Chemistry Unit VIII The Mole Concept

Transcription

1 Chemistry Unit VIII The Mole Concept PRE-TEST QUESTIONS 1. What is the mass of a carbon atom? 2. What is the SI unit of mass? 3. What does the prefix kilo- mean? 4. What two particles determine atomic mass? 5. What is the formula of potassium chlorate? I. The Mole A. Measuring Matter 1. The amount of a substance is measured by one of three different methods: a. By b. By c. By 2. You re familiar with many terms regarding numbers: a. A pair means. b. A dozen means. c. A baker s dozen means. d. A gross is. e. A score is. B. What is a mole? 1. It is impractical to count individual atoms. 2. Many properties depend on the number of molecules in the sample, not on the mass of the sample. How can the number of molecules in the sample be measured? 3. A is defined as the amount of matter that contains the same number of particles as grams of. 4. This number, called Avogadro s number, has been found to be. a. This is the number of, be they atoms, molecules, ions, formula units, or whatever, in a mole of any substance. b. A mole of any substance weighs differently from any other substance. II. The Math of Moles A. Avogadro s Number 1. To convert the number of moles of any substance into representative particles, or viceversa, use one of the following conversion factors: EXAMPLE VIII-01:Moles Atoms 1. How many atoms of copper are in 1.50 moles of copper? Hendley UNIT VIII Notes, Page 1

2 2. How many atoms of helium does 3.50 moles of helium contain? EXAMPLE VIII-02:Atoms Moles 1. How many moles are atoms of silver? atoms of nickel are equal to how many moles? EXAMPLE VIII-03:Molecules and Atoms 1. How many molecules of oxygen are in 12.5 moles of oxygen? (Remember: Oxygen is a diatomic molecule, meaning the formula is O 2.) 2. How many atoms of oxygen are in those 12.5 moles of oxygen? B. Molar Mass 1. In the laboratory, we cannot weigh in units of moles. However, we can weigh mass in. 2. The atomic mass of an element, when expressed in, is the of a of the element. a. A helium atom weighs approx amu i. A mole of helium weighs. b. A carbon-12 atom weights approx amu. i. A mole of carbon-12 weighs. 3. The molar mass of an element is the mass of a mole of that element. a. Also referred to as molar weight, and several other similar names. b. Measured in units of. c. For the purposes of this class, round to 2 places after the decimal. (This is because the reference table Periodic Table usually goes two after the decimal) Hendley UNIT VIII Notes, Page 2

3 EXAMPLE VIII-04:Molar Mass of Elements 1. What is the molar mass of neon? 2. How many grams would 3.25 moles of neon weigh? C. Molar Mass of a Compound 1. To find the mass of a mole of a compound, you must know the formula of the compound. 2. Best get to it EXAMPLE VIII-05:Molar Mass of a Compound 1. What is the molar mass of water? 2. What is the molar mass of ammonia? 3. What is the molar mass of sodium chloride? 4. What is the molar mass of carbon dioxide? 5. What is the molar mass of copper(ii) sulfate? 6. What is the molar mass of calcium nitrate? Hendley UNIT VIII Notes, Page 3

4 D. Mole-Mass Relationship 1. Now that you know how to calculate the molar mass, you will use this the rest of the year. 2. Use the molar mass to convert between units of moles and grams. a. Always do the conversion factors for 1 mole. EXAMPLE VIII-06: Moles Mass 1. What is the mass of 7.00 mol NaCl? 2. What is the mass of 2.50 mol of water? EXAMPLE VIII-07: Mass Moles 1. How many moles is 45.2 grams of CuSO 4? 2. A mass of 90.1 grams of water is how many moles? E. Mole-Volume Relationship 1. Solids and liquids have molar volumes that tend to vary quite a bit, because their particles are close together. Thus, small changes in size have big changes in volume. a. Gases, however, have fairly volumes because their particles are far apart, so the size, or mass, has little impact on the space they occupy. 2. Avogadro s Principle, stated by Amedeo Avogadro, states that equal volumes of gases at the same temperature and pressure contain numbers of particles. a. Avogadro s Principle is true at temperature and pressure, as long as they are kept. b. Avogadro s number was named in his honor. 3. Standard Temperature and Pressure (STP) refers to the specific temperature of ( ) and the pressure of ( = = ). a. This is a standardized temperature and pressure for measuring many other quantities that may vary due to changes in these values. 4. At STP conditions, 1 mole of any gaseous substance, regardless of what it is, occupies a volume of. a. This is called standard molar volume. Hendley UNIT VIII Notes, Page 4

5 EXAMPLE VIII-08: Moles Volume 1. What volume will moles of CO 2 gas occupy at STP? EXAMPLE VIII-09: Volume Moles 1. How many moles is 45.0 liters of O 2 gas at STP? F. Calculating Molar Mass from Density 1. Though a mole of different gases might occupy the same volume at the same temperature and pressure, they do have different masses and thus, different densities. 2. The density of a gas is usually stated in grams per liter (g/l) [instead of per ml like solids and liquids are], along with a specific. 3. You can determine the density of a gas at STP using the molar mass and molar volume: EXAMPLE VIII-10: Gas density calculations 1. What is the molar mass of a gas that has a density of 4.52 g/l at STP? 2. What is the density of CO 2 at STP? G. The Mole Road Map 1. You have now examined moles in terms of: a. (Avogadro s Number) b. (molar mass) c. (molar volume at STP) 2. Remember that to convert between any of those units, it must be in terms of moles. III. Percent Composition and Empirical Formulas A. Percent Composition of a Compound 1. The percent composition is the percent, by mass, of each element in a compound. Hendley UNIT VIII Notes, Page 5

6 2. To find the percent composition, take the mass of each individual element and divide it by the total molar mass of the entire compound, then multiply by to get the percentage. EXAMPLE VIII-11: Percent Composition from Formula 1. What is the percent composition of K 2 CrO 4? 3. Percent Composition from Mass a. Sometimes, chemists perform experiments and get actual numbers and do not know the formula, necessarily. EXAMPLE VIII-12: Percent Composition from Mass 1. A sample of 50.0 grams of water is broken up into hydrogen and oxygen gas. There are 5.60 grams of hydrogen liberated. What is the percent composition of this compound? B. Empirical Formulas 1. When chemists determine the formula for compounds, they always find it by the simplest ratio between compounds. 2. The empirical formula is the formula with the lowest whole-number ratio of elements in a compound. a. For example, the empirical formula of N 2 O 4 is NO 2. b. An empirical formula may or may not be the same as a molecular formula. i. CO 2 is both the empirical and molecular formula. (It can t be simplified more) 3. There are four steps to solving for an empirical formula. EXAMPLE VIII-13: Empirical Formula 1. A compound is 78% Boron and 22% Hydrogen. What is this compound s empirical formula? a. Assume 100 grams total. Hendley UNIT VIII Notes, Page 6

7 b. Convert grams into moles using molar mass. c. Divide all moles by the smallest value. d. Round it off by 0.1 if more than that, multiply by some constant to balance. EXAMPLE VIII-14: Empirical Formula 2 1. A compound is composed of 26.56% Potassium, 35.41% Chromium, and 38.03% Oxygen. Determine its empirical formula. a. Assume 100 grams total. b. Convert grams into moles using molar mass. c. Divide all moles by the smallest value. d. Round it off by 0.1 if more than that, multiply by some constant to balance. Hendley UNIT VIII Notes, Page 7

8 C. Molecular Formulas 1. The molecular formula (the formula of the actual molecule) is either the same as the empirical formula, or it is a simple, whole-number multiple of the empirical formula. 2. For example, most sugars share the empirical, but glucoses molecular formula is C 6 H 12 O 6. a. Similarly, vinegar is made of acetic acid, CH 3 COOH, which is C 2 H 4 O To determine the molecular formula from the empirical formula, chemists use a tool called a mass spectrometer. 4. To solve a molecular formula, you will be given the mass of the molecule. Determine the empirical formula, take its mass, and divide the two. a. This gives you the whole number ratio to multiply the empirical formula by in order to get the true molecular formula. EXAMPLE VIII-15: Molecular Formula 1 1. Consider Example VIII-13 above. Suppose the molar mass of the molecular compound was determined to be amu. Determine the molecular formula. EXAMPLE VIII-16: Molecular Formula 2 1. Hydrogen peroxide has an empirical formula of HO and a molecular mass of amu. Find its molecular formula. Hendley UNIT VIII Notes, Page 8