Unit 4 Notepack
Chapter 7 Chemical Quantities
Qualifier for Test
NAME____________________
Section 7.1 The Mole: A Measurement of Matter
A.
What is a mole?
1. Chemistry is a quantitative science. What does this term mean?
2. How do we measure quantities of matter?
a.
-examples:
b.
-examples:
c.
-examples:
3. Some of the units we use to measure matter indicates specific numbers:
a. pair –
b. dozen c. gross –
d. ream –
4. The mole is just like this. Define mole:
B.
The Number of Particles in a Mole
1. Counting atoms, ions, molecules, and formula units is impractical. They’re too small!
2. We use a counting unit, called the mole, to count these representative particles.
3. How many “representative particles” are contained in 1 mole?
4. This is an experimentally determined number. It is called _____________________, in honor of this Italian
scientist.
Sample Calculations:
1. How many moles of magnesium is 1.25 x 1023 atoms of magnesium?
2. How many moles is 2.80 x 1024 atoms of silicon?
3. How many molecules is 0.360 mole of water?
4. How many moles are equal to 2.41 x 1024 formula units of sodium chloride (NaCl)?
**Now suppose you want to determine how many atoms are in a mole of a compound.
For example, each molecule of carbon dioxide (CO 2) is composed of three atoms: One carbon and two oxygen
atoms.
Sample Calculation:
1. How many atoms are in 2.12 moles of propane (C3H8)?
2. How many atoms are there in 1.14 mole of sulfur trioxide?
3. How many moles are there in 4.65 x 1024 molecules of nitrogen dioxide?
4. How many atoms of Carbon are in 2.0 moles of C12H22O11, sucrose sugar?
C.
The Mass of a Mole of an Element
1. Define gram atomic mass –
a. What is the gram atomic mass of carbon?
b. What is the gram atomic mass of hydrogen?
c. What is the gram atomic mass of sulfur?
** This is the amount of an element needed to weigh out exactly one mole of atoms of that element. (6.02 x
1023 atoms)
D.
The Mass of a Mole of a Compound
1. Define gram molecular mass –
a. What is the gram molecular mass of sulfur trioxide?
b. What is the gram molecular mass of hydrogen?
c. What is the gram molecular mass of carbon dioxide?
**This is the amount of a molecular compound needed to weigh out exactly one mole of molecules of that
compound. (6.02 x 1023 molecules)
2. Define Gram formula mass –
a. What is the gram formula mass of sodium chloride?
b. What is the gram formula mass of ammonium carbonate?
c. What is the gram formula mass of potassium oxide?
** This is the amount of an ionic compound needed to weigh out exactly one mole of formula units of that
compound. (6.02 x 1023 formula units)
Section One Review Problems:
1. Find the gram formula mass of each compound:
a. lithium sulfide
b. iron (III) chloride
c. calcium hydroxide
2. How many oxygen atoms are in a representative particle of each substance?
a. ammonium nitrate
b. acetylsalicylic acid (C9H8O4), the chemical name of aspirin
c. ozone (O3), a disinfectant and natural molecule found in the atmosphere
d. nitroglycerine (C3H5(NO3)3), an explosive
3. How many moles in each of the following:
a. 1.50 x 1023 molecules of ammonia, NH3?
b. 1 billion (1 x 109) molecules of oxygen, O2?
c. 6.02 x 1022 molecules of bromine, Br2?
d. 4.81 x 1024 atoms of lithium, Li?
7.2 Mole-Mass and Mole-Volume Relationship
A. The Molar Mass of a Substance
1. Define molar mass –
2. There are situations that this term doesn’t work well for…
Ex: What is the molar mass of oxygen?
What would be better?
Under most situations, this general term works fine… just be careful for the seven naturally occurring diatomic
molecules.
List those here:
3. We use molar mass to convert between grams and moles of a substance
Example: How many grams are in 9.45 mol of dinitrogen trioxide?
Find the number of moles in 92.2 grams of iron (III) oxide
Sample Calculations:
1. Find the mass, in grams of each.
a. 3.32 mole of potassium atoms, K.
b. 4.52 x 10-3 mole C6H12O6
c. 0.0112 mole of potassium carbonate
2. Calculate the number of moles in 75 grams of each substance.
a. dinitrogen trioxide
b. nitrogen gas
c. sodium oxide
B. The Volume of a Mole of Gas
1. Unlike liquids and solids, the volumes of moles of gases are much more predictable under the same physical
conditions of temperature and pressure.
2. Since the volume of a gas varies directly with its temperature, and indirectly with its pressure, we have
“standard conditions” we use to report volumes of gases.
3. What is standard temperature and pressure (STP)?
4. One mole of any gas at standard conditions occupies ____________Liters.
5. This quantity is called the __________________ __________________.
Examples:
Determine the volume, in liters, of 0.60 moles of sulfur dioxide gas at STP.
Determine the number of moles of oxygen gas in 11.5 L at STP.
Sample Calculations:
1. What is the volume at STP of these gases?
a. 3.20 x 10-3 mol carbon dioxide
b. 0.960 mol methane, CH4
c. 3.70 mol nitrogen gas
2. Assuming STP, how many moles are in these volumes?
a. 67.2 liters of sulfur dioxide gas
b. 0.880 liters of helium gas
c. 1,000 liters of neon gas
**The density of a gas is usually measured in the units of grams per liter. (g/L). We can use the experimentally
determined density of a gas at STP to calculate the molar mass of the gas.
Example: The density of a gaseous compound containing carbon and oxygen is 1.964 g/L at STP. Determine
the molar mass of the compound. What is the formula for this gas?
Sample Calculations:
1. A gaseous compound composed of sulfur and oxygen that is linked to the formation of acid rain has a
density of 3.58 g/L at STP. What is the molar mass of this gas? What is the formula for this gas?
2. What is the density of krypton gas at STP?
3. What is the density of oxygen gas at STP?
C. The Mole Road Map
1. You have now examined a mole in terms of particles, mass, and volume of gases at STP. To convert between
one unit and another, the mole must be used as an intermediate step.
Sample Calculations: Using the mole map, solve the following problems:
Turn your page to 186 and place questions 24-28 here: Show all work.
24a.
b.
c.
d.
25a.
b.
c.
26.
27a.
b.
c.
d.
28.
7.3 Percent Composition and Chemical Formulas
A. Calculating the Percent Composition of a Compound
1. The relative amounts of each element in a compound are expressed as the percent composition, or the
percent by mass of each element in a compound.
Example: An 8.20 gram piece of magnesium combines completely with 5.40 grams of oxygen gas to form a
compound. What is the percent composition of this compound?
Sample Calculations:
1. 9.30 grams of magnesium combine completely with 3.48 grams of nitrogen gas to form a compound.
What is the percent composition of this compound?
2. 29.0 grams of silver combine completely with 4.30 grams of sulfur to form a compound. What is the
percent composition of this compound?
3. What a 14.2 gram sample of mercury (II) oxide is decomposed into its elements by heating, 13.2 grams
of mercury is obtained. What is the percent composition of this compound?
2. To calculate the percent composition of a known compound, use the chemical formula to calculate the molar
mass. Then for each element, calculate the percent by dividing the mass of each element in one mole of the
compound by the molar mass and multiplying by 100%.
Example: Calculate the percent composition of potassium dichromate.
Sample Calculations:
1. Calculate the percent composition of these compounds:
a. sodium bicarbonate
b. ammonium chloride
c. sulfur trioxide
2. Calculate the percent nitrogen in these common fertilizers:
a. CO(NH2)2
b. NH3
c.
NH4NO3
B. Using Percent as a Conversion Factor
1. You can use percent composition to calculate the number of grams of an element contained in a specific
amount of a compound. To do this, you multiply the mass of the compound by a conversion factor that is
based on the percent composition.
Example: Calculate the mass of carbon in 82 grams of propane (C 3H8).
Sample Calculations:
Calculate the grams of nitrogen in 125 grams of each fertilizer.
a. CO(NH2)2
b. NH3
c.
NH4NO3
C. Calculating Empirical Formulas
A. Determining the percent composition of a compound has an important application – calculating the empirical
formula:
B. Define empirical formula:
C. Define molecular formula:
D. The empirical formula may or may not be the same as the molecular formula
E. Practice: These are all molecular formulas, in other words, true formulas. Write their empirical formulas:
H20
H2O2
CO2
N2O4
F. We can use the percent composition of a compound to determine its empirical formula.
Example:
What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen?
Sample Calculations:
1. Calculate the empirical formula of each compound.
a. 94.1% O, 5.9% H
b. 79.8% C, 20.2% H
c. 67.6% Hg, 10.8% S, 21.6% O
d. 27.59% C, 1.15% H, 16.09% N, 55.17% O
D. Calculating Molecular Formulas
1. You can determine the molecular formula of a compound if you know its empirical formula and its molar mass..
Example: Calculate the molecular formula of the compound whose molar mass is 60.0 grams and empirical formula
is CH4N.
Step 1: Find the empirical formula mass
Step 2: Divide the molar mass by the empirical formula mass
Step 3: Distribute this through the subscripts of the empirical formula to arrive at the molecular
Formula
Sample Calculations
1. Find the molecular formula of each compound given its empirical formula and molar mass.
a. ethylene glycol (CH3O), used in antifreeze, molar mass = 62 g/mol
b. p-dichlorobenzene (C3H2Cl), mothballs, molar mass equal 147 g/mole