Final Exam Review Guide with Practice Questions
Unit 6 Stoichiometry
Standard: Solving for quantities in a chemical equation
 I can describe the units: mole, gram, Liter, formula units
What is the difference between all the units mentioned?
Mole: measure of the amount of a substance
Gram: measure of the mass of a substance
Liter: measure of the volume of a substance
Formula Units: Measure of the number of the smallest indivisible part of a substance i.e atoms or
molecules
 I can convert between units of mole, gram, Liter and formula units
How many liters are present in 13.95mol?
(19.95 mol)(22.4L) = 446.88 L
How many atoms are in 2.3 mol of N2 gas?
(2.3 mol)(6.022X1023 molecules) = 1.38X1024 molecules X 2 atoms per molecule = 2.77 X 1014
How many moles are in 32 grams of CH4?
32 g/ 16 g per mol = 2 mol
 I can explain what a chemical equation means with regard to how many reactants used and
products formed
What do the coefficients actually mean in: 2H2O  O2 + 2H2
Coeffients in an equation represent how many moles and/or molecules are required for each
reactant to produce each product. Very similar to how a recipe works in a cookbook.
 I can determine a mole ratio using a chemical equation
SiO2 + 4HF  SiF4 + 2H2O
How many moles of water are made for every hydrogen fluoride?
1 mol Hf X (2 H2O/4 HF) = 0.5 mole of water
 I can convert between reactants and products using mole ratios
4NH3 + 5O2  4NO + 6H2O
If you have 6 moles of ammonia, how much water can you make in moles?
6 mol NH3 X (6 H2O/4 NH3) = 9 mol of water
 I can convert moles, grams, Liters and formula units of reactants/products into units of
moles, grams, Liters and formula units of another reactant/product using mole ratios
4NH3 + 5O2  4NO + 6H2O. How many atoms of O2 oxygen are produced from required to form 199.4grams
of NO?
(199.4 g NO/ 30 g per mole) X (5 O2/ 4NO) X (6.022X1023) X (2 atoms per molecule) = 1.0X1023
atoms of oxygen
Standard: Solving for Limiting Reagent/Excess Reagent
 I can explain limiting and excess reactants
What is the difference between the two?
A limiting reactant will run out of its designated quantity first. This means the reaction will
be limited by how much of that particular reactant you have.
An excess reactant means you have more than enough of the particular reactant.
 I can determine the limiting and excess reactant using mole ratios
SiO2 + 4HF  SiF4 + 2H2O. If 4.6mol HF is reacted with 3.9mol SiO2 which is the limiting reactant?
4.6 mol HF/4 = 1.15 reactions (Limiting Reactant)
3.9 mol SiO2/ 1 = 3.8 reactions
 I can determine the limiting and excess reactant by converting from grams, Liters or
formula units and then using the mole ratio of the reactants
SiO2 + 4HF  SiF4 + 2H2O. If 50 grams HF is reacted with 30 grams SiO2 which is the limiting reactant?
(50 grams HF/20 g per mol)/ 4 = 0.625 reactions
(30 grams SiO2/ 1 g per mol)/ 1 = 0.499 reactions (Limiting Reactant)
 I can determine how much product is made by using the limiting reactant
SiO2 + 4HF  SiF4 + 2H2O. If 120 grams HF is reacted with 3.9mol SiO2, how much water is produced?
(120 grams HF/20 g per mol)/ 4 = 1.5 reactions (limiting reactant)
(3.9 mol SiO2/ 1 mol) = 3.9 reactions
6 mol of HF from above X (2 H2O/4 HF) = 3 mol of water produced
Standard: Calculating Percent Yield
 I can determine how much product should be made using stoichiometry
*See above practice problems
 I can identify the difference between how much product should be produced compared to
what was actually produced in an experiment
A scientist calculated he will collect 123 grams of water in a decomposition reaction. When he massed the
amount of water from his experiment, he collected 120 grams. Which is the theoretical and which is the
experimental?
123 grams = theoretical
120 grams = experimental
 I can determine the percent of product collected from an experiment based on how much
should have formed
You calculate a reaction to produce 15.62grams but it actually produced 11.08grams. What is the %
yield?
(11.08g/15.62g) X 100 = 70.93%
Unit 7: Gas Laws
Standard 1: Combined Gas Laws
 I can explain all three gas laws and how the variables affect one another
Describe each variable in the combined gas law equation.
Pressure: The amount of collisions the gas is exerting on itself and the container
Volume: The measure of space the gas occupies
Temperature: Always in Kelvin is the average kinetic energy of all the atoms of gas.
Explain how the other variables are affected if T increases. What is V decreases? What is P increases?
If temperature increases, particles are moving faster and colliding more often which increases
pressure and would therefore increase the volume of the container. *Be careful with this though
because we usually talk about only two variables at a time due to the fact that the three will interact
in different ways. For example, when pressure increases the container would increase in volume,
but once the container has more space you have less collisions and the pressure returns to the
original value. So, try explaining only two variables at a time and assume the other one is constant
like in a rigid container that cannot change volume.
If volume were to decrease the pressure would increase because the gas would collide more in tight
space. If pressure were constant, then the temperature were increase instead.
If pressure increased on the container the volume would decrease. If volume were constant then
the temperature would increase.
 I can solve for unknown variables in the combined gas law equation
Calculate the pressure of a gas with volume= 126mL, n=4.53, and T= 35C.
PV=nRT
(P)(.126 L)=(4.53 mol)(0.0821 atm*L/mol*K)(308K)
P = 909.1 atm
 I can convert units of pressure, temperature and volume to fit into a gas law problem correctly
How many kPa are in 1.56 atm?
1.56 atm X (101.3 kPa/1atm) = 158.028 kPa
Standard 2: Ideal Gas Law
 I can explain what the ideal gas law is and how each variable will impact the others
What happens to pressure and temperature when the numbers of moles increases in a container?
When the amount of gas is increase there are more particles colliding and more energy so the
temperature and pressure would increase.
 I can explain the assumptions behind what an ideal gas even is and how it acts
At the atomic level, describe an atom of gas and how it collides and the volume according to ideal gas law.
We assume each formula unit occupies zero space and that all particles have zero attractive forces
to eachother allowing for the PV=nRT math model to work correctly. However, real gases would
have to account for some size of particles and slight attraction. (There is an equation for that )
 I can solve for unknown variables in the ideal gas law equation
Calculate the number of grams of H2O(g) with a vol= 250mL, P=2.06atm. and Temp= 416K.
(2.05atm)(.25L)=n(0.0821 atm*L/mol*K)(416K)
n= 0.015 mol of water X (18 g per mol) = 0.27 grams of H2O
 I can apply stoichiometry to help solve and explain ideal gas law problems
A teacher wants to produce hydrogen gas by adding magnesium metal to hydrochloric acid,
according to the balanced equation below.
Mg + 2HCl  MgCl2 + H2
If the teacher wants to produce 11.78 Liters of gas at room conditions of 25 °C and 1.35 atm, what
mass of magnesium is required?
(1.35atm)(11.78 L)=n(0.0821 atm*L/mol*K)(298K)
n = 0.65 mol of H2
0.65 mol H2 X (1Mg/1H2) = 0.65 mol of Mg X (24.3 g per mol) = 15.8 g Mg
 I can apply the ideal gas law to determine densities of gases
What is the density of Cl2 gas at STP?
1 mol of Cl2 gas = 70.9 g
70.9g/ 22.4 L = 3.17 g/L
Standard 3: Graham’s Law
 I can explain Graham’s law
What is the rate of diffusion/effusion of gases based on?
How fast a gas travels is based on how much mass each particle has. The smaller the mass the faster
is can diffuse or travel through holes.
 I can identify the rate of effusion of various gases compared to one another
Which gas effuses faster: Methane or Oxygen?
Methane has a molar mass of 16 g/mol while oxygen gas (O2) has a molar mass of 32 g/mol, so
methane would effuse faster.
 I can determine the rate of effusion of gases using molar mass and the equation
What would be the rate of effusion of Hydrogen Sulfide (33 g/mol) compared to Xenon?
RHS/RXe = √(131.3 g/mol)/(33 g/mol)
1.99 = how fast HS is compared to xenon (This means HS effuses twice as fast)
Unit 8: IMFs
Standard: Phase Changes and Heating Curves
 I can describe all types of phase changes
List every phase change and what phase of matter is the starting point and which is the end point
Deposition – From a gas to a solid, exothermic/releases energy…Why? Going from a high energy state to
low energy state would require a loss of energy
Sublimation – From solid to gas, endothermic/requires energy…Why?
Melting – From solid to liquid, endothermic/requires energy…Why?
Freezing- from liquid to solid, exothermic/releases energy---Why?
Boiling – From liquid to gas, endothermic/ requires energy…Why?
Condensation- gas to liquid, exothermic/releases energy…Why?

I can interpret a phase change diagram and create one from data
Label the phases, normal boiling and freezing points, and triple point on the diagram. Show an arrow for
each phase change and label it.
Liquid
Solid
Gas
Triple Point

I can interpret a heating curve and create one from data
Label all the phase changes and phases on the heating curve. (Be able to read the graph too.)
90
E
80
T
E
M
P
(oC)
70
D
60
50
40
C
30
20
10
B
0
-10
A
10
20 30 40 50 60 70 80 90 100 110 120
TIME (min)
A = Solid
B = Melting
C = Liquid
D = Boiling
E= Gas

I can explain why phase changes happen in terms of IMFs and energy
Why does temperature remain constant during a phase change?
All the energy, which temperature measures, is either being absorbed to break IMFs or bond, or the
energy is being released to the environment when forming the bonds, and the overall substance is not
gaining or losing energy.
Why do liquids boil? (Involve both variables in your explanation)
A liquid will boiling when the kinetic energy of the molecules is enough to escape the attractive forces
holding them together. When the particles escape they exert a vapor pressure on the air around them,
and if the vapor pressure exceeds the atmospheric pressure surrounding them, then the liquid will boil.
Essentially, once the pressure exerted by the liquid exceeds the pressure from the air around it, then the
liquid boils. You can either lower the atmospheric pressure (Vacuum) or increase the temperature of the
liquid so more particles escape and the vapor pressure increases enough.
Standard: Properties and Phases of Matter

I can describe the properties of a solid, liquid and gas
1. Describe the properties of the following phases

gas
liquid
solid
assumes the shape and
volume of its container
particles can move past
one another
assumes the shape of the
part of the container which
it occupies
particles can move/slide
past one another
retains a fixed volume and
shape
rigid - particles locked into
place
compressible
lots of free space between
particles
not easily compressible
little free space between
particles
not easily compressible
little free space between
particles
flows easily
particles can move past
one another
flows easily
particles can move/slide
past one another
does not flow easily
rigid - particles cannot
move/slide past one
another
I can explain the kinetic molecular theory and how it relates to each phase
Use the kinetic molecular theory to describe the particles energy and movement in each phase of
matter.
a. Solid: Low energy, strong attractive forces, fixed positions, vibrate, little space inbetween.
b. Liquid: Medium energy, attractive forces (IMFs), particles can move freely, but are held
down by attractive forces, some space in-between, but not much
c. Gas: High energy, few if any attractive forces, particles are free to move randomly, lots of
space in between particles

I can diagram what each phase would look like at the atomic level
Draw a diagram of the particles in each phase of matter to represent your descriptions above.
Standard: Intermolecular Forces
 I can describe each type of IMF
List the IMFs and explain the requirement for a molecule to have each type and what it is.
Van der Waals/London Dispersion – Nonpolar particles have a dipole moment when electrons are unevenly
distributed causing a slightly positive and negative end (poles) causing particles to have a weak attraction
with one another
Dipole-Dipole : Polar molecules have a permanent uneven distribution of electrons causing the molecule to
have a positive and negative end like a magnet. This allows the molecules to have an attraction with each
other that is stronger than London dispersion
Hydrogen bonds : Polar molecules that have H bonded with O,N or F will form hydrogen bonds which are
basically the strongest dipole-dipole attractions

I can explain the strength of each IMF
List the IMFs in order of increasing strength AND explain why each IMF exhibits that strength.
London Dispersion only slight + and – ends, Dipole-Dipole larger + and – ends for stronger attraction,
Hydrogen bonds with the largest + and – ends due to O, N and F having large electronegativities

I can determine what IMF an atom or molecule would exhibit based on its structure or formula
What is the IMF of each of the following:
H2O
H-Bond
CH4
CH2O2
C3H8
CO2
HBr
HF
Dispersion
H-Bond
Dispersion
Dispersion
Dipole
H-bonds
Standard: Specific Heat
 I can explain what specific heat capacity is
What does heat capacity represent?
The amount of energy required to change the temperature of a substance OR how easily substances can
transfer energy.
 I can identify which materials have higher or lower specific heats
Why does water take so long to warm up, but sandy beaches are always hot during the day in summer?
Water does not transfer/absorb energy easily so its specific heat is high meaning it requires a lot of energy to
heat up. Sand is the opposite
 I can calculate all variables involved with specific heat
A 25.75-g piece of iron absorbs 1200.75 joules of heat energy, and its temperature changes from 25ºC to
110ºC. Calculate the heat capacity of iron.
Q=mcΔT
1200.75 J = 25.75 g * C * (110-25)
C = 0.5486 J/g*Celsius
How many joules of heat are needed to raise the temperature of 15.0 g of aluminum from 22ºC to 44ºC, if the
specific heat of aluminum is 0.95 J/gºC?
Q=15g*0.95*(44-22)
Q = 313.5 J
To what temperature will a 150.0 g piece of glass raise if it absorbs 10,275 joules of heat and its heat capacity
is 0.50 J/gºC? The initial temperature of the glass is 20.0ºC.
10,275 J = 150 g * 0.5 * (Tf – 20)
137 = Tf – 20
157 degrees C= Tf
Unit 9: Solutions and Colligative Properties
Standard: Solubility and Properties of Solutions
 I can describe the parts to a solution
Describe solute and solvent.
Solute gets mixed in with a solution and is dissolved Ex: Sugar, salt, ingredients
Solvent is the base of the solution and makes up the majority. Also does the dissolving. Ex: Water
 I can explain “like dissolves like”
What does this phrase mean?
Polar molecules dissolve Polar molecules
Nonpolar molecules dissolve nonpolar molecules
 I can predict solubility based on solute and solvent polarity
What types of molecules dissolve in water? Will solid carbon dioxide dissolve in oil?
Other polar molecules. Yes
 I can identify electrolytes
NaCl or Sugar?
NaCl
 I can explain properties of electrolytes in solution
What happens to electrolytes in a solution? What can they do?
They dissociate into the ions that make them up. These free moving ions can conduct electricity
since that is what electricity is….
 I can determine if a solution is saturated, unsaturated or supersaturated
How do you know when a solution is one of these terms?
Unsaturated solutions can dissolve more solute
Saturated solutions have the maximum solute dissolved in the solution
Supersaturated solutions have more than the maximum solute dissolved in solution
Standard: Net Ionic Equations
 I can identify whether a molecule/compound will become ions in a solution or not
Can you use the solubility rules correctly?
See reference
 I can explain what happens when two reactants dissolve into a solution
Draw NaCl and Ca(NO3)2 as they would exist in a solution.
 I can write a complete ionic equation
What is the ionic equation for: K2CO3 + Ca(NO3)2 
K2CO3 + Ca(NO3)2  2KNO3 + CaCO3
2K+1 + CO32- + Ca2+ + 2 NO31-  2K+1 + 2 NO31- + CaCO3(s)
 I can write a complete net ionic equation
What is the net ionic for the problem above?
CO32- + Ca2+  CaCO3(s)
Standard: Molarity
 I can explain concentration
What does concentration mean?
Concentration is a measure of how many solute particles there are in a given volume. In short, how
much stuff is packed into a specific area.
 I can interpret pictures and diagrams based on their concentration
What would a diagram of a highly concentrated solution look like at the atomic level? What about a dilute solution?
 I can determine the concentration of a solution using an equation
Calculate the molarity of a solution made with 3.44moles in 16L of water.
Standard: Colligative Properties
 I can describe what colligative properties are
What are they?
The properties of solutions that change based on how much solute is in the solution. Vapor pressure,
freezing point and boiling point
 I can explain how colligative properties affect a solution
What happens to a solution when you add solute? (3 properties change…)
When solute is added:
1. The solvent particles have trouble forming IMFs and bonds to become solid. So, the freezing point
is depressed
2. The solvent particles have a harder time escaping their IMFs to become gas, so the vapor pressure
is decreased
3. Because the solvent particles are vaporizing less, the boiling point is elevated
 I can determine how much a solution will be affected by colligative properties based on number of
ions added to the solvent
Which molecule would change the colligative properties the most: Salt, Sugar or aluminum nitrate? How do you
know?
Salt = 2 ions
Sugar = 1 particle
Aluminum nitrate = 4 ions which would have the largest effect on colligative particles
Unit 10: Acids and Bases
Standard: Properties
 I can define acids and bases
What is an Arrhenius acid and base? What about Bronsted-Lowry?
Arrhenius acids and bases add H+ and OH- to solution
Bronsted-Lowry acids and bases donate and accept protons
 I can name acids and bases
What is the formula of lithium hydroxide and hydroiodic acid?
LiOH and HI
Name the following: HCl
Sr(OH)2
H3PO4
Hydrochloric acid
strontium hydroxide
Phosphoric acid
 I can identify different types of acids and bases according to their formula
Label the compounds above as Arrhenius and Bronsted-Lowry acid and bases.
They are all both
 I can describe the properties of acids and bases
What are the properties of an acid and base?
Acid:
•
•
•
•
•
Tastes sour
Reacts with bases to produce water and salt
Reacts with metals to produce hydrogen gas
Changes colors with indicators
Conducts electricity
•
•
•
•
•
Feel slippery
Taste bitter
Change colors with indicators
Conduct electricity
React with acids to form water and salt
Base:
 I can identify products we use as acids and bases
What types of everyday items tend to be acidic? Basic? Can you give specific examples?
Food items are acidic and cleaning items are basic.
Lemons and citric juices vs. Bleach and detergents
Standard: Reactions
 I can write equations showing how certain compounds form acidic and basic compounds
What is the equation for water to react with itself and form acidic and basic ions?
H2O + H2O  H3O+ + OHWhat is the equation for potassium hydroxide mixing with water? What about hydrobromic acid with water?
KOH + H2O  K+ + OH- + H2O (Remember bases are electrolytes and become ions)
HBr + H2O  Br- + H3O+
(Same as the above, but the proton will exist as hydronium, H3O+
in solution)
 I can predict the products of an acid base reaction or a neutralization reaction
Finish the equations:
LiOH + HI  H2O + LiI
Base Acid C.A C.B.
NaOH + HNO3  H2O + NaNO3
Base Acid
C.A.
C.B
NH3 + H2O  NH4+ + OH-
Base Acid C.A.
C.B.
 I can explain societal effects of acids and bases like acid rain and limestone lakes
How does acid rain form? What are the effects on the environment? Why do limestone lakes not become acidic due
to acid rain?
Acid rain is rainwater with a lower pH. It forms when gaseous pollutants like carbon diocide and
nitrous oxide react with water vapor in the air to form acidic compounds like nitric acid or carbonic
acid. This rainwater can corrode metal and lower the pH of lakes and rivers which arms
ecosystems.
 I can identify different types of acids and bases in a reaction (Bronsted, Lewis, Arrhenius)
Label the three reactions you predicted the products for above.
Standard: pH
 I can identify various testing methods to classify acids and bases
List all the indicators we used in class to identify acids and bases.
pH paper, red cabbage, Phenolphthalein
 I can describe how various testing methods work to classify acids and bases work
How exactly does an indicator work?
Indicators react to form new compounds which absorb and reflect different wavelengths of light,
thus giving us the different colors we see when acids and bases are present
 I can classify acids and bases according to pH value
Draw the pH scale and label the areas that are acidic, basic and neutral.
Alkaline means basic
 I can rank acids and bases in order of strength bases on their concentrations and pH
values
Rank the following solutions from most acidic to least: pH=2
Most 2  5  8  12 Least
Which solution is the most acidic:
[H+] = 2 X 10-11
pH=12 pH=5
pH= 8
[OH-] = 2.4 X10-3
HCl with 0.035 M
 I can identify equations used for pH of substances
HCl = .035 M
What are the equation(s) used for the ionization of water?
Kw = [OH-][H+]
What are the equation(s) used to determine pH?
pH = -log[H+]
 I can calculate pH based on solution concentrations or concentrations based on pH
Solve the following problems:
pH of a 1.2 M acid?
-log [ 1.2] = -0.792
[H+] of a solution that has a pH of 4.76?
10-4.76 = 0.0000174 M
pH of a solution that contains .0046 NaOH?
1X10-14 = [.0046][H+]
[H+] = 2.174 X 10-12 M
-log[2.174 X 10-12 M ] = pH = 11.66
 I can use stoichiometry to determine mass of acid and base needed or used to make
specific concentrations
What mass of HNO3 would be required to make a solution that is 2M and 1.3L?
2M = mol/1.3L
Mol of HNO3 = 2.6 X (63 g per mol HNO3) = 163.8 grams
If you massed out 12.67 grams of HI and put it into 200mL of water, what pH would the solution
be?
12.67 g HI/(128 g per mol of HI) = 0.099 mol of HI/ 0.2L = 0.495 M
-log [ 0.495] = 0.3054 = pH
What mass of HCl would be required to make a .45L solution have a pH of 3?
10-pH = 10-3 = [H+] = 0.001 M = mol of HCl/ 0.45 L
Mol of HCl = 0.00045 X (36.45 g per mol of HCl) = 0.0164 grams of HCl
Standard: Titrations
 I can describe an acid-base titration and the method used
What is a titration?
The technique used to determine the unknown concentration of an acid or base using a standard
known concentration of the opposite type of solution
 I can explain why a titration works to determine a concentration
What does the neutralization point mean?
Equal number of particles of base and acid are present
Describe what is happening in the solution when the indicator changes colors during a titration.
See above statement on how indicators work
 I can determine concentrations and pH of solutions based on titration data
Here are some buret readings from titrating 25mL of an unknown acid with a 0.3 M standard solution of NaOH
0.3 M NaOH
initial reading:
48.7 mL
final reading:
32.4 mL
What is the molarity of the acid? What is the pH of the acid? What is the pH of the base?
MaVa=MbVb
Ma(25mL) = (0.3 M)(48.7mL)
Ma = 0.5844
-log [0.5844] = pH of acid = 0.233
pH of base (which is an entirely separate solution):
Kw=[0.3][H+]
[H+] = 3.33X10-14
-log[3.33X10-14] = 13.48 = pH of base
Unit 11: Organic Chemistry
Standard: Naming and Drawing
 I can name and draw simple alkane hydrocarbons
What are the ten alkanes? What do they look like?
Methane 1C
Ethane 2C
Propane 3C
Butane 4C
Pentane 5C
Hexane 6C
Heptane 7C
Octane 8C
Nonane 9C
Decane 10C

I can name and draw hydrocarbons with branches
Draw 2,3-methylhexane 2,3-dimethylhexane and 4-ethyl-5-propyloctane
Name:
3,7,-diethyl-6-propyl-decane (the 3 carbon is actually a different name, but we will call it propyl because it
is 3 carbons)

I can name and draw hydrocarbons with double and triple bonds
Draw ethane and 2-methyl-5-hexyne (The triple bond is on the 3 carbon, obviously it should be on the fifth)

I can name and draw hydrocarbons with functional groups
1,5-dipentanol
 I can name and draw hydrocarbons with halogens
Name both:
2-bromo-propane
1-bromo-propane
Standard: Concepts and Patterns of Organic Compounds
 I can explain (in general) why certain compounds exhibit different characteristics like smell
What role do functional groups serve on organic compounds? Why is a benzene ring referred to as an aromatic
ring? (relate your explanation to its structure)
Functional groups are called such because each group has a particular function due to their structure on
an organic molecule.
Benzene is called an aromatic rings because of its characteristic smell when present in molecules.

I can explain hydrocarbons
What is a hydrocarbon?
A carbon chain with hydrogens surrounding them. Very stable molecules that have a lot of potential
energy stored between all their bonds

I can explain the difference between saturated and unsaturated fats/hydrocarbons
What is the difference? Draw two saturated fats and two unsaturated fats.
Saturated fats are hydrocarbons that have the maximum number of hydrogens bonded onto the
carbons. This is a very stable molecule that requires immense amount of energy to break apart.
Unsaturated fats have less than the maximum number of carbons and are slightly easier to break down
because they are less stable and not as uniform in structure.