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Name ____________KEY______________________ Period ______________
Chemistry Final Exam STUDY/ Review Guide
MATTER:
Fill in the following chart for TYPICAL Solids, Liquids and Gases (not water)
State of
Particle
Shape
Density Rank
Energy Rank
Matter
Arrangement
Definite or
1 = highest
1 = highest
Regular or
Indefinite
3 = lowest
3 = lowest
Irregular
Compressibility
Yes or No
Gas
irreg
Indef
3
1
Y
Liquid
irreg
Def
2
2
N
Solid
reg
Def
1
3
N
Atoms – draw a diagram of an atom of Lithium (mass number 7), labeling the following parts – nucleus,
electron cloud, protons, neutrons, electrons
Nucleus – center of atom
Electron cloud – outside of nucleus
Protons – in nucleus with (+) symbol
Neutrons – in nucleus with no symbol
Electrons – in energy levels outside of nucleus
Label as: molecule, atom, compound, substance, homogeneous mixture, heterogeneous mixture
 Substance/Molecule - H2O
 Substance / atom - A single particle of pure potassium
 Heterogeneous mixture - A sample of sand scooped from a beach
 Substance / ionic compound - Sodium chloride
 Homogeneous mixture - Iced tea
 Atom - The smallest particle of an element that retains all of its chemical properties
 Heterogeneous mixture - Chicken soup
PHYSICAL AND CHEMICAL CHANGES AND PROPERTIES
Classify the following as a chemical or physical property or chemical or physical change
________ 1. PC - Liquid collects on the side of a cold drink glass
________ 2. CP - Hexane is flammable
________ 3. CC - Baking soda decomposes in the cake, releasing CO2 to make it fluffy.
________ 4. PP - Copper II sulfate is blue.
________ 5. PC - Ice is melting
________ 6. PP - Pure water has a density of 1 g/cm3
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STATES OF MATTER:
Mark the statements as true or false. If false, make changes to correct the statement.
1.
Liquids have the greatest amount of kinetic energy among the phases of matter – FALSE, gases do.
2. In general, a solid is the densest phase of matter. - TRUE
3.
Particles in a solid demonstrate rotation, vibration and translation. FALSE – no translation
4. The particles of liquids have low amounts of translation unless they are stirred. - TRUE
5. When a solid is melting, its temperature is still rising as it melts. – FALSE – holds steady
6. Both solids and liquids can flow. – FALSE – only liquids can flow
7. Translation is the type of kinetic energy in which a particle moves from place to place. - TRUE
DENSITY CALCULATIONS
Solve the following problems involving density calculations. Show all work.
What is the formula for density? D= m/v
A cube is 17 cm3 in volume with mass of 4.0 grams. Calculate its density. 0.24g/cm3
If a ball has a density of 27.5 g/cm3, determine its volume if its mass is 626 g. 22.8g
METRIC UNITS
Perform the following determinations.
________ 1.
1 kilometer (km) = _____1000____ meters (m)
________ 2.
1 milligram (mg)= _____.001___ grams (g)
________ 3.
1 liter (L)
________ 4.
2.3 kilograms (kg)= ____2300_______ grams (g)
________ 5.
12 meters (m) = ___________1200_______ centimeters (cm)
= __1000____ milliliters (mL)
ATOMIC STRUCTURE:
Fill in the following table with the appropriate information
Subatomic
Charge
Location
Relative Mass
Particle
Nucleus or
electron cloud
Proton
+1
Nucleus
1 amu
Neutron
O
Nucleus
1 amu
Electron
-1
Electron cloud
1/1640th amu
Order of Size
1= largest
3 = smallest
1
1
(same!)
3
2
PERIODIC TABLE:
Answer the following questions
What number of valence electrons indicates the most stable arrangement? Which group on the periodic table fits
this description without forming ions?
8 valence e-; noble gases
Elements on the periodic table are organized in order of increasing ____atomic number________.?
Elements are identified by the number of ___protons_____ in their nucleus. The number of ___electrons_______
and ____neutrons_______ may vary.
ORGANIZATION OF THE PERIODIC TABLE
Label the following on the diagram below. Then, label the number of valence electrons inside the space for
each group.
s-block
metals
Aluminum group
Halogens
p-block
non-metals
Carbon group
Noble Gases
d-block
alkali metals
Nitrogen group
transition metals
f-block
alkali earth metals
Oxygen Group
inner transition metals
Non-metals 
Metalloids
Border
stairstep
s - block
Metals
p - block
d – block
Transition Metals
f – block
Inner Transition Metals
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ELECTRON CONFIGURATION:
Write the full electron configuration for the following elements. Circle the valence electrons.
Strontium 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Helium 1s2
Oxygen
Barium 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
1s2 2s2 2p4
5p6 6s2
Neon 1s2 2s2 2p6
Fluorine
1s2
2s2
2p5
BONDING
Draw the covalent structure for the following chemical formulas. Name the shape of the molecule.
a) PBr3
c) SO3
b) SO2
d) PO43-
Indicate the correct word(s) to complete each sentence.
________ 1. Ionic bonds are formed between metals and non-metals
________ 2. Covalent bonds are formed between non-metals and other non-metals.
________ 3. Metals do not form bonds with other metals
________ 4. The transition metals lose electrons to form ions.
________ 5. When comparing degree of polarity between two bonds, the bond with
the greatest polarity has the largest difference between the two atoms
in the bond.
ELECTRON DOT STRUCTURES
Draw the electron dot structure of the valence electrons for the following elements. Indicate which, if any, fulfill the
“octet rule” when neutral. If they do not fulfill the octet rule when neutral, indicate the number of electrons gained
or lost for the atom to fulfill the rule.
Sodium
Silicon
Magnesium Mg
Phosphorus
Aluminum
Al
Sulfur
Si
P
Chlorine
Cl
Argon
Ar
S
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ION FORMATION
Fill in the following table regarding ions
Element
Group Name
Ion Type
cation/anion
Sodium
Alkali metals
cation
Charge
# and +/+1
Metal /
Non-Metal
Metal
Selenium
Oxygen group
anion
-2
Non-metal
Boron
B/Al group
cation
+3
Metalloid
Phosphorus
Nitrogen group
anion
-3
Non-metal
Calcium
Alkaline Earth
cation
+2
Metal
Fluorine
Halogens
anion
-1
Non-metal
Radon
Noble Gases
none
none
Non-metal
Phosphorus
Nitrogen group
anion
-3
Non-metal
BONDING
Fill in the following table regarding ionic bonding
Cation
Charge
Anion
Charge
Calcium
+2
Sulfur
-2
Balanced
Formula
CaS
Strontium
+2
Chlorine
-1
SrCl2
Aluminum
+3
Iodine
-1
AlI3
Cesium
+1
Nitrogen
-3
Cs3N
Magnesium
+2
Oxygen
-2
MgO
Boron
+3
Phosphorus
-3
BP
Sodium
+1
Fluorine
-1
NaF
Chemical
Name
Calcium
sulfide
Strontium
chloride
Aluminum
iodide
Cesium
nitride
Magnesium
oxide
Boron
phosphide
Sodium
fluoride
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POLYATOMIC IONS
Fill in the following chart
Cation
Formula and Anion
Charge
Calcium
Ca +2
Hydroxide
Formula and Balanced
Charge
Formula
-1
OH
Ca(OH)2
Strontium
Sr +2
Nitrate
NO3 -1
Sr(NO3)2
Ammonium
NH4 +1
Acetate
C2H3O2 -1
NH4C2H3O2
Cesium
Cs +1
Chlorate
ClO3 -1
CsClO3
Magnesium
Mg +2
Cyanide
CN -1
Mg(CN)2
Boron
B +3
Sulfate
SO4 -2
B2(SO4)3
Sodium
Na +1
Nitrite
NO2 -1
NaNO2
Chemical
Name
Calcium
hydroxide
Strontium
nitrate
Ammonium
acetate
Cesium
chlorate
Magnesium
cyanide
Boron
sulfate
Sodium
nitrite
BALANCING EQUATIONS AND REACTION SYMBOLS
Balance all equations. Label reaction type for all reactions.
1.
N2(g) +
3H2(g)
2.
2KClO3(s)
∆
3.
2NaCl(s)
+
4.
2AgNO3(aq) + MgCl2(aq)  2AgCl↓ + Mg(NO3)2(aq) DOUBLE REP.
5.
C3H8(l)

+

2NH3(g)
SYNTHESIS
3O2(g)
DECOMPOSITION
2KCl(aq)
+
F2(g) 
2NaF(aq)
5O2(g) 
3CO2(g)
+
+
Cl2(g) SINGLE REP.
4H2O(l) COMB.
Answer the following questions based upon the formulas above.
1. Rewrite #4 as a word equation. Dissolved silver nitrate and dissolved magnesium chloride
react to form solid solver chloride and dissolved magnesium nitrate.
2. What does the subscript “(g) “in #1 and subscript “(s)” indicate in equation 2? Gas; solid
3. Does equation 3 take place in water? How do you know? Yes; the NaF is dissolved
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4. Write the formula for the precipitate in equation 4. Rewrite its name in word format.
AgCl↓, silver chloride
5. Write the 3 criteria for identifying a combustion reaction.
Hydrocarbon and oxygen are
reactants; products are always CO2 and H2O
6. What does “∆” indicate in equation 2? How can you indicate a catalyst in a reaction?
Heat;
symbol will be written over the arrow.
REACTIONS STOCHIOMETRY and Limiting Reactants:
Perform the following stochiometric calculations. Make sure you have a balanced equation first!
CH4
+
2O2

CO2
+
2H2O
How many grams of O2 are needed to react with 9.02 g of CH4? 36.1 g O2
2.
N2
+
3H2

2NH3
How many grams of NH3 are formed when 11.7 g of H2 react? 65.9 g H2
3.
2C
+
O2

2CO
How many grams of carbon are needed to produce 0.98 g of CO? 0.82 g C
4. Silicon dioxide (quartz) reacts with hydrogen fluoride to produce silicon(IV)fluoride and water by the
reaction SiO2 + 4HF  SiF4 + 2H2O. Determine which is the limiting reactant if 2.0 mol HF is
combined with 4.5 mol SiO2? 2.0 mol HF is limiting
5. Using the reaction above, determine which is limiting and which is excess when 46.3 g of SiO2 react
with 25.0 g of HF.
25.0 g HF is limiting
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GRAPH OF TEMPERATURE VS. TIME FOR A SUBSTANCE BEING HEATED
T
E
M
P
E
R
A
T
U
R
E
(°C)
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
E
D
B
C
A
TIME (seconds)
Phase Changes – temperature vs. time:
Phase changes (step diagram)
Indicate the following on the graph above.
melting, vaporization, solid being heated, liquid being heated, gas being heated, freezing, condensation
Answer the following questions based upon the diagram above.
1. Which interval(s) on the graph above indicate a change in kinetic energy? A—B; C-D and E above
2. Which intervals on the graph above indicate a change in potential energy? B-C and D-E
3. At lines A-B and C-D the graph shows plateaus - the temperature is not changing while heat energy is
still being added. Describe what the heat is doing to particles of the substance. Breaking bonds
4. Explain the difference between vaporization and evaporation. Give an example for each. Vaporization
(boiling) is where the substance becomes liquid from inside and the surface. Evaporation is vaporization from the surface only.
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GASES:
Name and describe the three types of particle motion. Indicate the phases of matter that exhibit
each.
Rotation (spinning in place) – all phases
Vibration (shaking in place) – all phases
Translation (moving from place to place) – Liquid and Gas only
Describe the “Kinetic Theory of Gases” and list the three assumptions associated with it. What
volume does one mole of any gas occupy at STP? 22.4 L
Kinetic theory states that all matter is composed of particles and the particles are in constant motion. Particles are
small hard spheres which are not attracted or repelled from each other. They move in straight lines until they impact
with something. Between the particles is empty space.
Define “phase change.” This is a physical change which involves changes in the bonds between particles, types
of particles, and the types of particle motion. No chemical bonds are made or broken.
Define these types of phase changes. Give an example of where you might experience each.
Melting - solid to liquid
Vaporization - liquid to gas
Freezing - liquid to solid
Condensation - gas to liquid
Evaporation - liquid to gas (surface only)
Sublimation - Solid to gas
GAS LAWS:
Solve the following problems using the gas laws
1.
A gas with volume of 5.3 L at 17°C decreases to -3°C. Find the new volume.
2.
A 2.1 L sample of gas a standard pressure is moved to a 5.2 L container. Find the new pressure.
3.
An aerosol can contains gas at 15 atm and 25°C. Find the pressure inside the can at 100.°C
4.
Find the number of moles of gas in a container at the following conditions – 6.0 atm, 3.6 L, and
118°C.
5.
A gas at 742 mmHg and -18°C occupies 1.7 L. Find the volume of the gas at STP.
6.
Define absolute zero. Give the temperature for it in both °C and K. -273°C and 0K; temp at
which all particle motion ceases
7.
A mixture of four gases is stored at 1187.3 mmHg. H2 is 127 mmHg, O2 is 350 mmHg and CO2 is
642 mmHg. Find the partial pressure of the Helium in the mixture. 68.3 mmHg
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8.
Under what conditions does a gas stop acting “ideal” and begin acting “real?” low temps / high pressure
9.
What volume does one mole of any gas occupy at STP?
22.4L
CONVERSIONS
Perform the following conversions.
1. 26C to K
299K
5. 23.5 atm to mmHg
2. -52C to K
325K
6. 89 kPa to Torr
3. 89 K to C
184°C
7. 768 mmHg to atm
4. 304 K to C
31°C
8. 11.2 atm to kPa
17900 mmHg
670 Torr
1.01 atm
1130 kPa
WATER AND SOLUTIONS:
Draw a picture of a water molecule. Indicate the poles on this polar molecule.
Write a definition and example for each of the following.
Polar – charges on the water molecules (O is slight
negative while H is slight positive)
Specific heat – amount of energy required to raise
Surface tension – “skin” on water (water strider)
Hydrogen bonds – bond b/w adjacent water molecules b/w
– O and + H (drops on a penny)
one g of water by 1°C (1 cal)
Surfactant - breaks H-bonds (soap)
a) Why does solid ice float on liquid water? H-bonds hold water molecules apart;
increases spaces between molecules
b) How does soap help to clean dirty dishes? Soap is a surfactant which breaks Hbonds and allows the soap and oil to mix
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MOLAR MASS AND PERCENT COMPOSITION
Find the gram formula mass for the following formulas
BaCl2
208.236 g/mol
CuSO4
159.608 g/mol
NaNO3
84.994 g/mol
Ba(OH)2
171.3438 g/mol
Determine the percent composition of the following
S in CuSO4 20.1%
NO3 in NaNO3 72.9%
H2O in CuSO4 * 5 H2O 36.1%
Empirical and Molecular Formulas
A compound sample is made of 40.68 g C, 5.08 g H and 54.24 g O. Its molar mass is 118.1 g/mol.
Determine the empirical and molecular formula.
A compound sample contains 174.86 g iron and 75.14 g oxygen. Molar mass is 160.0 g mole.
Determine the molecular formula.
Acids and Bases
What helps you to recognize an acid when you look at a chemical formula? What are the general characteristics of
acids?
H (hydrogen) ion; smell strongly, corrosive, react with bases and metals; conduct electricity
What helps you to recognize a base when you look at a chemical formula? What are the general characteristics of
bases?
OH (hydroxide) ion; slippery corrosive; react with acids
What two products are always formed from the neutralization of an acid and base? Water and salt
Write the chemical equation for the neutralization of hydrochloric acid and sodium hydroxide
HCl + NaOH  NaCl + H2O
Write the chemical equation for the neutralization of sulfuric acid and potassium hydroxide
H2SO4 + 2 KOH  K2SO4 + 2 H2O
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Data Analysis, Significant Figures and Sources of Error:
Students want to determine the percent of water in a compound. They heat the
substance in a watch glass in order to drive off the water. The raw data they
collected is found below. Use this information to answer the questions.
Item
CuSO45H2O
Mass of Empty Watch Glass
46.25 g
Mass of Watch Glass and Compound before
70.65 g
heating
CuSO45H2O
Mass of Watch Glass and Compound after
63.33 g
heating
CuSO4
a. What is the mass of the sample of CuSO45H2O they started with?
24.40g
b. What is the mass of the water lost?
7.32g
c. From the given data, calculate the percent water in the compound.
30.%
d. What is the mass of one mole of this entire compound?
249.684 g/mol
e. What is the mass of just the water that is trapped in 1 mole of this compound?
90.074 g/mol
f. From the two previous questions, determine the actual percent (accepted value)
of water that makes up this compound.
36.1%
g. Assuming that the accepted value of % water should be 36.1%, determine the %
error for this lab group’s results.
16.9 %
h. If their experiment showed the percentage of water to be too low, how might this
error have occurred?
Incomplete evaporation
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Determine the number of significant figures in each of the following numbers.
a. 13,000
2
b. 13,000.
5
c. 0.00013
2
d. 9.0 x 109
2
Solve the following correctly, and include the correct label and number of significant
figures in your response.
a. A box measures 1.613 m by 6.213 m by 5.15 m. Determine the volume of the box.
51.6 m3
b. You measured the mass of a beaker as 16.255 g. When you added water to the
beaker, you measured the mass as 36.04g. Determine the mass of the water in
the beaker.
19.79g
c. In B above, what should you have done to allow you to record the mass of the
water to the thousandth of a gram?
Read your mass decimal measurement to a thousandth of a gram. (0.001g)
State the rules for use of significant figures within calculations:
For addition and subtraction?
Round to the same number of decimal places as the addend with the fewest decimals
For multiplication and division?
Round to the same number of significant figures as the multiplicand or divisor with the
fewest significant figures.
That’s it! If you would like additional practice for any of the topics covered on this exam – see me!
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