Warm Up

• What is a mole?

• What is molar mass?

• What is Avogadro’s number?

Chapter 7

The Mole and Chemical

Compostion

How can chemical composition be determined?

Unit Essential Question:

How is the mole used in conversions?

Lesson Essential Question:

Section 1: Avogadro’s Number and Molar Conversions• 1 mole = 6.022 x 1023 particles

−SI unit for amount of substance.•It’s a counting unit (like a dozen).

−Remember that the unit of particles can be: ions, molecules (mcs.), atoms, formula units (f.u.), etc.

Recall that formula units = simplest ratio of ions in an ionic compound.

covalent compounds

ionic compounds

Recall your mole map!

Converting moles particles• Same as Chapter 3, but it will involve

molecules, formula units, or ions instead

of just atoms.

• Steps:1) Need 1mol = 6.022 x1023 molecules, etc.2) Use dimensional analysis- turn this into a fraction!*Be sure to place the correct units on the top and bottom so they cancel!

Sample Problems 1 & 2: Moles & Particles

• Find the number of molecules in 2.5 mol of sulfur dioxide.

• A sample contains 3.01 x 1023 molecules of sulfur dioxide. Determine the amount in moles.

1.5 x 1024 molecules SO2

0.500mol SO2

Molar Mass• Amount of mass (in grams) in 1 mole of a

substance.

• Use molar masses from the periodic table.−Round to 2 decimal places!−Use units of g/mol.−Example:

• C: 12.01g/mol means that 1 mol C = 12.01 g• Cl: 35.45g/mol means that 1 mol Cl = 35.45g

• Use to convert between moles and mass.

Sample Problems 3 & 4: Moles & Mass

• What is the mass of 5.3mol Be?

• If you have 27.0g of manganese, how many moles do you have?

48g Be

0.491mol Mn

Molar Masses of Compounds• Add together the molar masses of all

elements or ions present.−Ex: CH4

−C: 12.01g/mol H: 1.01g/mol−12.01g/mol + 4(1.01g/mol) = 16.05g/mol

−This means that 1 mole of CH4 has a mass of 16.05g.

• You will need to calculate the molar mass

of a compound whenever you are

converting between mass and moles!

Additional Molar Mass Examples:• Element

−Ag = 107.87 g/mol

• Diatomic Element/molecule−Br2 = 79.90 x 2 = 159.80 g/mol

• Molecule (Covalent compound)−H2O = (1.01 x 2) + 16.00 = 18.02 g/mol

• Formula unit (Ionic compound)−Ca(NO3)2 = 40.08 + (2 x 14.01) + (6 x

16.00) = 164.10 g/mol

Sample Problem 5: Mass to Moles with a Compound

• Find the number of moles present in 47.5 g of glycerol, C3H8O3.

• Hint: you will need to calculate the molar mass of glycerol!

Glycerol’s molar mass: 92.11g/mol

0.516mol C3H8O3

Sample Problem 6: Number of Particles to Mass

• Remember- you can’t go directly between mass (g) and the number of particles! You must convert to moles first!

• Find the mass in grams of 2.44 x 1024 atoms of carbon.

48.7 g C

More Practice

• How many moles of iron (III) sulfate,

Fe2(SO4)3, are there in a 178g

sample?

0.445mol

How are molar masses on the periodic table

determined?

Lesson Essential Question:

Mole Ratios in Chemical Formulas• Ratios can be formed between amounts of elements or ions within

a compound.−Look at the subscripts.

• Example #1: CaCl2

−For every 1mol of CaCl2 there is 1mol of Ca+2 ions and 2mol of Cl- ions.

• Example #2: Na2CO3

−For every 1mol of Na2CO3, there are 2mol of Na+ ions and 1mol of CO3

-2 ions.

• Example #3: N2O3

−For every 1mol of N2O3 there are 2mol of N atoms and 3mol of O atoms.

Practice

• If you have one mole of strontium

cyanide, Sr(CN)2, how many moles of

strontium ions are there? How many

moles of cyanide ions are there?

• Given the compound P2O5 what is the

mole ratio of P atoms to O atoms?

Section 2: Relative Atomic Mass and Chemical

Formulas• Periodic table masses are averages of

all isotopes present. −Recall that we said a weighted average is

used- takes into account the amount of each isotope.

−Average atomic mass: (% x atomic mass)+(% x atomic mass)+…

100−Note: % is the percent abundance (how

often the element is found as that isotope in nature).

Sample Problem• The mass of a Cu-63

atom is 62.94 amu, and that of a Cu-65 atom is 64.93 amu. If the abundance of Cu-63 is 69.17% and the abundance of Cu-65 is 30.83%, what is the average atomic mass of copper?

What information can be determined from formulas?

How can formulas be determined?

Lesson Essential Questions:

Calculating Percent Composition

• Tells you the percent each element

makes up of the whole compound.Step 1: Determine the molar mass of the entire compound.Step 2: Divide each element’s total molar mass by the molar mass of the compound.Step 3: Multiply by 100 to get percent.Step 4: Check your answer by adding up the percentages to makes sure they equal 100%.

Percent Composition Cont.

• Calculating the percent composition of

a compound can be helpful in

determining the formula/identity.

• Example:−Iron and oxygen form two compounds:

•Fe2O3 and FeO

−Fe2O3 = 69.9% Fe and 30.1% O

−FeO = 77.7% Fe and 22.3% O

Sample Problem #I• Calculate the percent composition of

copper (I) sulfide.

• Calculate the percent composition of isopropyl alcohol, (CH3)2CHOH.

Sample Problem #2

Determining Empirical Formulas

• The empirical formula shows the simplest ratio of elements/ions in the compound.−Ionic compounds are represented with empirical

formulas.

• Given percent composition data, you can determine the empirical formula of a compound.Step 1: Assume 100 g of the sample- put ‘g ’ in

for ‘% ’. Ex: 18.2% O 18.2gStep 2: Convert grams to moles.Step 3: Divide each mole value by the smallest

mole value. This will tell you the number of each element that appears in the formula.

Step 4: If you get a decimal, multiply ALL numbers by a whole number to turn the decimal into a whole number.

•The numbers you will need to multiply by should be relatively small (2, 3, etc.)

Determining Empirical Formulas Cont.

Sample Problem #1• Chemical analysis of a liquid shows

that it is 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula of this substance.

• A compound is found to contain 38.77% Cl and 61.23% O. What is the empirical formula?

Sample Problem #2

Molecular Formulas• Show the actual numbers of elements in the

compound- not necessarily the simplest

formula.−Often seen for covalent compounds.

• They will be a whole number multiple of the

empirical formula (can’t be a decimal).−In other words: n(empirical formula) = molecular formulawhere n is a whole number.

−Ex: 6(CH2O) C6H12O6

• Molecular and empirical formulas can be the

same!

Molecular Formulas Cont.

Molecular Formulas Cont.• The molecular formula can be determined from

the empirical formula and experimental molar

mass of a compound.Step 1: Determine the molar mass of the given empirical formula.Step 2: Solve for n by dividing the experimental molar mass by the molar mass of the empirical formula.*Remember: n(empirical formula) = molecular formula

Step 3: Multiply the subscripts in the empirical formula by n.

Sample Problem #1• The empirical formula for a compound is

P2O5. Its experimental molar mass is 284g/mol. Determine the molecular formula of the compound.

• A brown gas has the empirical formula NO2. Its experimental molar mass is 46g/mol. What is the molecular formula?

Sample Problem #2

Hydrates- Honors Only

• Not in the textbook.

• Hydrates – ionic compounds that

contain water molecules within the

crystal structure.−Example: CuSO4•5H2O

−Anhydrous – without the water = CuSO4

Determining Hydrate Formulas

• Formula can be determined if given: the

mass of the hydrate, the anhydrous

mass, and the formula of the ionic

compound.Step 1: Determine the mass of water in the hydrate (subtract anhydrous mass).Step 2: Convert the anhydrous ionic compound mass and water mass to moles.Step 3: Divide both molar amounts by the smallest number. This gives you the number of water molecules in the hydrate.

Sample Problem #1

• A 5.82 g sample of Mg(NO3)2· XH2O in an evaporating dish is heated until it is dry.  The mass of the anhydrous sample is 2.63 g Mg(NO3)2.  What is the formula for the hydrate?

Determining % Water in a Hydrate

• Formula can be determined if given

the formula of the hydrate.Step 1: Calculate the mass of the entire hydrate and the mass of just the water. Step 2: Divide the mass of the water by the mass of the entire hydrate and multiply by 100 to get a percent.

• What percentage, by mass, of water is found in the hydrate CuSO4·5H2O?

Sample Problem #2