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Chemistry
Mrs. Baxley, Mr. Harvey, Mrs. Ozment
FINAL EXAM REVIEW PACKET
Fall Semester 2012
This packet is optional and will not count for, replace or be used for ANY grade (or extra credit)
Chemistry Final Review Material
The format of the final exam is MULTIPLE CHOICE. This review packet is intended to inform you of the
content in each unit covered and allow you practice with the content from this semester. It is not intended to
address any specific test question on the final exam. Completion of this packet does not guarantee success on
the Final Exam, but practicing with the content is a great idea.
Unit 1: Scientific Method, Graphing, and Safety
DEFINE:
a. Chemistry b. scientific method
c. Dependent variable
d. independent variable
1. Identify the following pieces of lab equipment:
Test tube
Graduated Cylinder
Beaker Tongs
Bunsen Burner
Watch Glass
Test tube holder
Beaker
Crucible Tongs
Test tube rack
Crucible and lid
Evaporating Dish
Chemistry Final Review Material
1. Safety: Describe an important rule in lab that applies to the following items:
Example: Food: Food is never allowed in lab.
a. Goggles: Always on your eyes when in lab
b. Accidental Spills Tell etacher immediately
c. Clothing/ Shoes: - Shoes MUST cover entire foot, long pants, no holes, cover entire leg
d. Smelling Chemicals: - Waft gently towards nose
2. Which unit(s) of measurement are usually dependent variables? Which are most often independent variables?
Time is an independent variable types of measurements are usually independent varaiables.
3. Label each kind of graph shown and answer the following questions about the graphs
Pie Chart
Bar Graph
a. What percent of the sources of chlorine in the stratosphere are CFCs? 57%
b. During which month of the year does Jacksonville usually get the most precipitation? The least?
August, November
4. Sequence the following steps for plotting a line graph
____7___
____3___
____1___
____5___
____2___
____6___
____4___
a. Give the graph a title.
b. Choose the ranges for the axes.
c. Identify the independent and dependent variables.
d. Plot the data points.
e. Determine the range of the data that needs to be plotted for each axis.
f. Draw the "best fit" line for the data.
g. Number and label each axis.
Chemistry Final Review Material
Unit 2: Measurement
Uncertainty in Measurements, SigFigs, Sci Notation, Dimensional Analysis, Rounding Rules
DEFINE:
a. accuracy
b. precision
c. scientific notation
d. significant figures
1. In the measurement 0.503 L, which digit is the estimated digit?
a. 5
b. the 0 immediately to the left of the 3
c. 3
d. the 0 to the left of the decimal point
2. Which two of these are equivalent lengths?
a. 6000 cm
b. 0.0600 km
c. 60 mm
d. 0.600 m
3. Which of these is the smallest length?
a. 6000 cm
b. 0.0600 km
d. 0.600 m
c. 60 mm
4. Convert 600 kilograms to grams
600 kg * 1000g = 6 x 105
1 kg
5. Which of the following conversions is/are incorrect.
a. 1 kilometer = 1000 meters
c. 500 centimeters = 0.5 meters
b. 100 millimeters = 1 centimeter
d. 1 meter = 1000 millimeters
6. The measure of the amount of three-dimensional space that an object occupies is known as:
a. volume
b. density
c. weight
d. mass
7. Convert 350. mL to Liters
350. mL * 1L___ = .350 L
1000 mL
8. Which metric distance is equal to 0.62 miles?
a. One centimeter
b. One millimeter
c. One meter
d. One kilometer
9. In the metric system, the base unit for mass is the:
a. Gram
b. Meter
c. Liter
d. Pound
10. Convert: 500 meters to kilometers
500 m * 1 km = 0.5 km
1000 m
11. Convert 5.0 x 103 mL to liters
5.0 x 103 mL * 1L
=5L
1000 mL
12. How many milliliters are in one deciliter?
a. 1 million, or 106
b. 1 thousand, or 103
c. 100, or 102
d.
10, or 101
Chemistry Final Review Material
13. In the metric system, the base unit for length is the:
a. yard
b. foot
c. mile
d. meter
14. Convert 75 mL to cm3
75 mL * 1cm3 = 75 cm3
1 mL
15. Convert 0.050 kilograms to grams 0.050 kg * 100,000 g = 5 x 103 g
1 kg
16. Which of the following are equivalent lengths?
a. 45 000 kilometers
b. 45 000 millimeters
c. 45 meters
17. Which of the following is the largest mass?
a. 800 centigrams
b. 5 kilograms
c. 1 000 milligrams
d. 450 centimeters
d. 100 grams
18. Determine the value of the missing measurement. SHOW YOUR WORK! USE UNITS!
mass = 75 g; volume = 10 cm3; density =
d = 7.5 g/cm3
mass = 400 g; density = 15 g/cm3 ; Volume = 15 g/cm3 = 400 g/x cm3 = 27 cm3
volume = 25 cm3; density = 5 g/cm3 ; Mass = 5 g/cm3 = x g/ 25 cm3 = 125 g or 100 g (sf)
19. Identify the number of sig figs in each of the following:
520 mL _2______
10.002 ns __5_____
0.0102 ms__3_____
0.4051 Pa _4_____
0.230 kg ___3____
0.001 cm _1______
256,000 L __3____
23.0 m _3______
20. Calculate using sig fig rules: (Review sigfig rules in calculations http://www.phys.unt.edu/PIC/significant_figures.htm)
a. 0.3287g x 45.2g =
14.9 g2
b. 0.258 mL / 0.36105 mL =
c. 12.5kg + 52.68 kg + 2.1 kg = 67.3 kg
0.715
d. (1250 cal – (234.207 cal / 52.69 cal) = 1246 cal
21. Write in scientific notation:
8960 __8.96 x 103_______
36,000,000 _3.6 x 107_________
0.00023 2.3x 10-3_________
0.000 000 025 3 2.5 x 10 -8___________
86,000 __8.6 x 101_________
2.04 x 103 ___2040________
1.23 x 10-5 _.0000123_______
1.20 x 10-2 ___.0120________
Chemistry Final Review Material
22. Determine the numeric measurement (include units) associated with each of the following pictures:
54.0 mL
11.23 mL
73.0 mL
Unit 3: Matter
Types of Matter, States of Matter, Physical/Chemical Properties, Physical/Chemical Changes
DEFINE:
a. matter
b. homogeneous
c. heterogeneous
d. mixture
e. pure substance
f. solution
g. physical property
h. chemical property
i. phase changes/changes of state
Chemistry Final Review Material
Use the word bank below to complete each sentence.
Evaporation
Element
Compound
Condensation
Solution
Mixture
Freezing
Heterogeneous
Substances
Melting
Homogeneous
Atoms
1. __Condensation___ occurs when a gas becomes a liquid.
2. All matter is made up of tiny particles called __atoms___.
3. When a solid becomes a liquid, _melting_____ occurs.
4. An _element_____ is made up of only one type of atom.
5. __freezing___ changes a liquid into a solid.
6. A mixture is made up of 2 or more substances that are physically combined (and can be separated).
7. When a liquid becomes a gas, __evaporation_____ occurs.
8. A mixture that is uniform (evenly spread) throughout the sample is said to be _homogenous__. These types
of mixtures are also known as ___solutions___.
9. _Compounds__ are 2 more atoms chemically bonded in a definite ratio.
10. A mixture that has uneven distribution of 2 or more substances is called ___heterogenous__.
11. Matter is divided into 2 categories: __solutions_____ and Mixtures.
12. Use the words below to complete the concept map.
heterogeneous
mixtures
13.
salt-water mixture
solutions
sand-water mixture
water
1= mixtures
2 = water
3. heterogenous
4.sand-water
mixture
5. solutions
6. salt water
mixture
Chemistry Final Review Material
Identify the following as either a chemical or physical change:
a.
Torn car seat ___P____
f. Crack in the sidewalk ____P_______
b.
Car’s faded paint _C ______
g. Burning a log ___C ___________
c.
A car’s rusting hood _C_______
h. Crushing a pop can ___P________
d.
butter melting _ __P______
i. leaves changing color____C______
e.
alcohol evaporating___P _______
j. wood rotting ___C_________
14. Identify as an element or a compound:
a. water __compound____
c. helium __element__
b. carbon dioxide ___compound___
d. arsenic ___element______
15. Identify as a pure substance, heterogeneous mixture or a homogeneous mixture:
a. Alphabet soup___heterogenous____
d. salt___pure substance____
b. sea water ___heterogenous__
e. granite ___heterogenous___
c. air__if pure - homogenous___
f. sugar ___pure substance____
Unit 4: Atomic Structure
History of the Atom, Subatomic Particles, Isotopes, Average Atomic Mass
DEFINE:
a. atom
b. electron
c. proton
d. neutron
e. nucleus
f. isotope
g. atomic number
h. mass number
Chemistry Final Review Material
i. average atomic mass
1. Describe the evolution of the model of the atom from Democritus through the present day model.
2. How many protons and electrons are present in a vanadium atom? 23
3. How many protons and electrons are present in a nitrogen atom? 7
4. What is the name of the element that has atoms that contain 5 protons? Boron
5. Write the chemical symbol for the ion with 95 protons and 89 electrons. Am+6
6.
Write the chemical symbol for the ion with 33 protons and 36 electrons. As-3
7.
Where are protons located in an atom? What is its charge?
Nucleus, +1
8. Where are electrons located in an atom? What is its charge? Around the outside of the nucleus, -1
9. Where are neutrons located in an atom? What is its charge? Nucleus, neutral
10. What is the atomic number of helium? 2
11. What is the atomic mass of oxygen? 32.00 g/mole
12. Complete the table below.
Isotope
Symbolic
Notation
Hydrogen-1
Hydrogen – 3
Number of
Protons
Number of
Electrons
Number of
Neutrons
1
1
0
1
Oxygen - 16
1
8
2
8
10
36
Copper-65
29
29
Uranium = 235
92
92
143
13. The element copper, Cu, has two naturally occurring isotopes. 69% of all copper consists of atoms with 34
neutrons, 31% of all samples consist of samples with 36 neutrons.
a. What are the two mass numbers? ____63__ and ___65____
b. Calculate the average atomic mass of copper atoms.
.69(63 amu) + .31(65 amu) = 64 amu
Chemistry Final Review Material
Unit 5: Electron Configuration
Arrangement of electrons in the electron cloud
DEFINE:
a. principal energy level
b. valence level
c. sublevel
d. “s” block
e. “p” block
f. “d” block
g. “f” block
1. List the first 4 places electrons can be found in the electron cloud and identify how many electrons each can
contain.
S orbital – 2 electrons; P orbital – 6 electrons; d orbital – 10 electrons; f orbital – 14 electrons
2. Write the orbital diagram notation for:
Ca-20
F-9
Ni-28
Sb-51
3. Write the electron configuration/spectroscopic notation for:
Ca-20 1s2 2s2 2p6 3s2 3p6 4s2
F-9
1s2 2s2 2p5
Ni-28 1s2 2s2 2p6 3s2 3p6 4s2 3d8
Sb-51 1s2 2s2 2p6 3s2 3p6 4s2 3d104p6 5s24d105p3
Chemistry Final Review Material
4. Write the noble gas notation for:
Ca-20 [Ar] 4s2
F-9
[He] 2s22p5
Ni-28 [Ar]4s23d8
Sb-51 [Kr] 5s24d105p3
5. How many valence electrons are present in each of the above atoms?
Ca – 2, F – 7, Ni – 2, Sb - 5
6. How many electrons can be placed into an “s” sublevel? Into a “p” sublevel? Into a “d” sublevel? Into an “f”
sublevel?
S = 2 electrons, p = 6 electrons, d= 10 electrons, f = 14 electrons
7. Which elements will gain electrons to form an ion? Where are these elements located on the Periodic Table?
Non-metals, top right, right of the staircase
8. Which elements will lose electrons to form an ion? Where are these elements located on the Periodic Table?
Metals, to the left of the stairstep line
Unit 6: Periodicity
History of the Periodic Table, Characteristics of Elements, Periodic Trends
DEFINE:
a. period
b. group/family
c. metal
d. nonmetal
e. metalloid
f. electronegativity
g. atomic radius
h. ionic radius
i. ionization energy
1. Answer the following about the Periodic Table:
Chemistry Final Review Material
a. What is the special group name in which Sodium is found? Alkali Metals
b. What is the special group name in which Fluorine is located? Halogens
c. What is the special group name in which Neon is found? Noble Gasses
2. Use the word bank to fill in the appropriate term below.
Groups
Hydrogen
Actinides
Ductile
periods
Inert
Halogens
malleable
metals
right
Alkali Metals
Transition Metals
nonmetals
metalloids
left
staircase
Noble Gases
Lanthanides
Alkaline Earth Metals
a. Metals are found primarily on the __left__ side of the periodic table (with the exception of __Al, ____).
Nonmetals are found on the __right__ side of the table.
b. Columns in the periodic table are called __groups________ and the rows are called ___periods_______.
c. The _alkali metals____ are found in group 1. They are the most reactive metals. Group 2 metals are
called the__alkaline__ __earth___.
d. The last group (18) are called the ___Noble Gasses____. They are non-reactive or ____inert____.
e. Salt forming compounds come from group 17, the_halogens________. They are the most reactive nonmetals.
f. Metalloids are found along the “___stair step line___”, and are often used as semiconductors, like
computer chips.
g. Groups 3 – 13 are called the __transition metals__ and include the 2 series below the table, the
__Lanthadines___ &__Actadines_____.
h. ___Metals______ have high luster and can conduct electricity and heat. They are also __malleable___ and
_ductile____.
i. __Non-metals_ have no luster, and are poor conductors of heat and electricity.
j. _Metalloids___ have properties of both metals and nonmetals.
3. Determine the charge assigned to each of the following ions:
Aluminum ___+3 ______ Chloride ___-1 _____ Copper(II) ___+2 ______
Zinc ___+2 _____
Magnesium ___ +2_____ Sulfur ___-2_ ____ Phosphide ___-3 ______
Silver ___+1 _____
4. How does the size of an ion compare to the size of the neutral atom from which it was created? Ions are bigger
5. How does an atom’s position on the periodic table provide information on that atom’s size (atomic radius)?
The farther left in the period, the larger the atom, the further down a group it is, the larger it is.
Chemistry Final Review Material
6. What is electronegativity and why do nonmetals have high values for it?
How attractive an element can make itself to an electron, non metals want to gain electrons so they will have
high electronegativity values
7. Describe the difference in the tendencies of metals and nonmetals to form ions (which is more likely to form a
cation and which is more likely to form an anion)
Metals will form cations and non metals will form anions
8. Which atom is the largest?
a. Li
b. B
c. N
d. F
e. He
9. Which atom is the largest?
a. Li
b. K
c. Rb
d. Cs
e. Kr
10. Which atom has the largest ionization energy?
a. Li
b. B
c. N
d. F
e. He
11. Which atom has the largest ionization energy?
a. Li
b. K
c. Rb
d. Cs
e. Kr
Unit 8: Ionic Compounds
Transferring electrons, ions, nomenclature
DEFINE:
a. ionic bond
b. ionic compound
c. polyatomic ion
d. nomenclature
e. stock system
f. traditional system
g. empirical formula
1. Name or write the formulas for the following:
potassium iodide
KI
KOH
potassium hydroxide
barium chloride
BaCl2
LiI
lithium iodide
lithium bromide
LiBr
AlF3
aluminium fluoride
Chemistry Final Review Material
iron(III) sulfate
Fe2(SO4)3
chromium(III) sulfide
Cr2S3
calcium carbonate CaCO3
cobalt(II) fluoride
silver oxide
FeCl2
iron (II) chloride
MgO
magnesium oxide
Co(NO3)2 calcium nitrate
CoF2
Ag2O
Zn(PO4)2
zinc phosphate
(NH4)SO4
ammonium sulfate
magnesium hydroxide
Mg(OH)2
NO
nitrogen monoxide
Nickel(II) nitrate
Ni(NO3)2
AgI
silver iodide
Unit 9: Covalent Compounds
DEFINE:
a.
molecule
b. Lewis structure
c. single bond
d. double bond
e. triple bond
f. multiple bond
g. VSEPR theory
h. molecular geometry
i. polar bond
j. polar molecule
k. nonpolar molecule
1. Name the following molecules:
SiO2
silicon dioxide
PCl3
phosphorous trichloride
SiF4
silicon tetraflouride
N2 O
dinitrogen monoxide
SO3
sulfur trioxide
N2O5
dinitrogen pentaoxide
Chemistry Final Review Material
P3O4
diphosphorous tetraoxide
2. Write the formula or name for these compounds.
dinitrogen pentoxide N2O5
phosphorus trichloride
selenium monofluoride
PCl3
SeF
silicon hexacarbide
SiC6
trisulfur heptabromide S3Br7
carbon monoxide
CO
carbon dioxide
CO2
carbon tetrachloride
CCl4
diphosphorus octoxide P2O8
3. How are bonds different in ionic and covalent compounds.
Ionic bonds – transfer of electrons
Covalent Bonds – Sharing of electrons
4. Why do most atoms form bonds? They become more stable due to exsisiting at a lower energy level
5. Which nonmetal groups on the periodic table can form double bonds?
Oxygen family
6. Which nonmetal elements can form triple bonds? Nitrogen family
7. Which nonmetal elements can only form single bond? Halogens
8. What major assumption of the VSEPR theory means that bond angles will be as large as possible and that compound
will exist in 3 dimensional space?
Unpaired electron clouds push ligands away from central atom.
9. In H2S, two hydrogen atoms are bonded to one sulfur atom. Why isn’t the molecule linear?
There are 2 unbonded pair of electrons on the central atom.
11. What two factors determine whether or not a molecule is polar? The polarity of the bonds on the central atom
and the symmetry of the molecule.
12. Draw the Lewis Structure for the following molecules, then determine if the molecule is polar or nonpolar.
HCN
N2
HF
H 2S
Chemistry Final Review Material
NH3
Unit 10: Chemical Reactions
Words to formulas, Balancing, types, predicting products, Activity Series, Solubility Rules
DEFINE:
a. law of conservation of matter
b. reactant
c. product
d. coefficients
e. precipitate
f. soluble
g. insoluble
h. activity series
1. List the observable indications that a chemical change has occurred.
Generation of heat/cold, light, solid from liquid, smoke, odor
2. Use the word bank to match the correct term with its definition.
decomposition reaction
Law of Conservation of Mass
double replacement reactions
neutralization reaction
___products__
products
single replacement reactions
reactants
synthesis reactions
coefficients
combustion reactions
a. compounds found after the yield sign (right side)
organic combustion_____ b. reaction using oxygen to form carbon dioxide, water, and energy
Law of Cons. Of Mass c. the number and type of atoms on each side of a chemical reaction must be balanced (matter is
not created or destroyed!)
double replacement___
__synthesis
d. reaction where 2 elements “switch” places to form 2 new compounds
e. reaction that forms a more complex substance
Coefficients____f. numbers placed before a compound to balance the equation
_Reactants_____g. compounds found before the yield sign (left side)
Chemistry Final Review Material
_Single replacement______ h. reaction where one element replaces another
decomposition___ i. reaction that breaks down a complex substance
3. Classify each reaction by type and balance them:
_D___________ a.___2__H2O2  ___2__H2O + _____O2
__S__________ b. ___4__Fe + __3___O2  2 Fe2O3
___SR_________ c._____Fe2O3 + __3___H2  _2____Fe + 3_____H2O
_____SR_______ d. _____Zn + _____H2SO4  _____ZnSO4 + _____H2
4. Predict the products of the following reactions and classify the reaction by its type:
a. NaCl + KNO3  NaNO3 + KCl
b.2 Na + O2  2 Na2O
S
c. Cl2 + KF  NR
SR
d. H2O  H2 + O2
D
DR
5. Balance the following equations:
a.__ 2__ AgNO3 + ___H2S  ___Ag2S +__2_ HNO3
b.___MnO2 + __4_HCl  ___ MnCl2 + _2___H2O + ___Cl2
6. Predict end products and write balanced chemical equations and use (aq), (g), (l) or (s) after each formula.
a.
Sodium chloride reacts with lead (II) nitrate to produce…
NaCl (aq) + Pb(NO3)2 (aq)
NaNO3 (aq) + PbCl2(s)
b. Calcium nitrate reacts with sodium sulfide to produce…
Ca(NO3)2 (aq) + Na2S (aq)
NaNO3 (aq) + CaS (aq)
c.
Ni + 2AgCl  NiCl2 + 2Ag
SR
Chemistry Final Review Material
d. 2 Fe + O2 2 FeO
S
e. CH4 + 2O2  CO2 + 2H2O
OC
f. Ca(CO3)2  Ca (s) + 3CO2 (g)
D
g. H2O + KCl  HCl + KOH
DR
Unit 11: The Mole
DEFINE:
a. mole
b. molar mass/formula mass
c. percent composition
d. empirical formula
1. Calculate the molar mass of the following compounds: (Remember to add units)
3.
a.
CaCO3___________
b.
MgSO4___________
c.
NaOH___________
d.
KCl ___________
e.
Co(NO3)2________________
f.
Zn(PO4)2________________
Find the mass of 1.112 mol of HF.
Chemistry Final Review Material
4.
If you have 66.38 g of KMnO4 find how many moles it is made up of?
5.
How many moles of ethane (C2H6) contain 8.46 x 1024 formula units?
6.
How many molecules are in 5.1 g of H2O?
7.
What is the mass of 22.4 L of H2O?
8.
Find the % composition of H in H2O?
9. Find the % composition when 9.02 g of Mg combine completely with 3.48 g of N to form a compound.
10. Calculate the empirical formula for a compound made up of 94.1% O, and 5.9% H.
11. Calculate the empirical formula for a compound made up of 79.8% O, and 20.2% H.