Chapter 7 Chemical Quantities

1
Chapter 7Chemical Quantities

• Charles Page High School
• Dr. Stephen L. Cotton

2
Section 7.1The Mole A Measurement of Matter

• OBJECTIVES
• Describe how Avogadros number is related to a
  mole of any substance.

3
Section 7.1The Mole A Measurement of Matter

• OBJECTIVES
• Calculate the mass of a mole of any substance.

4
What is a Mole?

• 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.

5
Moles (abbreviated mol)

• Defined as the number of carbon atoms in exactly
  12 grams of carbon-12.
• 1 mole is 6.02 x 1023 particles.
• Treat it like a very large dozen
• 6.02 x 1023 is called Avogadros number.

6
Representative particles

• The smallest pieces of a substance.
• For a molecular compound it is the molecule.
• For an ionic compound it is the formula unit
  (ions).
• For an element it is the atom.
• Remember the 7 diatomic elements (made of
  molecules)

7
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

8
Types of questions

• How many molecules of CO2 are there in 4.56 moles
  of CO2 ?
• How many moles of water is 5.87 x 1022
  molecules?
• How many atoms of carbon are there in 1.23 moles
  of C6H12O6 ?
• How many moles is 7.78 x 1024 formula units of
  MgCl2?

9
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.

10
Gram Atomic Mass (gam)

• Equals the mass of 1 mole of an element in grams
• 12.01 grams of C has the same number of pieces as
  1.008 grams of H and 55.85 grams of iron.
• We can write this as 12.01 g C 1 mole C
• We can count things by weighing them.

11
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?

12
What about compounds?

• in 1 mole of H2O molecules there are two moles of
  H atoms and 1 mole of O atoms
• To find the mass of one mole of a compound
• determine the moles of the elements they have
• Find out how much they would weigh
• add them up

13
What about compounds?

• What is the mass of one mole of CH4?
• 1 mole of C 12.01 g
• 4 mole of H x 1.01 g 4.04g
• 1 mole CH4 12.01 4.04 16.05g
• The Gram Molecular Mass (gmm) of CH4 is 16.05g
• this is the mass of one mole of a molecular
  compound.

14
Gram Formula Mass (gfm)

• The mass of one mole of an ionic compound.
• Calculated the same way as gmm.
• What is the GFM of Fe2O3?
• 2 moles of Fe x 55.85 g 111.70 g
• 3 moles of O x 16.00 g 48.00 g
• The GFM 111.70 g 48.00 g 159.70 g

15
Section 7.2Mole-Mass and Mole-Volume
Relationships

• OBJECTIVES
• Use the molar mass to convert between mass and
  moles of a substance.

16
Section 7.2Mole-Mass and Mole-Volume
Relationships

• OBJECTIVES
• Use the mole to convert among measurements of
  mass, volume, and number of particles.

17
Molar Mass

• Molar mass is the generic term for the mass of
  one mole of any substance (in grams)
• The same as 1) gram molecular mass, 2) gram
  formula mass, and 3) gram atomic mass- just a
  much broader term.

18
Examples

• Calculate the molar mass of the following and
  tell what type it is
• Na2S
• N2O4
• C
• Ca(NO3)2
• C6H12O6
• (NH4)3PO4

19
Molar Mass

• The number of grams of 1 mole of atoms, ions, or
  molecules.
• We can make conversion factors from these.
• To change grams of a compound to moles of a
  compound.

20
For example

• How many moles is 5.69 g of NaOH?

21
For example

• How many moles is 5.69 g of NaOH?

22
For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles

23
For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH

24
For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g

25
For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g
• 1 mole NaOH 40.00 g

26
For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g
• 1 mole NaOH 40.00 g

27
For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g
• 1 mole NaOH 40.00 g

28
Examples

• How many moles is 4.56 g of CO2?
• How many grams is 9.87 moles of H2O?
• How many molecules is 6.8 g of CH4?
• 49 molecules of C6H12O6 weighs how much?

29
Gases

• Many of the chemicals we deal with are gases.
• They are difficult to weigh.
• Need to know how many moles of gas we have.
• Two things effect the volume of a gas
• Temperature and pressure
• We need to compare them at the same temperature
  and pressure.

30
Standard Temperature and Pressure

• 0ºC and 1 atm pressure
• abbreviated STP
• At STP 1 mole of gas occupies 22.4 L
• Called the molar volume
• 1 mole 22.4 L of any gas at STP

31
Examples

• What is the volume of 4.59 mole of CO2 gas at
  STP?
• How many moles is 5.67 L of O2 at STP?
• What is the volume of 8.8 g of CH4 gas at STP?

32
Density of a gas

• D m / V
• for a gas the units will be g / L
• We can determine the density of any gas at STP if
  we know its formula.
• To find the density we need the mass and the
  volume.
• If you assume you have 1 mole, then the mass is
  the molar mass (from PT)
• At STP the volume is 22.4 L.

33
Examples

• Find the density of CO2 at STP.
• Find the density of CH4 at STP.

34
The other way

• Given the density, we can find the molar mass of
  the gas.
• Again, pretend you have 1 mole at STP, so V
  22.4 L.
• m D x V
• m is the mass of 1 mole, since you have 22.4 L of
  the stuff.
• What is the molar mass of a gas with a density of
  1.964 g/L?
• 2.86 g/L?

35
Summary

• These four items are all equal
• a) 1 mole
• b) molar mass (in grams)
• c) 6.02 x 1023 representative particles
• d) 22.4 L at STP
• Thus, we can make conversion factors from them.

36
Section 7.3Percent Composition and Chemical
Formulas

• OBJECTIVES
• Calculate the percent composition of a substance
  from its chemical formula or experimental data.

37
Section 7.3Percent Composition and Chemical
Formulas

• OBJECTIVES
• Derive the empirical formula and the molecular
  formula of a compound from experimental data.

38
Calculating Percent Composition of a Compound

• Like all percent problems
• Part whole
• Find the mass of each component,
• then divide by the total mass.

x 100
39
Example

• Calculate the percent composition of a compound
  that is 29.0 g of Ag with 4.30 g of S.

40
Getting it from the formula

• If we know the formula, assume you have 1 mole.
• Then you know the mass of the pieces and the
  whole.

41
Examples

• Calculate the percent composition of C2H4?
• How about Aluminum carbonate?
• Sample Problem 7-11, p.191
• We can also use the percent as a conversion
  factor
• Sample Problem 7-12, p.191

42
The Empirical Formula

• The lowest whole number ratio of elements in a
  compound.
• The molecular formula the actual ratio of
  elements in a compound.
• The two can be the same.
• CH2 is an empirical formula
• C2H4 is a molecular formula
• C3H6 is a molecular formula
• H2O is both empirical molecular

43
Calculating Empirical

• Just find the lowest whole number ratio
• C6H12O6
• CH4N
• It is not just the ratio of atoms, it is also the
  ratio of moles of atoms.
• In 1 mole of CO2 there is 1 mole of carbon and 2
  moles of oxygen.
• In one molecule of CO2 there is 1 atom of C and 2
  atoms of O.

44
Calculating Empirical

• We can get a ratio from the percent composition.
• Assume you have a 100 g.
• The percentages become grams.
• Convert grams to moles.
• Find lowest whole number ratio by dividing by the
  smallest.

45
Example

• Calculate the empirical formula of a compound
  composed of 38.67 C, 16.22 H, and 45.11 N.
• Assume 100 g so
• 38.67 g C x 1mol C 3.220 mole C
  12.01 gC
• 16.22 g H x 1mol H 16.09 mole H 1.01
  gH
• 45.11 g N x 1mol N 3.219 mole N 14.01
  gN

46
Example

• The ratio is 3.220 mol C 1 mol C
  3.219 molN 1 mol N
• The ratio is 16.09 mol H 5 mol H
  3.219 molN 1 mol N
• C1H5N1
• A compound is 43.64 P and 56.36 O. What is
  the empirical formula?
• Caffeine is 49.48 C, 5.15 H, 28.87 N and
  16.49 O. What is its empirical formula?

47
Empirical to molecular

• Since the empirical formula is the lowest ratio,
  the actual molecule would weigh more.
• By a whole number multiple.
• Divide the actual molar mass by the empirical
  formula mass.
• Caffeine has a molar mass of 194 g. what is its
  molecular formula?

48
Example

• A compound is known to be composed of 71.65 Cl,
  24.27 C and 4.07 H. Its molar mass is known
  (from gas density) to be 98.96 g. What is its
  molecular formula?
• Sample Problem 7-14, p.194