Chapter10Mole(12-13)_1.ppt

Chapter 10

10.1 Measuring Matter

What do you ask for when you buy:

2 shoes

12 eggs

48 donuts

1 pair

1 dozen

Pair, Dozen, Ream

These are all ways to batch a

group of objects to make them

easier to count!

The object may change

Eggs to Donuts

What do eggs have to do with

Chemistry?

How many carbon atoms are in a teaspoon of carbon?

200,666,666,666,666,666,666,667

atoms!!!

It would be nice if chemists

had a batch like a dozen (but

Meet the Mole!

Counting Atoms

• Chemistry is a quantitative science - we need a "counting unit."

The chemist’s “dozen” is called the:

MOLE (or the unit mol)

• The MOLE_MOLE

How Big is a Mole?

One mole of marbles would cover the entire Earth (oceans included) for a depth of three miles.

One mole of $100 bills stacked one on top of another would reach from the Sun to Pluto and back 7.5 million times.

11.1 Measuring Matter – Moles and

Avogadro’s number

• NOTES START HERE!!

• Mole (mol) – the amount of a substance that contains the same number of particles, 6.02x1023 representative

particles, as the number of atoms in 12 g of carbon-12. • Like donuts are counted in dozens, the mole is a SI unit

for counting the amount of a substance.

– 1 dozen pencils contains the same number of particles as 1 dozen donuts

– 1 mole of carbon atoms contains the same number of particles as 1 mole of water molecules

This does not mean they both weigh the same,only that

10.1 Measuring Matter – Moles and

Avogadro’s number

• Avogadro’s Number – the number of

representative particles (6.02  1023) in exactly

one mole of a pure substance.

– 1 mole pencils = 6.02  1023 pencils

10.1 Measuring Matter – Moles and

Avogadro’s number

If Avogadro’s number is the number of particles in 1 mole, how do you know what kind of particle you have?

Remember…

10.1 What kind of particle?

Particles can be an

atom

,

molecule

, or

formula unit

?

• Atom – one atom

• Molecule – more than one nonmetal atom • Formula unit - compound with a metal or

10.1 Measuring Matter – Moles and

Avogadro’s number

How do you know what kind of particle you have?

Examples

1. NaCl 2. H₂O 3. H₂ 4. Na

formula unit molecule molecule

10.1 Measuring Matter – Moles and

Avogadro’s number

Avogadro’s number is really an equality!

1 mole C = 6.021023 atoms C

What can you use equalities to do?

Stoichiometry Island Diagram

Mass

Particles

Volume _{Mole Mole}

Mass

Volume

Particles

Known Unknown

Substance A Substance B 1 m

ole = _mo

lar m ass

(g) _on PT

Use coefficients from balanced chemical equation 1 mole = 22.4 L @ STP

1 mole = 6

.02 x 1023 pa

rticles

(atom s, m

olecu les, o

r form ula un

its)

1 mole = 22.4 L @ STP

1 m ole

= _6.0 2 x ₁₀

23 pa rtic_les (ato ms_,m ole cule s, o

r fo_rm ula

un its) 1 mole

= m olar m

ass ( g) on

PT

(gases) (gases)

Atoms Form Formu_la Atoms

10.1 Measuring Matter – Moles and

Avogadro’s number : Problem #1

How many and what type of particles are in 2.0 moles of copper?

Cu mol

2.0

= atoms Cu

mol Cu

atoms Cu

1

6.021023

10.1 Measuring Matter – Moles and

Avogadro’s number: Problem #2

How many atoms are in 3.2 mol of C?

C mol

3.2 _{= atoms C}

mol C

atoms C

1

6.021023

10.1 Measuring Matter – Moles and

Avogadro’s number: Problem #3

How many moles of nitrogen are in 2.78  1023 molecules of nitrogen?

= mol N₂

mol N₂

molecules N₂ 1

6.021023

.462

10.1 Measuring Matter – Moles and

Avogadro’s number: Problem #4

How many moles of water molecules are in 3.7  1021 molecules of H

2O?

= mol H₂O

mol H₂O

molecules H₂O 1

6.021023

.0061

3.7  1021 _{molecules H}

The Formula Conversion Factor

Think about what a formula really tells you!

1 molecule of H₂O

Doesn’t this mean:

1 molecule H₂O = 2 hydrogen atoms

The Formula Conversion Factor

How many hydrogen atoms are in 6.5x1014

molecules of water?

atoms H

molecule H₂O 2

1

6.5x1014 molecules H

2O

10.2-10.3 Mass and the Mole

• Molar mass – a term used to refer to the

mass, in grams, of one mole of a substance.

• How do you find this mass?

The atomic mass printed on the periodic table has two meanings –

1.It is the average mass of an atom in atomic mass units (amu)

10.2-10.3 Mass and the Mole

• Finding molar mass

– Atoms – the mass of 1 mole of any atom is the same as the atomic mass in grams.

– Round all molar masses to the tenths place.

– The unit for molar mass is g/mol.

1mole H atoms =1.00794 g

1mole C =

 1.0 g/mol

12.0107 g  12.0 g/mol

10.2-10.3 Mass and the Mole

• Molar mass continued

– Compounds – the mass of 1 mole of a

compound is the sum of the masses of the atoms.

1mole H₂O = 2 H =

1 O =

2(1.0g) = 2.0g 16.0g

10.2-10.3 Mass and the Mole

Molar mass is really an equality!

For water:

1mole H₂O = 18.0 g

What can you use equalities to do?

10.2-10.3 Mass and the Mole

Molar Mass Problem #1

How many moles are in 242 g of water?

mol H₂O

g H₂O 1

18.0

13.4

242 g H₂O _{= mol H}

10.2-10.3 Mass and the Mole

Molar Mass Problem #2

What is the mass of 3.77 mol of Au?

g Au

mol Au 197.0

1

743

3.77 mol Au

10.2-10.3 Mass and the Mole

Multi-step problem

How many water molecules are in 242 g of water?

mol H₂O

g H₂O 1

18.0

8.09x1024

242 g H₂O

=

molecules H₂O

1 mol H₂O

6.021023 _{molecules H}

10.4 Percent Composition

How would you calculate the percent females in this room?

All percents are calculated in the same way! 100

× people #

females #

100 ×

10.4 Percent Composition

• Percent Composition – the percent by

mass of each element in a compound.

100 ×

compound mass

element mass

molar mass

What is the percent composition of each element in water?

Mass H =

Mass O = 16.0 g

10.4 Percent Composition

2(1.0 g) = 2.0 g

+ 18.0 g 100 × O H mass H mass = H % 2

100

×

g

18.0

g

2.0

=

Using Percent Composition

If a glass of water contains 648 g of water, how many grams of hydrogen would it

hold?

Remember, water is 11% hydrogen

What is 11% of 648 g?

= (.11)

× g)

(648

71.28 g

Dimensional Analysis: Another way

to solve for the previous problem

If a glass of water contains 648 g of water, how many grams of hydrogen would it hold?

g H

g H₂O 2.0

18.0 72

648 g H₂O

10.4 Empirical Formula

What does H₂O mean?

Does it mean 2 atoms of H for every atom of O?

YES

Does it mean 2 g of H for every 1 g of O?

10.4 Empirical Formula

What does H₂O mean?

Does it mean 2 moles of H for every mole of O?

Always!

Formulas are ratios of moles!

Form

ulas a

re rat

ios of

mole s!

Formulas are ratios of moles!

Fo rm ula s a re r atio s o f m ole s!

10.4 Empirical Formula

• Empirical Formula: simplest whole

number ratio of moles of the atoms in a substance. Experimental method that is the first step in finding the formula of a compound.

Circle the empirical formulas!

H₂SO₄ H₂O

N₂O₄

C₂H_{6 NO2}

10.4 Empirical Formula

Finding the empirical formula

1. Find the mass of each element in the compound.

– Usually given

– If given as %, then change % to g.

75.0% C and 25.0% H  75.0 g C and 25.0 g H

2. Convert masses to moles.

– Use molar masses.

75.0 g C

= mol C g C

mol C 1

12.0

6.25 25.0 g H

= mol H g H

mol H 1

1.0

(Check animation)

10.4

Empirical Formula

3. Find the smallest whole number ratio of moles.

a. Write the results of step 2 like a formula. If C = 6.25 mol and H = 25 mol

C_6.25H₂₅

b. Divide by the smallest mole amount. In this example, we divide by 6.25

C_6.25H₂₅

c. If not all whole numbers, multiply by 2,3, or 4 ends in .5, multiply by 2

ends in .3, multiply by 3 ends in .25, multiply by 4

10.4 Empirical Formula

Examples for step 3

X = .029 mol X = .009 mol

Y = .039 mol Y = .006 mol

X_.029Y_.039

.029 .029

X₁Y_1.3

Multiply by 3

X₃Y₄

X_.009Y_.006

.006 .006

X_1.5Y₁

Multiply by 2

10.4 Empirical Formula

More Examples for step 3

X = 1.47 mol X = 2.4  10-4

mol

Y = 3.68 mol Y = 7.3  10-4

mol X_{1.47 1.47}1.47Y3.68

X₁Y_2.5

Multiply by 2

X₂Y₅

2.410-4 _2.4₁₀-4

X₁Y_3.04

X₁Y₃

4 4

- _7.3×10

10 × 2.4

10.4 Empirical Formula - Example

A sample of an unknown gas contains 43.2 g of carbon and 115.8 g of oxygen. What is the empirical formula? 1. Find Masses

43.2 g C 115.8 g O 2. Change to moles

43.2 g C

= mol C g C

mol C 1

12.0

3.60

115.8 g O

= mol O g O

mol O 1

16.0

10.4 Empirical Formula - Example

3. Get whole numbers C_3.60O_7.24

3.60 3.60

C_1.00O_2.01

10.4 Molecular Formula

• Molecular formula - is some whole number

multiple of the empirical formula.

– HO is an empirical formula – H₂O₂ is twice HO

– (HO)_X and X = 2

– For C₆H₁₂O₆, the empirical formula is CH₂O and X=6

• To convert an empirical formula to a molecular formula you must find X.

mass empirical

mass molar

10.4 Molecular Formula - Example

An unknown gas is found to have an empirical formula of NO₂ and a molar mass of 92.0 g/mol. What is the molecular formula?

mass empirical

mass molar

= X

g/mol 46.0

g/mol 92.0

=

X =2

Class Practice of Empirical and

Molecular Formulas

• 7.36 grams of a compound decomposes into 6.93 grams of oxygen and the rest is hydrogen. The molar mass is 34.0 g/mol.

First, find the empirical formula.

1. Find the mass of each element in the compound.

– Total Mass: 7.36g – Mass of O: 6.93g

– Mass of H : 7.36g-6.93g= 0.43gH

2. Convert masses to moles.

– Use molar masses.

0.43 g H

= mol H g H

mol H 1

1.0

0.43 6.93 g O

= mol O g O

mol O 1

16.0

3. Find the smallest whole number ratio of moles.

H_0.43O_.433

.43 .43

HO mass empirical mass molar = X g/mol 17.0 g/mol 34.0 =

X =2

Class Practice of Empirical and

Molecular Formulas

• Find the molecular formula if you are given the empirical formula and the molar mass. HgCl, 472.2 g/mol

mass empirical

mass molar

= X

g/mol 236.1

g/mol 472.2

=

X =2

Water of Hydration

(or Water of Crystallization) • Hydrate: is a compound that has a

specific number of water molecules

bound to its atoms.

• The water can be driven off by heating:

• CuSO₄

.

• Called copper(II)sulfate pentahydrate- heat . + heat

Naming Hydrates

• For hydrate formulas:

– The number of water molecules associated with each formula unit of the compound is written following a dot

Naming Hydrates

– For the name, write the name of salt compound followed by the prefix of the number of water molecules and the word hydrate.

• Example: Na₂CO₃10H₂O

• Example: sodium carbonate decahydrate • Prefixes are the same as prefixes used in

Class Practice of Determining the

Formula for a Hydrate

A mass of 2.50 g of blue, hydrated copper(II) sulfate (CuSO₄xH₂O) is placed in a crucible and heated. After heating, 1.59 g white

anhydrous copper(II) sulfate (CuSO₄) remains. What is the formula for the hydrate?

First, find the empirical formula.

1. Find the mass of anhydrous salt and water in compound. – Known:

• Total Mass of hydrate:

• Mass of anyhdrous compound:

• Molar mass of H₂O =

• Molar mass of CuSO₄ =

2.50 g CuSO₄xH₂O 1.59 g CuSO₄

18.0 g/mol H₂O

159.6 g/mol CuSO₄

To find the mass of water, subtract the hydrate from the anhydrous compound.

Mass of anyhdrous compound:

1.59 g CuSO₄

0.91 g CuSO₄

1. Convert masses to moles.

– Use molar masses.

1.59 g CuSO₄

= mol CuSO₄ g CuSO₄

mol CuSO₄ 1

159.6

0.00996

0.91 g H₂O

= mol H₂O g H₂O

mol H₂O 1

18.0

3. Find the smallest whole number ratio of moles.

CuSO_40.00996H₂O_0.051

0.00996 0.00996

CuSO₄ (H₂O)₅

Formula: CuSO₄xH₂O = CuSO₄5 H₂O