AP Chemistry: Chapter 1-3 Notes Outline
Objectives:
Memorize key ions and atoms symbols and their charges
Matter
Be able to count and use significant figures in calculations
Be able to convert measurements using dimensional analysis
Be able to name ionic compounds
Be able to name acids
Balance chemical equations
Use dimensional analysis to solve stoichiometric problems
Use dimensional analysis to do limiting reactant problems
Use dimensional analysis to calculate percent yield
Calculate percent composition
Calculate empirical and molecular formulas
Memorize Key ions and atoms:
You must score a 95/100 to receive credit. See memory sheet for these items.
Matter:
Pure Substances vs Mixtures
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Mixtures
Homogeneous
Definition
Examples
Heterogeneous
Definition
Examples
Separation of Mixtures
Based on Physical Properties
Distillation
Chromatography
Solubility
Others
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Significant Figures:
Why significant figures?
Rules for counting significant figures
1. All nonzero numbers count
2. Leading zeros don’t count
3. Trailing zeros count if there is a decimal
4. Trailing zeros don’t count if there is no decimal
Examples of counting significant figures:
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Rules for calculating with significant figures:
1. Addition and Subtraction: You are only as good as your least accurate place value
2. Multiplication and Division: You are only as good as your least accurate number of significant figures
Examples of calculations involving significant figures
Convert with Dimensional Analysis
The principle is that you are multiplying several fractions tog tether, but each is the same measurement so that you are essentially multiplying
by 1.
Dimensional Analysis Rules:
1. All numbers must be written as fractions—this includes numbers that are given—so 6.7 grams is
2. All subsequent multipliers are written such that the units must cancel: For example:
6.7 g
1
6.7 g
1lb
x
 0.015lbs . Notice how the grams
1
454 g
cancel
3. All numbers on the top of a series of fractions are multiplied and all numbers on the bottom of a series of fractions are divided. For
example:
143.55mL
1L
0.2642gal 4qt 2 pt
x
x
x
x
 0.30341pt . In your calculator you would: 143.55*0.264*4*2/1000=.
1
1000mL
1L
1gal 1qt
Type 1: One Dimensional Problems
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Type 2 : Working with multiple dimensions—Convert 55 miles per hour into meters per second.
Type 3: Area and Volume: Hint a m3 = m m m
Type 4: Complex Problems
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Naming Ionic Compounds
Rules
1. Name Cation
2. Use Roman Numerals if the metal is a transition metal (Exceptions: Zn, Ag, Al)
3. Name the Anion (It may be a polyion)
Examples: From Name to Formula
Examples: From Formula to Name
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Naming Acids
Rules: Always check the anion of the acid
1. If the anion ends in ide change to hydo________ic acid
2. If the anion ends in ate change to ________ic acid
3. If the anion ends in ite change to ________ous acid
4. If the anion is per____ate change to per_____ic acid
Examples
Balance chemical equations
Balance the following:
__O2 
__C6H12O6
+
__CO2
__H2
+
__O2 
__H2O
__H2
+
__N2 
__NH3
__NO +
__O2 
__NO2
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+
__H2O
Stoichiometry Problems using dimensional Analysis
Rules
1. Always start with a balanced equation
2. Write what you know underneath the chemical that is known (5.23grams)
3. Place a question mark (?) underneath what you are looking for.
4. Use dimensional analysis to convert to the chemical you are trying to find information about. Note that you will most likely need to
use the mole-mole ratio (also called the stoichiometric ratio) to convert from one chemical to another.
Examples
How many liters of hydrogen gas is formed when 13.5-grams of calcium reacts with sulfuric acid?
Nitrogen gas reacts with oxygen to form dinitrogen trioxide. How many molecules of oxygen are needed to make 5.5 L of N2O3 at STP?
Calcium Carbonate decomposes into calcium oxide and a common gas. When 45.5 grams of calcium oxide is formed how many liters of
gas is also formed from this reaction.
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Limiting Reactant Problems
Rules
1. Convert all reactants to moles
2. Divide by the coefficient in the balanced equation
3. The chemical with the lowest number of moles is the limiting reactant (LR)
4. Use the LR for all further calculations
Example:
When 114.0 g of iron and 252.7 g of chlorine gas reacts, iron(III) chloride is formed.
a. Write a balanced equation
b. What is the limiting reactant?
c. How many grams of Iron (III) chloride is formed
d. How much excess reagent is left over at the end of the experiment?
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33.6 grams of sulfur dioxide reacts with 55.3 grams of water to form sulfurous acid.
a. Write a balanced equation.
b. What is the limiting reactant?
c. How many grams of sulfurous acid will be formed?
d. How much excess reagent will remain?
Percent Yield
Rules
1. Balanced Equation
2. Write amount given under the reactant that has a value
3. You will be given an amount of product (write this under the product) this is the actual amount produced
4. Convert given reactant amount to same product (and units) that was given
Actual Pr oduced
5. PercentYield 
x100
TheoreticallyMade
Examples:
Nitrogen gas reacts with hydrogen gas to make ammonia (NH3). 15.5-L of N2 reacts at STP to make 30-L of ammonia. What is the
percentage yield?
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Percent Composition
Find the percentage composition of a chemical by finding the total molar mass of the compound find the mass percent of each element in a
compound:
Example: NH4NO3
Percent Error
Observed  Actual
% Error 
x100
Actual
Empirical Formula:
Rules
1. Percent to Mass
2. Mass to Moles
3. Divide by small
4. Times till whole
Example:
Determine the empirical formula of a compound with 52.8% Sn, 12.4% Fe, 16% C and 18.8% N.
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Molecular Formulas
1. Molecular Formulas are ________ _________not just the simplest _____________
2. the molar mass of the empirical formula (EF) is always a multiple of the MW.
3. Use the MW to determine by what factor you need to multiply the empirical formula to find the molecular formula.
Example
 Determine the molecular formula for a compound that contains 22.5% Na, 30.4% P and 47.1% O and a molar mass of 306 g/mol
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Example: Complex Molecular and Empirical Formula Problems—Combustion Problems
 Many homes in rural America are heated by propane gas, a compound that contains only carbon and hydrogen. Complete combustion
of a sample of propane produced 2.641 grams of carbon dioxide and 1.442 grams of water as the only products. Find the empirical
formula of propane.
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 Galactose (Gal) (also called brain sugar) is a type of sugar found in dairy products, in sugar beets and other gums and mucilages.
When 2.315 grams of galactos is completely burned it produces 3.953 grams of carbon dioxide and 1.389 grams of water. The molar
mass of galactose is between 172 g/mol and 186 g/mol. What is the empirical and molecular formula of galactose?
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: Three volatile compounds X, Y, and Z each contain element Q. The percent by weight of element Q in each compound was
determined. Some of the data obtained are given below.
Compound
Percent by weight Molecular
of Element Q
Weight
X
64.8%
?
Y
73.0%
104.
Z
59.3%
64.0
(a) The vapor density of compound X at 27˚C and 750. mm Hg was determined to be 3.53 grams per liter. Calculate the
molecular weight of compound X.
(b) Determine the mass of element Q contained in 1.00 mole of each of the three compounds.
(c) Calculate the most probable value of the atomic weight of element Q.
(d) Compound Z contains carbon, hydrogen, and element Q. When 1.00 gram of compound Z is oxidized and all of the carbon
and hydrogen are converted to oxides, 1.37 grams of CO2 and 0.281 gram of water are produced. Determine the most
probable molecular formula of compound Z.
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