CHAPTER 6: CHEMICAL QUATITIES

6.1 Measuring Matter

1. MOLE: SI unit that measures "amount of substance".

2. MOLE:

measures # particles, mass, volume.

6.2 The Mole

1. MOLE: = 6.02 X 10²³ representative particles of that substance - AVAGADRO'S NUMBER

2. REPRESENTATIVE PARTICLE - refers to the species present in a substance: atoms (elements), molecules (diatomic molecules or molecular compounds, or formula units (ionic compounds).

Examples: elements: Fe, Ca, C, H₂, O₂

Molecules: H₂0, CO₂

Formula Units: NaCl, CaCl2

3. A mole of each of these substances contains 6.02 X 1023 particles.

4. How many atoms are in a mole of a compound? Example CO₂

Answer: 3 x Avagadro's Number

5. To find the number of atoms in a mole of a compound, you must determine the number of atoms in a representative formula of that compound. See example 2, page 146; Example 3, page 147.

6. Reminder - Steps in problem solving.

Unknown

Known

Conversion Units

Calculations

Finish - S.F., S.N.

6.3 The Gram Formula Mass

1. GRAM ATOMIC MASS (gam): Atomic mass of an element expressed in grams. Larger than AMUs. Use grams.

2. AMU is really "relative" amount calculated with weighted average masses as compared to C-12.

3. GRAM ATOMIC MASS of any two elements must contain the same number of atoms. Example: 12.0 g of C has same # atoms as 16.0 g O, which is Avagadro's Number = Mole.

4. MOLE: defined as the amount of substance that contains as many representative particles as the number of atoms in 12.0 g of carbon-12. (Atomic mass)

5. The gram atomic mass (gam) is the mass of one mole of atoms of any element.

6. GRAM MOLECULAR MASS (gmm) of any molecular compound is the mass of one mole of that compound. Calculate by adding atomic masses of a given formula.

Example: 1 mol SO₃ = 80.1 g

7. Do practice problems, 13 & 14.

8. The GRAM FORMULA MASS (gfm) is the mass of 1 mol of an ionic compound. Calculate the same as gmm.

6.4 The Molar Mass of a Substance. 1. MOLAR MASS is a term which refers to all: elements, molecules, ionic compounds gam, gmm, gfm.

2. Gfm can also be used in a general way. Do problems 16, 17.

6.5 Mole-mass Conversions.

1. Gfm is used to convert moles to grams and grams to moles:

Grams<---------->Moles

2. Calculate gfm first to covert grams to moles or moles to grams. Use Gfm/mole as conversion unit. See eg. 6&7, problem 18.

6.6 Volume of mole of Gas

1. Volume of gas is measured at a standard temperature and pressure (STP).

• Standard temperature = O^oC
• Standard pressure = 101.3 kPa or 1 atmosphere (atm).

2. 1 mole of any gas occupies volume of 22.4 L. (Molar Volume)

3. 1 mole of a gas has 6.02 x 10²³ particles and a volume of 22.4 L at STP.

Moles<------->Volume

4. Example 8,9; problems 20.22.

6.7 Gas Density and the Gram Molecular Mass

1. Can determine Gfm using experimentally determined density in g/L. Start with mole of gas, i.e., 22.4 L and use density as conversion factor. Find Gfm, then compare with calculated gfms.

2. See Example 10.

6.8 Converting Between Units with Mole

1. To change from one unit to another, use mole as intermediate step.

2. Conversion units: 1 mol/1gfm; 1mol/22.4 l gas; 1 mol/6.02 x 10²³ particles. 1gfm/6.02 x 10²³

3. See example 11, 24, 26, 27

6.9 Calculating Percent Composition

1. To determine the formula of a compound, we need to know the relative amounts of elements.

2. The relative amounts are expressed as the percent composition, which is the percent by mass of each element in the compound.

3. Each element in the compound has its own percent, which together at up to 100%.

4. %mass of element = grams of element/grams of compound X 100%.

5. Example 12, 28. Use given masses in grams.

6. May also use formula to calculate gfm, which = mole of compound.

7. Calculate % of each element.

8. % mass = grams of element in 1 mole/gfm of compound x 100%

9. Example 13, 29

10. Use % composition to calculate number of grams of an element in a compound. Make a conversion factor based on % composition.

11. Example 14, 30.

6.10 Calculatin Empirical Formulas

1. May calculate % composition experimentally, then calculate empirical formula.

2. Empirical formula gives the lowest whole number ratio of the atoms of the elements in a compound.

3. Empirical formual may or may not be the same as the molecular formula. If not it is a simple multiple of empirical formula.

4. Gives relative counts of count of atoms or moles of atoms in molecules or formula units.

5. May be viewed on microscopic level----> molecule, or macroscopic ----> mol.

5. E.g., H₂O₂ -> HO: N₂H₄ -> NH₂.

6. Example 15, 33, 34.

7. Use % composition as grams in 100g. Convert to moles, using conversion factors with mole of atom. Devide # moles of each atom by one of the # moles. Multiply if not in whole number ratios.

6.11 Calculating Molecular Formulas

1.The molecular formula of a compound is either the same as its experimentally determined empirical formula, or it is some simple whole-number multiple of it.

2. If you know empirical formula and gfm, you can determine molecular formula.

Determine empirical mass (efm).

Devide by know gfm. This gives appropriate multiple.

3. Example 16, 36, 37