • So what’s left?• We have to classify reactions• You need to know 5 classes of chemical

reactions….

Chemical reactions

• This is going to be an overview• These reactions can be tricky but we are going

to keep it simple• The five reactions are : Synthesis,

Decomposition, Single replacement, Double replacement and Combustion

Chemical reactions

• First some vocabulary:• Binary Compound= compound consisting of

two elements• Binary salt= metal combined with nonmetal• Metallic oxide= a compound composed of a

metal and oxygen• Carbonate= CO3

• Chlorate=ClO3

Chemical reactions

• Synthesis• That means that two pieces join together to

produce one, more complex compound. • These pieces can be elements or simpler

compounds. • Complex simply means that the product

compound has more atoms than the reactant molecules. Usually!!

Chemical reactions

• Written using generic symbols, it is usually shown as: A + B ---> AB

These are some examples: Mg + O2 ---> MgO

H2 + O2 ---> H2OK + Cl2 ---> KClFe + O2 ---> Fe2O3

Chemical reactions

• Notice that two elements are combining in each example.

• Synthesis can also be two compounds making a more complex compound (or a compound and an element joining together) as in these examples:

• CaO + CO2 ---> CaCO3

Na2O + CO2 ---> Na2CO3

KCl + O2 ---> KClO3

Ba(ClO3)2 ---> BaCl2 + O2

Chemical reactions

• Notice how, in every case so far, there is only one substance on the right-hand (product) side. This is not always the case in a synthesis reaction.

• Categories: 1) Direct union of two elements will produce a binary

compound.2) Metallic oxides and carbon dioxide react to produce carbonates.3. Binary salts and oxygen react to produce a chlorate.

Chemical reactions

• Decomposition: During decomposition, one compound splits apart into two (or more pieces). These pieces can be elements or simpler compounds

• Written using generic symbols, it is usually shown as:

• AB ---> A + B

Chemical reactions

• However, that really only works for splitting apart into the elements, like these examples.

• HgO ---> Hg + O2

H2O ---> H2 + O2

MgCl2 ---> Mg + Cl2

FeS ---> Fe + S

Chemical reactions

• Decomposition can also split one compound into two simpler compounds (or compound and an element) as in these examples:

• CaCO3 ---> CaO + CO2

Na2CO3 ---> Na2O + CO2

KClO3 ---> KCl + O2

Ba(ClO3)2 ---> BaCl2 + O2

Chemical reactions

• Notice how, in every case so far, there is only one substance on the left-hand (reactant) side. This is always the case in a decomposition reaction. Don't forget that!!

• Figuring out what the products are in decomposition is harder (maybe you'll think it's easier!!) because you will have to recognize several categories of decomposition reactions

Chemical reactions

• Here are your first three: • 1) All binary compounds (like the four in the first

example set above) will break down into their elements.2) All carbonates (like the first two in the second example set above) break down to the oxide and carbon dioxide.3. Chlorates (like KClO3 and Ba(ClO3)2 in the example) will break down to the binary salt and oxygen.

Chemical reactions

• Single replacement, one element replaces another element in a compound. There are two different possibilities:

• 1. One cation replaces another. Written using generic symbols, it is:

AX + Y ---> YX + A

Chemical reactions

• Element Y has replaced A (in the compound AX) to form a new compound YX and the free element A. Remember that A and Y are both cations (postively-charged ions) in this example.

• Some examples are: • Cu + AgNO3 ---> Ag + Cu(NO3)2

Fe + Cu(NO3)2 ---> Fe(NO3)2 + Cu

Chemical reactions

2. One anion replaces another. Written using generic symbols, it is:

A + XY ---> XA + Y

Chemical reactions

• Element A has replaced Y (in the compound XY) to form a new compound XA and the free element Y. Remember that A and Y are both anions (negatively-charged ions) in this example.

• The only examples the ChemTeam knows about involve halogens, so here are two examples:

• Cl2 + NaBr ---> NaCl + Br2

Br2 + KI ---> KBr + I2

Chemical reactions

• In single replacement, one reactant is always an element. It does not matter if the element is written first or second on the reactant side. The other reactant will be a compound.

Chemical Reactions

• During double replacement, the cations and anions of two different compounds switch places.

• Written using generic symbols, it is: AB + XY ---> AY + XB

Chemical Reactions

• A and X are the cations (postively-charged ions) in this example, with B and Y being the anions (negatively-charged ions).

Here is another way to look at the above generic example: a) the outside portions (the cation A and anion Y) combine

to make a formula called AY.b) The inside portions (the anion B and the cation X)

switch order so that X (postively charged) goes first and B (negatively charged) goes second making a formula called XB.

Chemical Reactions

• Keep in mind that, when it comes to writing actual formulas, you MUST write chemically correct formulas. Please do not assume from the AY and XB examples that the product formulas will always be one-to-one in terms of positive and negative.

Some examples of actual reactions are: • KOH + H2SO4 ---> K2SO4 + H2O

FeS + HCl ---> FeCl2 + H2SNaCl + H2SO4 ---> Na2SO4 + HCl

Chemical Reactions

• Combustion, at its most general, can mean the reaction of oxygen gas (O2) with anything.

• However, we will understand combustion to mean the reaction of oxygen with an compound containing carbon and hydrogen. A common synonym for combustion is burn.

• Written using generic symbols, it is usually shown as: • CxHy + O2 ---> CO2 + H2O

Molar ratios

• The molar ratio will assume a place of central importance in solving stoichiometry problems. The sources for these ratios are the coefficients of a balanced equation.

• We will look at what a molar ratio is and then a brief word on how to recognize which ratio to use in a problem.

Molar ratios

• 2 H2 + O2 ---> 2 H2O

• What is the molar ratio between H2 and O2?

• Answer: two to one. So this ratio in fractional form is:

• The ChemTeam recommends you explicitly write a one in the denominator of the ratio.

2 H2 + O2 ---> 2 H2O

• What is the molar ratio between O2 and H2O?

• Answer: one to two. • As a fraction, it is:

2 H2 + O2 ---> 2 H2O

• What is the molar ratio between H2 and H2O?

• Answer: two to two or:

Stoichiometry

• Stoichiometry deals with calculations about the masses (sometimes volumes) of reactants and products involved in a chemical reaction.

• Chemists and chemical engineers must perform calculations based on balanced chemical reactions

• to predict the cost of processes.• • These calculations are used to avoid using

large excess amounts of costly chemicals.

Stoichiometry

• 2 NO(g) + O2(g) à 2 NO2(g)• • We can now read the balanced chemical equation as

“_______________________________ gas• react with _________________________________ gas to

produce ________________________• gas”.• • The coefficients indicate the

________________________________, or the ratio of the moles, of

• reactants and products in every balanced chemical equation.

Stoichiometry Mass-Mass


• Here is a typically-worded problem: • Given 20.0 grams of A and sufficient B, how many

grams of C can be produced? • You will need to use molar ratios, molar masses,

balancing and interpreting equations, and conversions between grams and moles.


• The Steps Involved in Solving Mass-Mass Stoichiometry Problems

• Make sure the chemical equation is correctly balanced. • Using the molar mass of the given substance, convert

the mass given in the problem to moles. (1st bridge) • Construct a molar ratio and use it to convert to moles

of the unknown. (2nd bridge) • Using the molar mass of the unknown substance,

convert the moles just calculated to mass. (3rd Bridge)


• How many grams of chlorine can be liberated from the decomposition of 64.0 g. of AuCl3 by this reaction: 2 AuCl3 ---> 2 Au + 3 Cl2

Bye-Bye Stoichiometry

Hello solutions…

• Take a minute and see if you can define the word solution and come up with a couple of examples…

• Homogeneous mixtures• The components are uniformly dipsersed• You cannot see the individual components• A solution is a homogenous mixture made up of

very small particles that are individual molecules, atoms or ions

Properties of Solutions

• They are clear• They cannot be filtered • They will generally last forever• All solutions contain two parts• #1 Solvent= does the dissolving• #2 Solute= is dissolved• Solutions in which water is the solvent are

called aqueous

Properties of Solutions

A solution is a A solution is a ______________________________ mixture of 2 or more mixture of 2 or more substances in a single substances in a single phase. phase.

One constituent is One constituent is usually regarded as usually regarded as the the SOLVENTSOLVENT and the and the others as others as SOLUTESSOLUTES..

Types of Solutions

• Gaseous = consist of gases or vapors dissolved in one another

• Air• Liquid = liquid solvent and solid liquid or gas

as solute• When liquids dissolve in each other they are

miscible• When they don’t dissolve they are immiscible

Types of Solutions

• Solid=mixtures of solids uniformly spread throughout one another at the molecular or atomic level

• Alloys like brass

Solubility..the ability to go into a solution

• Insoluble= cannot be dissolved• Like dissolves like• That means that chemically similar substances

will dissolve one another• Like ionic and ionic• When water dissolves an ionic like salt the salt

ions dissociate or break apart and become uniformly distributed (dissolved)

Solubility

• Factors that effect solubility:• Nature of the solvent/solute• Temperature• Pressure Henry’s law---Nucleation and

mentos

Solubility curves

• Shows how much solute will dissolve in a given amount of solvent over a range of temperatures

Solubility

• Saturated solution= a solution that has dissolved in it all the solute that it can normally hold at given conditions

• Usually a precipitate will form if a solution is saturated

• A precipitate means that solute material is not dissolving

Solubility

• A solution that contains less solute than it can handle is said to be unsaturated

• Under special conditions some solvents can become supersaturated.

• This means that they will hold temporarily more solute than they normally would

• A small event will cause this solution to precipitate..like pouring

Solution

• The rate of solution: how fast a substance dissolves

• Factors that effect the rate of solution• The lab

Molarity

• Ya…more moles• The molarity of a solution is definedas the

number of moles of solute per liter of solution• The unit is M (capital m)• So a 3.0 molar solution of nitric acid is

abbreviated 3.0M HNO3

• That means the solution contains 3.0 moles of nitric acid per liter of solution

Molarity

• The molarity of a solution is calculated by taking the moles of solute and dividing by the liters of solution.

Molarity

• Example #1 - Suppose we had 1.00 mole of sucrose (it's about 342.3 grams) and proceeded to mix it into some water.

• It would dissolve and make sugar water. We keep adding water, dissolving and stirring until all the solid was gone.

• We then made sure that when everything was well-mixed, there was exactly 1.00 liter of solution.

Molarity

• What would be the molarity of this solution?

• 1.0 M• Easy right?

Molarity

• What is the molarity when 0.75 mol is dissolved in 2.50 L of solution?

• 0.300 M

Molarity

• Suppose you had 58.44 grams of NaCl and you dissolved it in exactly 2.00 L of solution.

• What would be the molarity of the solution?• Step One: convert grams to moles. • Step Two: divide moles by liters to get

molarity. • 1mol/2.00L= 0.5M

Chemical Kinetics

• The rates of chemical reactions• The collision theory states that in order for a

chemical reaction to take place particles must collide

• Not only must reactants collide, they must collide effectively

• That means that they have to hit at the appropriate angle and stick together (magnets)

Chemical Kinetics

Chemical Kinetics

• Four factors that affect the rate of a chemical reaction:

1.The nature of reactants2.The concentration of reactants3.The temperature4.The use of catalysts

Chemical Kinetics

• The nature of reactants:• Ionic substances usually react very quickly

while covalents not so much• Concentration: Usually the more concentrated

the reactants are the faster the reaction will occur

• Temperature: An increase in temperature will always increase the rate of a reaction

Chemical Kinetics

• Activation energy: the amount of energy needed to create an activated complex

• Activated complex is the in-between stage that only lasts for a brief moment where reactants have effectively collided and are on their way to becoming products.

• Catalysts: speed up or lower the activation energy without being used up or altered

Potential Energy Diagrams


• These show the relationship between the activation energy (pushing the boulder up) and the energy absorbed or given off during a reaction (the boulder rolling down)

• Sometimes the reactants have more energy than the products and sometimes the products have more energy than the reactants.

so what is chemistry? the study of matter matter pure substances elements na, h, o compounds...

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