The Mole - Weebly€¦ · mole Avogadro’s number • Main Idea - Chemists use the mole to count...

The MoleCh 11

Measuring matter11.1

11.1 Vocabulary• Review

o molecule: two or more atoms that covalently bond together to

form a unit

• Newmole

Avogadro’s number

• Main Idea - Chemists use the mole to count atoms,

molecules, ions, and formula units.

• NOTE – YOU WILL NEED A SCIENTIFIC

CALCULATOR FOR THIS CHAPTER!

How do we measure items?

You can measure mass,

or volume,

or you can count pieces.

We measure mass in grams.

We measure volume in liters.

We count pieces in MOLES.

What is the mole?

We’re not talking about this

kind of mole!

Counting Particles

• Chemists need a convenient method for accurately

counting the number of atoms, molecules, or

formula units of a substance.

• The mole is the SI base unit used to measure the

amount of a substance.

Moles (is abbreviated: mol)

It is an amount, defined as the number of carbon atoms in exactly 12 grams of carbon-12.

1 mole = 6.022 x 1023 of the representative particles.

Treat it like a very large dozen!

6.022 x 1023 is called Avogadro’s number.

Similar Words for an amount

Pair: 1 pair of shoelaces = 2 shoelaces

Dozen: 1 dozen oranges = 12 oranges

Case: 1 case of Dr. Pepper

= 24 cans Dr. Pepper

Gross: 1 gross of pencils = 144 pencils

Ream: 1 ream of paper = 500 sheets of paper

What are Representative Particles?

The smallest pieces of a substance:

1) For a molecular compound: it is

the molecule.

2) For an ionic compound: it is the

formula unit (made of ions).

3) For an element: it is the atom.

• Remember the 7 diatomic

elements? (made of molecules)

Types of questions• How many oxygen atoms in the

following?

CaCO3

Al2(SO4)3

• How many ions in the following?

CaCl2

NaOH

Al2(SO4)3

3 atoms of oxygen

12 (3 x 4) atoms of oxygen

3 total ions (1 Ca2+ ion and 2 Cl1- ions)

2 total ions (1 Na1+ ion and 1 OH1- ion)

5 total ions (2 Al3+ + 3 SO42- ions)

Converting Between Moles and Particles

• Conversion factors must be used.

• Moles to particles

Example:

Number of molecules in 3.50 mol of sucrose

Converting Between Moles and Particles (cont.)

• Particles to moleso Use the inverse of Avogadro’s number as the conversion factor.

Practice problems (round to 3 sig. figs.)

How many molecules of CO2 are in 4.56 moles of CO2?

How many moles of water is 5.87 x 1022

molecules?

How many atoms of carbon are in 1.23 moles of C6H12O6?

How many moles is 7.78 x 1024 formula units of MgCl2?

2.75 x 1024 molecules

0.0975 mol (or 9.75 x 10-2)

4.44 x 1024 atoms C

12.9 moles

10.1 CheckWhat does the mole measure?

A. mass of a substance

B. amount of a substance

C. volume of a gas

D. density of a gas

10.1 CheckWhat is the conversion factor for determining the

number of moles of a substance from a known

number of particles?

A.

B.

C. 1 particle 6.02 1023

D. 1 mol 6.02 1023 particles

Mass and the mole11.2

11.2 Vocabulary• Review

o conversion factor: a ratio of equivalent values used to express the same

quantity in different units

• Newo molar mass

• Main Idea - A mole always contains the same

number of particles; however, moles of different

substances have different masses.

The Mass of a Mole• 1 mol of copper and 1 mol of carbon have different

masses.

• One copper atom has a different mass than 1

carbon atom.

Measuring Moles Remember relative atomic mass?

- The amu was one twelfth the mass of

a carbon-12 atom.

Since the mole is the number of atoms

in 12 grams of carbon-12,

the decimal number on the periodic table

is also the mass of 1 mole of those atoms in

grams.

The Mass of a Mole (cont.)• Molar mass (MM) is the mass in grams of one mole of

any pure substance.

• Also called Formula Weight (FW)

Molar Mass

Equals the mass of 1 mole of an element in

grams (from periodic table)

12.011 grams of C has the same number of

pieces as

• 1.008 grams of H

• 55.85 grams of iron.

We can write this as: 12.011 g C = 1 mole C

We can count things by weighing them.

Using Molar Mass• Moles to mass

Using Molar Mass (cont.)• Convert mass to moles with the inverse molar mass

conversion factor.

Examples How much would 2.34 moles of carbon

weigh?

How many moles of magnesium is 24.31 g of Mg?

How many atoms of lithium is 1.00 g of Li?

How much would 3.45 x 1022 atoms of U weigh?

28.1 grams C

1 mol Mg

8.72 x 1022 atoms Li

13.6 grams U

Using Molar Mass (cont.)• This figure shows the steps to complete conversions

between mass and atoms.

11.2 CheckThe mass in grams of 1 mol of any pure substance is:

A. molar mass

B. Avogadro’s number

C. atomic mass

D. 1 g/mol

11.2 CheckMolar mass is used to convert what?

A. mass to moles

B. moles to mass

C. atomic weight

D. particles

Moles of compounds11.3

Vocabulary• Review

o representative particle: an atom,

molecule, formula unit, or ion

• Main Idea -The molar mass of a compound can be

calculated from its chemical formula and can be

used to convert from mass to moles of that

compound.

Chemical Formulas and the Mole

• Chemical formulas indicate the numbers and types

of atoms contained in one unit of the compound.

• One mole of CCl2F2 contains one mole of C atoms,

two moles of Cl atoms, and two moles of F atoms.

The Molar Mass of Compounds

• The molar mass of a compound equals the molar

mass of each element, multiplied by the moles of

that element in the chemical formula, added

together.

• The molar mass of a compound demonstrates the

law of conservation of mass.

The Molar Mass of Compounds in 1 mole of H2O molecules there are two

moles of H atoms and 1 mole of O atoms (think of a compound as a molar ratio)

To find the mass of one mole of a compound

odetermine the number of moles of the elements present

oMultiply the number times their mass(from the periodic table)

oadd them up for the total mass

Calculating Molar Mass

Calculate the molar mass of

magnesium carbonate, MgCO3.

24.3050 g + 12.0107 g + 3 x (15.9994 g)

= 84.3139

Thus, 84.33139 grams is the formula

mass for MgCO3.

Examples Calculate the molar mass of the

following and tell what type it is:

Na2S

N2O4

C

Ca(NO3)2

C6H12O6

(NH4)3PO4

= 78.045 g/mol

= 92.011 g/mol

= 12.011 g/mol

= 164.088 g/mol

= 180.156 g/mol

= 149.087 g/mol

Moles to Mass Conversion for Compounds

• For elements, the conversion factor is the molar

mass of the elements.

• The procedure is the same for compounds, except

that you must first calculate the molar mass of the

compound.

Mass to Moles Conversion for Compounds

• The conversion factor is the inverse of the molar

mass of the compound.

37

For exampleHow many moles is 5.69 g of NaOH?

38

For example

How many moles is 5.69 g of NaOH?

5 69. g

39

For example

How many moles is 5.69 g of NaOH?

5 69. g mole

g

We need to change 5.69 grams NaOH to moles

40

For example

How many moles is 5.69 g of NaOH?

5 69. g mole

g

We need to change 5.69 grams NaOH to moles

1mole Na = 22.990 g

1 mol O = 15.999 g

1 mole of H = 1.008 g

41

For example

How many moles is 5.69 g of NaOH?

5 69. g mole

g

We need to change 5.69 grams NaOH to moles

1mole Na = 22.990 g

1 mol O = 15.999 g

1 mole of H = 1.008 g

1 mole NaOH = 39.997 g

42

For example

1 mole NaOH = 39.997 g NaOH

= 0.1422606696 mol NaOH

= 0.142 mol NaOH 3 Sig Figs

Mass to Particles Conversion for Compounds

• Convert mass to moles of compound with the

inverse of molar mass.

• Convert moles to particles with Avogadro’s number.

• This figure summarizes the conversions between

mass, moles, and particles.

11.3 Check

How many moles of OH— ions are in 2.50 moles of

Ca(OH)2?

A.2.00

B. 2.50

C.4.00

D.5.00

11.3 Check

How many particles of Mg are in 10 moles of MgBr2?

A.6.02 1023

B. 6.02 1024

C.1.20 1024

D.1.20 1025

Empirical and Molecular formulas

11.4

Vocabulary

• Review

o percent by mass: the ratio of the mass of each element to

the total mass of the compound expressed as a percent

• New

percent composition

empirical formula

molecular formula

• Main Idea - A molecular formula of a compound is

a whole-number multiple of its empirical formula.

Percent Composition

• The percent by mass of any element in a

compound can be found by dividing the mass of

the element by the mass of the compound and

multiplying by 100.

Percent Composition

• The percent by mass of each element in a

compound is the percent composition of a

compound.

• Percent composition of a compound can also be

determined from its chemical formula.

50

Calculating Percent Composition

of a Compound

Like all percent problems:

part

whole

1) Find the mass of each of the

components (the elements),

2) Next, divide by the total mass of

the compound; then x 100%

x 100 % = percent

51

Example

Calculate the percent composition of a

compound that is made of 29.0 grams of

Ag with 4.30 grams of S

(Assume you have one mol of substance)

29.0 g Ag

33.3 g totalX 100 = 87.1 % Ag

4.30 g S

33.3 g totalX 100 = 12.9 % S

Total = 100 %

52

Examples

Calculate the percent

composition of C2H4?

How about Aluminum

carbonate?

85.7% C, 14.3 % H

23.1% Al, 15.4% C, and 61.5 % O

Empirical Formula• The empirical formula for a compound is the smallest

whole-number mole ratio of the elements.

• You can calculate the empirical formula from percent by mass by assuming you have 100.00 g of the compound. Then, convert the mass of each element to moles.

o SO3 has a percent composition of 40.05% S and 59.95% O.

o To find mole ratio assume 100 g of Sulfur Trioxide; this means 40.05 g S and 59.95 g O. Multiply each by molar mass to determine moles.

o 1.249 mol S and 3.747 mol O. Turn this into a whole number ratio by dividing by the smallest. S becomes 1 and O becomes 3

o [If you still don’t get whole number, multiply by smallest number that will produce a whole number. If you had 1.5 mol C, 3 mol H, 1 mol O, multiply all by 2 to get 3 mol C, 6 mol H, 3 mol O = C3H6O2

• The empirical formula may or may not be the same as the molecular formula.

Molecular formula of hydrogen peroxide = H2O2

Empirical formula of hydrogen peroxide = HO

Molecular Formula• The molecular formula specifies the actual number

of atoms of each element in one molecule or

formula unit of the substance.

• Molecular formula is always a whole-number

multiple of the empirical formula.

• To distinguish the molecular formula from empirical

formula, chemists must experimentally determine

the molar mass of the compound. Then, this

relationship is used: o Molecular formula = (empirical formula)n

• n is the factor by which molecular formula is obtained.

o Unknown substance X has formula weight of 26.04 g/mol and its empirical

formula is CH. Determine the molecular formula.

o (26.04g/mol)/ (13.02 g/mol) = 2

o ANSWER – C2H2

Empirical vs. Molecular FormulaWhat is the empirical

formula for the

compound C6H12O6?

A.CHO

B. C2H3O2

C.CH2O

D.CH3O

Which is the empirical

formula for hydrogen

peroxide?

A.H2O2

B. H2O

C.HO

D.none of the above

Formulas of hydrates11.5

Vocabulary• Review

crystal lattice: a three-dimensional geometric

arrangement of particles

• New

hydrate

Main Idea -Hydrates are solid ionic compounds in

which water molecules are trapped.

Naming Hydrates• A hydrate is a compound that has a specific

number of water molecules bound to its atoms.

• The number of water molecules associated with

each formula unit of the compound is written

following a dot.

• Sodium carbonate decahydrate =

Na2CO3 • 10H2O

Naming Hydrates

Analyzing Hydrates• When heated, water molecules are released from a

hydrate leaving an anhydrous compound.

• To determine the formula of a hydrate, find the

number of moles of water associated with 1 mole of

hydrate.

Analyzing Hydrates (cont.)

1. Weigh hydrate.

2. Heat to drive off the water.

3. Weigh the anhydrous compound.

4. Subtract and convert the difference to moles.

5. The ratio of moles of water to moles of anhydrous

compound is the coefficient for water in the

hydrate.

Use of Hydrates• Anhydrous forms of hydrates are often used to

absorb water, particularly during shipment of

electronic and optical equipment.

• In chemistry labs, anhydrous forms of hydrates are

used to remove moisture from the air and keep

other substances dry.

11.5 CheckHeating a hydrate causes what to happen?

A.Water is driven from the hydrate.

B. The hydrate melts.

C.The hydrate conducts

electricity.

D.There is no change in the

hydrate.

11.5 CheckA hydrate that has been heated and the water driven

off is called:

A.dehydrated compound

B. antihydrated compound

C.anhydrous compound

D.hydrous compound

11.5 CheckTwo substances have the same percent by mass

composition, but very different properties. They must

have the same ____.

A.density

B. empirical formula

C.molecular formula

D.molar mass