Chapter 7 Chemical Quantities

1. Chapter 7Chemical Quantities Fall 2008

2. The Mole: A Measurement of Matter-What Is a Mole? • How do you measure matter? • You count things • You weigh things or determine mass • You get the volume • Can you count the grains of sand? Would you want to? • How do we determine the amount of something big or small without having to count it?

3. The Mole: A Measurement of Matter-What Is a Mole? • We use problem solving steps to figure out the amount of very small or very big forms of matter. That method is called…… Dimensional Analysis

4. Dimensional Analysis • Here is an example: • How many eggs can Bug Bunny buy for cupcakes for Elmer Fudd’s birthday if he has $10 to buy eggs with at $0.80 cent for a dozen? • First start with what you know • 12 eggs = 1 dozen, 6 eggs = ½ dozen • Eggs cost ¢80 for 1 dozen • Bugs has $10

5. Dimensional Analysis • Next we set up Conversion factors • 12 eggs/ 1 doz. Or 1 doz./12 eggs • $0.80/ 1doz. Or 1 doz./$0.80 • Now we can set up the problem • $10 1 doz. 12 eggs = 150 eggs $ 0.801 doz.

6. What is a Mole? • But now back to the mole… • In chemistry, it is not the furry little creature… • It is the measurement of the amount of particles in a substance • Particles can be atoms, ions, formula units or molecules

7. What is a Mole? • So, how much is a mole….. 6.0223 x 1023 particles • Particles are extremely small and that number is extremely big. • This is due to Amadeo Avagadro’s research leading to the discovery, not the discovery of the number itself.

8. What is a Mole? • It has also been discovered that 1 mole of all gases at a certain set of conditions called STP (standard temperature and pressure) have the same volume—22.4L! • So, 1 mole = 22.4L (gas @ STP)

9. Mass and Moles • How do we usually measure out chemicals in a lab? • NOT by counting the number of particles • We usually use mass to measure out chemicals. The problem with mass is that not all atoms or compounds have the same mass. • Avogadro’s number was established by using Carbon-12 as a standard. The number of particles in 12 grams of carbon 12 was set as 1 mole.

10. Mass and Moles • Conveniently, this is similar to how the atomic mass unit was established, so the atomic mass on the periodic table is also an element’s molar mass (the mass of one mole of an element, measured in g/mol). • Since compounds are just atoms bonded together, you can find the molar mass of a compound by adding together the molar mass of all the elements in the compound.

11. Mass and Moles • We have now discussed moles in three ways—as particles (1 mol = 6.02x1023 particles), as a mass ( molar mass = 1 mol) and as a volume (22.4 L = 1 mol) • To get from particles to grams, grams to liters or particles to liters of a substance, two steps are required, with moles always in between!

12. The Mole Map Gas Volume @ STP # Particles  x 22.4L 6.02 x 1023  x Mole  Molar Mass x Mass

13. Mole Road Map • Here is a graphic of the procedure steps: • Pick the box of the data you are given in the problem and follow the steps toward the box containing what you are asked for in the problem.

14. % Composition • The percent composition of an element is the relative amount of that element in a compound. • You find the percent composition of an element by dividing the element’s mass by the mass of the entire compound. mass of element x 100 mass of compound

15. For example: • What is the % composition of carbon dioxide? Mass of oxygen = 2(16.0)=32.0 g Mass of carbon = 12.0 g Mass of carbon dioxide = 44.0 g 12.0 g C 44.0 g CO2 32.0 g O 44.0 g CO2 27.3% Carbon x100 = x100 = 72.7% Oxygen

16. % Water in a Hydrate • A hydrate is an ionic compound with water as part of its crystalline structure. • To find the percent of water in a hydrate, use the same formula, except: • % Water = x 100 Mass water Mass compound

17. Empirical Formula • The empirical formula gives you the lowest, whole number ratio of elements in the compound. • The empirical formula may or may not be the same as the molecular formula. • For carbon dioxide, the molecular formula is CO2, and the empirical formula is CO2. One carbon and two oxygens are the lowest ratio of atoms. • The molecular formula for dinitrogen tetrahydride is N2H4, but the empirical formula is NH2.

18. What is the empirical formula for: • C6H12O6 • N2H2 • CH4

19. To Find the Empirical Formula from % Composition: • If given the percentages, assume there are 100.0 grams of the compound. • Convert the grams of each element to moles. • Divide by the smaller amount of moles, then manipulate the ratio so that all numbers are whole. • What is the empirical formula of a compound that is 27.3% carbon and 72.7% oxygen? • What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen?

20. To Find the Empirical Formula from % Composition: • If given the percentages, assume 100% = 100.0 grams . • Convert the grams of each element to moles. • Divide by the smaller amount of moles, then manipulate the ratio so that all numbers are whole.

21. From there, find it’s molecular formula: • If given the molar mass (how many g/mol of the compound) then you can calculate the molecular formula from the empirical formula. • Take the compound’s empirical formula mass and compare to the molecular mass. The molecular mass will be a multiple of the empirical formula’s mass.