Chapter 6

1. Chapter 6 Chemical Quantities

2. Introduction • We live in a quantitative world • Questions such as “how many?” “how much?” “how far?” and “how long?” are common. • Chemistry is a quantitative science • A common question in chemistry would be, “How many grams of the elements hydrogen and nitrogen must be combined to make 200g ammonia, NH3?”

3. 6.1 Measuring Matter • Matter is often measured in different ways: • Counting → There are 26 desks in the classroom. • Some terms refer to a specific number: a pair = 2; a dozen = 12 • Mass (or weight) → The book has a mass of 1 kg • Volume → The beaker contains 250 ml of water

4. 6.1 Measuring Matter • Using equivalent measurements to write conversion factors, you can convert between these different units of measurement. • Example: Apples Equivalent Measurements: • 1 dozen apples = 12 apples (count) Using the average size of apples you can determine that: • 1 dozen apples = 2 kg apples (mass) • 1 dozen apples = 0.20 bushel of apples (volume) Conversion Factors: or… 2 kg apples 1 dozen apples

5. 6.1 Example 1 • What is the mass of 90 average-size apples? Unknown: Known: 90 apples 1 dozen apples = 12 apples 1 dozen apples = 2.0 kg apples Plan: Convert: number of apples → dozens of apples → mass of apples mass of apples

6. 6.1 Example 1 Solution: Final Answer: 15 kg apples 180 kg apples 90 apples x 1 dozen apples x 2.0 kg apples = 1 12 apples 1 dozen apples 12 = 15 kg apples

7. 6.2 The Mole • The mole is the SI unit that measures the “amount of a substance” • The unit mole can be related to the number of particles (count), the mass, and the volume of an element or compound * Just like a dozen was used in the apple example

8. 6.2 The Mole • Mater is composed of different kinds of particles • atoms, molecules, ions • Since these particles are so small, even the smallest sample of a substance would have a very large number of particles • It is not practical or possible to count individual particles unless you use a term to represent a specific number of particles – that unit is the mole

9. 6.2 The Mole • A mole of a substance represents 6.02 x 1023representative particles of that substance • 6.02 x 1023 is called Avogadro’s number and it is a constant (it stays the same) representative particle–refers to the species present in a substance: usually atoms, molecules, or formula units (ions)

10. 6.2 The Mole • The representative particle of most elements is the atom. • Example: iron is composed of iron atoms • The representative of diatomic molecules (H2, N2, O2, F2, Cl2, Br2, and I2) and all molecular compounds is the molecule. • The formula unit is the representative particle of ionic compounds. * See Table 6.1, page 145

11. 6.2 Example 2 • How many moles of magnesium are 3.01 x 1022 atoms of magnesium? Unknown: Number of moles of Mg Known: 3.01 x 1022 atoms of Mg 1 mole = 6.02 x 1023 atoms

12. 6.2 Example 2 Solution: Final Answer (in Scientific Notation): 5.00 x 10-2 mol Mg 3.01 x 1022 atoms Mg x 1 mol Mg = 1 6.02 x 1023 atoms Mg 3.01 x 1022 mol Mg = 0.500 x 10-1 mol Mg = 6.02 x 1023

13. 6.2 The Mole • To determine how many atoms are in a mole of a compound you need to know how many atoms are in a representative unit of that compound (a molecule or formula unit). • Example: A single molecules of CO2 has 1 atom of carbon and 2 atoms of oxygen • The number of atoms (3) in a mole of CO2 (6.02 x 1023) is three times greater than Avogadro’s number → 3 x 6.02 x 1023