Honors ChemistryPart 1

Unit 4 – The Mole

WHAT IS A MOLE?

602214199000000000000000

6.02 x 1023

Mole Facts 6.02 X 1023  Pennies: Would make at least

7 stacks that would reach the moon. 6.02 X 1023  Watermelon Seeds: Would be

found inside a melon slightly larger than the moon.

6.02 X 1023  Blood Cells: Would be more than the total number of blood cells found in every human on earth.

1 Liter bottle of Water contains 55.5 moles H20

Definition of Mole

The amount of atoms in 12.0 grams of Carbon 12 (6.02 x 1023 atoms known as Avogadro’s number).

A sample of any element with a mass equal to that element's atomic weight (in grams) will contain precisely one mole of atoms (6.02 x 1023 atoms).

Molecular Mass

The sum of the masses of all the atoms in a molecule of a substance

CaCO3

1 atom of Ca = 40.08 amu1 atom of C = 12.00 amu3 atoms of O = 48.00 amu 100.08 amu

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Formula Mass (Molar Mass) The mass of 1 mole (in grams) Equal to average atomic mass but the

unit is grams1 mole of C atoms = 12.01 g1 mole of Na atoms = 22.99 g1 mole of Cu atoms = 63.55 g

Example Problem

Find the mass of 1 mole of KAl(SO4)2 ● 12H2O

1 K = 39.101 Al = 26.982 (SO4) 2(32.06 + ((16.00 x 4))=192.12

12 H2O 12(2.02 + 16.00)=216.24

Mass of 1 mole = 474.44 g/mol

Try These:

Find the molecular mass for these : HNO3

CO2

Find the molar mass for these compounds:

C6H10O5

H2SO4

The Mole

1 mole of gas always contains 6.02 x 1023 molecules of that gas

1 mole Cl2 gas = 6.02 x 1023 molecules of Cl2 1 mole NO2 gas = 6.02 x 1023molecules of NO2

1 mole CO gas = 6.02 x 1023 molecules of CO 1 mole CO2 gas = 6.02 x 1023molecules of CO2

The Mole

Also applies to other particles! (not only molecules in a gas)

1 mole C = 6.02 x 1023 C atoms 1 mole H2O = 6.02 x 1023 H2O molecules 1 mole NaCl = 6.02 x 1023 NaCl formula

units 1 mole of Na+ = 6.02 x 1023 Na+ ions 1 mole of Cl- = 6.02 x 1023 Cl– ions

Avogadro’s Number

We can use Avogadro’s # as a conversion factor:

1 mole 6.02 x 1023 particles

Or6.02 x 1023 particles

1 mole Note that a particle could be an atom,

molecule, formula unit, or ion !

Example Problems

How many molecules are in 3.5 moles of H2O?

How many moles are present in 4.65 molecules of NO2?

Mass and Mole Relationship 1 mole of any substance = the molar

mass of that substance (in grams) Find the number of moles present in

56.7 g of HNO3.

56.7 g HNO3 1 mole HNO3

63.01 g HNO3

Example Problems

Find the number of grams present in 4.5 moles of C6H10O5.

Find the number of moles present in 12.31 g of H2SO4.

How many molecules are in 4.5 grams of NaCl?

Gas Volumes and Molar Mass Avogadro’s Law Equal volumes of gases under the same

conditions of temperature and pressure contain equal numbers of molecules

1 mole of any gas at standard temperature and pressure (STP) occupies 22.4 liters

Standard temperature: 0ºC or 273K Standard pressure: 1 atm or 101.325 kPa

Gas Volumes and Molar Mass 32.00 g O2 = 1 mole = 22.4 L

2.02 g H2 = 1 mole = 22.4 L

44.01 g CO2 = 1 mole = 22.4 L

Example Problems

How many liters are present in 5.9 moles of O2?

How many liters are present in 3.67 moles of CO2?

How many atoms of O are present in 78.1 g of O2?

Percent Composition

Finding what percent of the total weight of a compound is made up of a particular element

Formula for calculating % composition:

Total mass of the element in the compound

Total formula mass

X 100

Example Problem:

Calculate the % composition of calcium in Ca(OH)2.

Example Problems

Find the percentage composition of a compound that contains 1.45 g of carbon, 4.23 g of sulfur, and 1.00 g of hydrogen in a 6.68 g sample.

21.7% carbon 63.3% sulfur 15.0% hydrogen

Example Problem

A sample of an unknown compound with a mass of 5.00 grams is made up of 75% carbon and 25% hydrogen. What is the mass of each element?

3.75 g of carbon 1.25 g of hydrogen

Formulas

Empirical Formula - expresses the smallest whole number ratio of atoms present

E.g. CH2O

Ionic formulas are always empirical formulas

Molecular Formula - states the actual number of each kind of atom found in one molecule of the compound.

E.g. C6H12O6

Empirical Formula

1. Determine mass in grams of each element

2. Calculate the number of moles of each

3. Divide each by the smallest number of moles to obtain the simplest whole number ratio

4. If whole numbers are not obtained in step 3, multiply all by the smallest number that will give whole numbers

Empirical Formula

Remember this: Percent to mass Mass to mole Divide by small Multiply ‘till whole

Example Problem

Given that a compound is composed of 60.0% Mg and 40.0% O, find the empirical formula.

Example Problem

A compound is found to contain 68.5% carbon, 8.63% hydrogen, and 22.8% oxygen. The molecular weight of this compound is known to be approximately 140.00 g/mol. Find the empirical and molecular formulas.

Hydrates

Ionic compounds Water is bonded to the crystal structure Ex: CuSO3 • 7H2O

The percentage of water in a hydrate can easily be calculated using the formula:

% Water = Mass of water x 100 Mass of hydrate

Example Problem

What is the percentage of water in CuSO3 • 7H2O?

A 3.5 g sample of a hydrate is heated and only 1.7 g of the anhydrous salt remain. What is the percentage of water?

Law of Definite Proportions Formulas give the numbers of atoms or

moles of each element Always a whole number ratio 1 molecule NO2 : 2 atoms of O for every

1 atom of N 1 mole of NO2 : 2 moles of O atoms to

every 1 mole of N atoms

Law of Multiple Proportions When any two elements, A and B,

combine to form more than one compound, the different masses of B that unite with a fixed mass of A have a small whole-number ratio

Example: In H2O, the proportion of H:O = 2:16 or

1:8 In H2O2, H:O is 2:32 or 1:16

How Do We Determine Concentration?

Molarity Molality

How do we make solutions?

M1 = m1/V1 rearrange to M1V1 = m1

M2 = m2/V2 rearrange to M2V2 = m2

If m1=m2

then, M1V1 = M2V2

M1V1 = M2V2

M1 = concentration of the first solutionV1 = volume of the first solutionM2 = concentration of the second solutionV2 = volume of the second solution

Let's consider a sample problem: You have 1 L of a 0.125 M aqueous solution of

table sugar.  You want to dilute the solution to 0.05 M.  What do you do?

Dilution

To solve the problem, you simply plug in the three numbers you know:

(0.125 M) (1 L) = (0.05 M) V2

2.5 L = V2

Using the equation, you determine that the volume of the diluted solution should be 2.5 L. 

So we simply add enough water to the first solution so that the solution's volume becomes 2.5 L.

What is Saturation? A solution is saturated if it contains as much solute as can possibly

be dissolved under the existing conditions of temperature and pressure

Unsaturated: Has less than maximum amount of solute that can be dissolved

Supersaturated: Contains more than maximum (How can this happen?)