Accelerated Chemistry
Final Exam Review
Get Ready to Celebrate!!
Unit 1. Safety and Measurement
• Mass is the amount of matter in a substance.
• Weight is the amount of gravity pulling on the
object’s matter.
• Weight can change depending on location, mass
can not.
Unit 1 Safety/Measurement
1. Scientific notation - can only have 1 digit to the left
of the decimal.
2. Significant Figures - know rules. What digit is the
estimated digit? (LAST ONE!)
3. Density, d = m/v
4. Factor Label
5. Basic Safety procedures for lab
6. Rounding Rules
>5, round up
<5, no change
=5 (Number before is even, stays the same)
(Number before is odd, round up).
Significant Figures
Significant Figures Rules
• 1. ALL NONZERO integers are significant.
• 2. Zeros that LEAD a number are NEVER significant.
Ex: 0.00052
• 3. Zeros that are between other nonzero integers ARE
significant. Ex: 105.2
• 4. Zeros at the end of a number are significant ONLY if
• there is a decimal somewhere in the number.
• Ex: not sig = 120,000
–
Sig = 120,000. .23600 1.200 x 105
Metric Conversions 1 mL = 1cm3
Accuracy vs. Precision
•Accuracy- The closeness of a measurement to the
true value.
•Precision- The extent to which an experiment can be
repeated with the same result.
Physical vs. Chemical
Physical Properties that are detected by the senses or can be
identified without a chemical reaction.
Physical Change does not result in a new substance.
Chemical Properties is the ability of a substance to form a
new one.
Chemical Change results in a totally different substance
Unit 2: Solubility
Solids - solubility increases as temperature increases.
Gases - solubility decreases as temperature increases.
Practice solubility graph problems!
•Saturated: A solution in which the solvent has as
much solute dissolved in it as possible for that
temperature.
•Unsaturated: A solution in which more solute can be
added and dissolved in the solvent.
•Supersaturated: A solution in which more solute is
dissolved than would normally be allowed for that
temperature.
Unit 3
Nomenclature/Balancing
Atom is smallest part of element that can exist. Atoms
have a neutral charge.
Ion is an atom or group of atoms that has a + or - charge.
Cations are positively charged ions.
Anions are negatively charged ions.
Monatomic ions have only one type of atom.
Polyatomic ions have more than one type of atom.
Ionic vs. Covalent
Ionic compounds consist of a metal and nonmetal.
Charges must be neutral.
Positive ion first, then negative ion.
If more than one charge, must include Roman
numeral.
Add suffix “ide” to second element.
Reduce to lowest terms. (Ex. NaCl, NaOH)
Covalent compounds consist of two nonmetals.
Not based on charges.
Use numerical prefixes (mono, di, tri, etc.)
Add suffix “ide” to second element.
Cannot reduce. (Ex. CO2, HF)
Write the following formulas for Binary
Ionic Compounds.
Calcium chloride
CaCl2
Beryllium oxide
BeO
Aluminum oxide
Al2O3
Potassium bromide
KBr
Lithium phosphide
Li3P
Examples with varying charges
Lead (II) bromide
PbBr2
Copper (I) nitride
Cu3N
Iron (II) Chloride
FeCl2
Lead (IV) oxide
PbO2
Write the names of the following
FeF2
CoCl2
Cu2S
NiO
Co3N2
CuBr
•Synthesis (Composition or Combination )- takes 2 reactants
to form 1 product A + B ------> AB
•Decomposition- one reactant breaks up into more than one
AB ----->A + B 2KClO3 ----> 2KCl + 3O2
•Single Replacement (Displacement)- Reactant composed of
more than one type of atom gives it’s “Partner” to another reactant.
A + BC-------> AC + B
•Double Replacement(Displacement)-Reactants composed of
more than one type of atom “switch partners”
AB + CD-------> AD + CB
•Combustion-Reaction of compounds containing C’s & H’s
with oxygen to form carbon dioxide and water.
Ex: C6H12O6 +6O2 ---> 6CO2 + 6H2O
Unit 4 Stoichiometry
Mole Road Map
Grams A
Liters A
Moles A
Moles B
Grams B
Liters B
Percent Composition by Mass
1. For each element, multiply atomic mass by subscript:
H2 2 x 1.00 = 2.00
O 1 x 16.00 = 16.00
2. Find formula mass:
2.00 + 16.00 = 18.00 amu
3. Plug in to percent composition for each.
H
2.00 / 18.00 = 11.11 %
O 16.00 / 18.00 = 88.89 %
Empirical Formulas
•Step 1:Convert grams to moles.
•Step 2: Divide each by the smallest number of moles to
get subscripts.
•Step 3: If you do not have a whole number, or one
that is reasonably close, multiply by an integer to obtain a
whole number.
Empirical Formula for a Hydrate
•Step 1: Find the moles of anhydride (compound
without water)
•Step 2: Find the moles of water.
•Step 3: Divide moles of water by moles of anhydride.
This should turn out to be a whole number
•Step 4: Place this whole number into the formula in
front of water.
Empirical Formula
Hg--> 67.6g x 1mole = .337 moles Hg / . 337 = 1
200.59g
S --> 10.8g x 1mole = .337 moles S / . 337 = 1
32.06g
O --> 21.6g x 1mole
16.00g
= 1.35 moles O / . 337 = 4
Mole-Mole Problem
Aluminum and oxygen react to form aluminum oxide. If
you have 2.3 moles of aluminum oxide after the
reaction how many moles of aluminum were used?
4Al + 3O2 ---> 2Al2O3
Mole-Mass Problem
You have 6.7 moles of oxygen gas reacting with unlimited
supply of hydrogen. How many grams of water are
produced?
H2 + O2 ---> 2H2O
Mass-Mass Problem
CaC2 reacts with water to produce C2H2 and
calcium hydroxide. If 51.6g of CaC2 are reacted,
how many grams of calcium hydroxide are
produced?
CaC2 + 2H2O ---> C2H2 + Ca(OH)2
Limiting Reactant
•The limiting reactant is the reactant that
gives you the least amount of product.
How will you know if it is a L.R. Problem?
--- You will be given 2 amounts of a reactant!
Do 2 stoich problems, the one that produces the least
amount is the limiting reactant.
Limiting Reactant
How many grams of CO2 can form when a mixture of 2.93g of
ethylene and 5.29g of oxygen are reacted?
•C2H4(g) + 3 O2(g) ---->2 CO2(g) + 2 H2O(g)
•O2--> 4.85g CO2 possible
•C2H4--> 9.19g CO2 possible
4.85g CO2 can be formed based on LR
which is O2.
Limiting Reactant Problems
How many grams of lithium nitride can be made by
reacting 56.0g of nitrogen with 56.0g of lithium?
N2 + 6Li ----> 2Li3N
Theoretical, Experimental, &
Percent Yields
Theoretical yield is the amount of product that you mathematically
determine (using stoich) from a chemical reaction.
Experimental yield is the amount of product that you actually get
when you perform the reaction in lab.
Percent yield
=
Experimental Yield
Theoretical Yield
x 100
Gases
•1atm = 760 mm Hg
•
= 760 torr
•
= 101,325 Pa
Standard Temperature and Pressure (STP) = 1 atm and 273 K
Mass/Volume Stoich @ STP
Consider the following equation:
C6H12O6(s) + 6O2(g) -----> 6CO2(g) + 6H2O(g)
1. How many liters of carbon dioxide can be produced when
50.0 of of sugar react with excess oxygen at STP?
Consider the following equation:
C6H12O6(s) + 6O2(g) -----> 6CO2(g) + 6H2O(g)
If 15.0L of oxygen react with excess glucose at STP, how many
grams of water will be produced?
Ideal Gas Law
•P = PRESSURE (in atm)
•V= VOLUME (in L)
•n = MOLES OF GAS
•R= GAS CONSTANT, 0.0821 L Atm/mol K
•T= TEMPERATURE (in K)
Gas Stoichiometry
If at STP, use stoichiometry.
If not at STP, use PV=nRT.
1 mole of any gas = 22.4L