Download Chemistry SOL Review Test

Name: _____________________
Chemistry SOL Review Test
Directions: Read each question and write a brief answer. Use your notes,
diagrams, etc. to help you.
 SHOW ALL WORK (no work = no credit)
 Your work and writing must be neat & legible (unreadable = no credit)
Unit 1: Measurement (Ch. 3)
1) Define the following
accuracy reflects how close a measurement is to
a) Accuracy
a known or accepted value
b) Precision
c) Quantitative Data
d) Qualitative Data
precision reflects how reproducible
measurements are, even if they are far from the
accepted value
quantitative analysis is used to tell 'how much' is
in a sample
data that determine the characteristics of a
substance chemical constituents.
2) Number of Significant Figures
a) 1.23 X 107
3
b) 100.1010
7
c) 33,000
2
d) 0.00004
1
3) Perform the calculations indicated. Round your answer to the
correct number of significant figures.
1500
a) 632,000  410
b) 100.0 + 45.0 + 35
180.
c) 98.012 +16
114
d) 891.000 * 1.0
890
4) Put the following numbers in proper scientific notation.
-5
a) 0.0000330
3.30 x 10
b) 95,200,000
9.52 x 107
5) Take the following numbers out of scientific notation.
a) 3.458 X 10-11
0.000 000 000 03458
10
b) 2.4 X 10
24,000,000,000
6) Answer the following: Be sure to put your answers in
PROPER SCIENTIFIC NOTATION
a) (3.10 x 103)+(1.68 x 102)
3.27 x 103
b) (1.11 x 108) – (1.1 x 103)
1.1 x 108
5
-9
c) (8.00 x 10 ) / (4.00 x 10 )
2.00 x 1014
d) (600 x 1025)*(0.00020 x 10-10)
1 x 1014
7) Use the solving problems method to answer the following
questions. SHOW ALL OF YOUR WORK.
a) How many micrograms are in 3.00 grams?
3.00g
1,000,000 μg =
3.00 x 106 μg
1g
b) How many seconds are in 24 days?
24 days 24 hrs 60 min 60 sec =
1 day
1 hr
1 min
c) How many grams are in 9.42 mL of
pentane (C5H12)?
(Density = 0.0980 g/mL)
9.42 ml
0.0980 g =
1 mL
d) How many milligrams are in 85
kilograms?
85 kg 1,000 g 1,000,000 μg =
1 kg
1g
2,073,600 sec
2,100,000 sec
0.923
g C5H12
8.5 x 1011 μg
Unit 2: Matter & Energy (Ch. 2)
8) Classify the following as either a physical or chemical change
a) Pizza is sliced
Physical change
b) Paper is torn
Physical change
c) NaCl is dissolved
Physical change
d) Wood burns
Chemical change
9) ID the following as either a heterogeneous mixture or a
homogeneous mixture.
a) Italian dressing
heterogeneous mixture
b) NaHCO3 solution
homogeneous mixture
c) Muddy solution
heterogeneous mixture
10) Predict whether the element is a metal, nonmetal, or metalloid.
nonmetal
a) selenium
metal,
b) magnesium
nonmetal
c) chlorine
d) antimony
metalloid
11) List four different characteristics of metals & nonmetals
Nonmetal
Metal
Not Ductile (Brittle)
Ductile
Low MP and BP
High MP and BP
Tend to be volatile
Luster (shiny)
Insulators
Conductors
12) Describe the following changes of state
a) Sublimation
Solid to Gas
b) Freezing
Liquid to solid
c) Vaporization
Liquid to Gas
d) Condensation
Gas to Liquid
e) Melting
Solid to Liquid
13) Differentiate between gases, liquids, & solids.
solids
liquids
gases
Slow
Faster
Fastest
Particle speed
close together no Far apart no
Particle spacing Packed
Volume
Regular
arrangement
Fixed
regular
arrangement
Fixed
regular
arrangement
Variable
(fixed or variable)
Shape
Definite
Not
Not
(definite or not)
Most dense
Somewhat dense
Least dense
Density
No
Not
easily
Yes
Compressible?
14) Identify the group/family number/s for the following:
a) Noble gases
18
b) Halogens
17
c) The Alkali Metals
1
d) Alkaline Earth Metals
2
e) Transition Metals
3-12
Unit 3: Atomic Structure and Theory (Ch. 4)
15) Describe the Dalton’s atom model. Atoms are indivisible and
indestructible. (Billiard ball model)
16) Describe Thomson’s atom model. Electron are like raisins
surrounded by a soup of positive charge to balance the electrons'
negative charges (Plum Pudding Model).
17) What experiment did Thomson do? The cathode rays tube
18) What was his model called?
Plum Pudding Model
19) Describe Rutherford’s model. Atoms have a center nucleus and
the rest of the atom is mostly empty space
20) What experiment did Rutherford do? He shot alpha particles
through a thin film of gold and to his amazement a few alpha
particles rebounded almost directly backwards.
21) What was his model called? “gold foil" experiment
22) Describe Bohr’s model. He proposed his quantized shell model of
the atom, electrons orbits around the nucleus
23) What was his model called? Planetary Model
24) What is the difference between an isotope and an ion?
Isotope
Ion
Different number of Neutrons
Different number of Electrons
25) Use your periodic table to fill in the following table.
(HINT: These are isotopes)
Protons Neutrons Electrons Atomic Mass
number number
33
44
33
33
77
a) As -77
153
74
54
53
127
b) I
30
35
30
30
65
c) Zn
2+
26
30
24
26
56
d) Fe
5
6
5
5
11
e) boron-11
26) A cation is positive “+” in charge
a anion is negative “‒“ in charge.
27) A horizontal row in the periodic table is called a Periods
28) A vertical column is called a Groups or Families
H
29) proton
B
F
G
30) isotope
31) electron
32) neutron
D
33) nuclear pull
E
34) shielding
a. the smallest particle of an element that retains the
properties of that element
b. atoms of the same element with different masses
c. the number of protons in the nucleus of an atom
d. attraction of positive nucleus for the outer negative
electrons
e. electrons in between the outer electrons and the nucleus
shield the positive nucleus from the negative outer electrons.
f. the first subatomic particle discovered
A
C
35) Atom
36) atomic
number
g. subatomic particle with no charge
h. .the kind of subatomic particle with a positive charge
37) The element iron is made up of the isotopes 58Fe (2.45%), 56Fe
(80.3%), 54Fe (10.2%), and 57Fe (7.05%). Calculate the average
atomic mass (weight) of iron.
Answer: 58 (0.0245)+56(0.803)+54(0.102)+57(0.0705) = 55.92
Unit 4: Electrons in atoms (Ch. 5)
38) Describe the following prop
39) erties and their periodic trend (you may draw a picture for
the trend)
Definition
Trend
a) electron affinity
→ increases
D
H
A
C
F
E
B
G
40) 1st quantum #
41) p orbital
42) Atomic spectrum
43) Electromagnetic
spectrum
44) 3rd quantum #
45) 2nd quantum #
46) d orbital
47) 4th quantum #
a) the characteristic spectrum of an element
b) complex shaped
c) A listing of all known forms of radiation
d) describes how far the electrons is from the nucleus
e) describes the shape of the orbital
f) describes the orientation of the orbital
g) describes the direction the electron is spinning in
h) dumbbell shaped
48) Which
color has the shortest wavelength? ROY G BIV (longest to shortest
so Violet)
b) electronegativity
→ increases
49) Define the following terms.
Frequency
The amount of time it takes for one cycle to complete
Amplitude
The distance from the equilibrium to the crest or to the trough
Wavelength The distance from Crest to Crest or Trough to Trough
c) ionization energy
→ increases
50) Write the complete electron configuration for the following :
2
2
6
2
6
2
2
a) Titanium (Ti). 1s , 2s 2p , 3s 3p , 4s , 3d
d) atomic radius
← increases
e) ionic size (cations, anions)
f) Atomic number
g) valence electrons
h) atomic mass
i) electron configuration (s, p, d, f blocks)
→ increases
→ ↓ increases
→increases
→ ↓ increases
2
2
6
2
6
b) Si-4 1s 2s 2p 3s 3p normally it would be 3p2 but there are 4
-4
additional electrons Si
giving it a noble gas configuration
Unit 5: Nomenclature (Ch. 9)
51) Name the following compounds: (molecular or ionic)
HINT: Molecular = 2 nonmetals – use prefixes; Ionic = anything else – just state
the name of the cation then anion (use roman numerals if it is a cation w/
multiple charges)
a) CaS
b) PO3
Calcium Sulfide
Phosphorus trioxide
c) N2O6
d) (NH4)3CO3
Dinitrogen hexoxide
Ammonium carbonate
e) C2I4
f) Fe2O3
Dicarbon tetraiodide
Iron (III) oxide
52) Write the formulas for the following compounds
b) rubidium carbonate
a) lead (IV) chloride
Rb2CO3
PbCl4
c) silicon monoxide
d) sodium arsenide
SiO
Na3As
e) strontium nitrate
Sr(NO3)2
Unit 6: The Mole (Ch. 10)
f) silver nitride
Ag3N
Use the solving problems method to answer the following questions.
SHOW ALL OF YOUR WORK and use significant figures.
53) Find the molar mass of Sr3(PO3)2
Sr 87.62(3)+ P 30.97(2) + O 16.00(6) = Answer: _420.80_
54) How many formula units are in 2.84 moles of potassium oxide
(K2O)?
2.84 mol K2O 6.02 x 1023 fu = Answer: 1.71 x 1024 fu K2O
1 mol K2O
55) How many moles are in 8.50 X 1023 molecules of water?
8.50 x 1023 mlcs H2O
1 mol H2O
= Answer: 1.41 mol H2O
6.02 x 1023 mlcs H2O
56) How many grams are in 4.42 moles of methane (CH4)?
4.42 mol CH4
16.05 g CH4
= Answer: 67.4 g CH4
1 mol CH4
57) How many moles are in 1.2 X 104 liters O2 gas at STP?
1.2 x 104 L O2
1 mol O2
22.4 L O2
= Answer: 540 mol O2
58) How many grams are in 0.9500 liters F2 (Fluorine Gas)?
1 mol F2
38.00 g F2
= Answer: 1.612 g F2
0.9500 L F2
22.4 L F2
1 mol F2
59) How many atoms of nitrogen are in 56.652 grams of nitrogen?
56.652 g N2 1 mol N2 6.02 x 1023 mlc N2=
Answer: 1.2171 x1023 mlc N2
28.02 g N2 1 mol N2
60) Find the percentage composition of a compound that contains
19.5 g of iron and 5.8 g or sulfur. The total mass of the compound
is 20.3 g.
Answer: omit_
61) 11.66 g iron, 5.01 g oxygen. Find the empirical formula.
11.66 g Fe 1 mol Fe
55.85 g Fe
5.01 g O 1 mol O
16.00 g O
= 0.20877 / 0.20877 = 1 x 2 = 2
= 3.13125 / 0.20877 = 1.5 x 2 =3
So the empirical formula is Fe2O3
62) What is the molecular formula? empirical formula NO2,
molar mass = 92.02 g/mol
N 14.01(1) +16.00(2) = 46.01
92.02/46.01 = 2
So there are 2 (NO2); so the molecular formula is N2O4
63) What is used to convert molecules to moles?
a) Name: Avogadro’s number
b) Number: 6.02 x 1023 mlcs/mole
Unit 7: Chemical Equations (Ch. 11)
64) Use the activity series of metals to determine which of the
following reactions will occur. If a reaction will take place,
complete and balance the equation. If the reaction will not occur,
write no reaction.
a) 2 Al + 3 CuSO4  3 Cu + Al2(SO4)3
b) 6 Ag + 2 H3PO4  3 H2 + 2 Ag3PO4
65) Determine the products of the following double replacement
reactions.
a) 3 Ag2CO3 + 2 Na3PO4  2 Ag3PO4 + 3 Na2CO3
b) Pb(NO3)2 + CoCl2  PbCl2
+ Co(NO3)2
66) What are the meanings of the following symbols?
a) (aq)
aqueous
b) ↑
given off (gas)
c) ↓
Precipitate (solid)
d) ∆
heat
67) Directions: Balance the following reactions. To the right of
each equation, write the type of reaction each represents. (Use S,
D, S-R, D-R, C rather than the name)
a) Zn + HCl → ZnCl2 + H2
Single Repl
b)
C7H6O2 + O2 → CO2 + H2O
c) Sb + O2 → Sb4O6
Combustion
Synthessis
d) H2O2 → H2O + O2
Decompostion
e) Al2(SO4)3 + Ca(OH)2 → Al(OH)3 + CaSO4
Double Repla
Unit 8: Stoichiometry (Ch. 12)
68) Define the following:
It is the reactant that is completely used up
a) Limiting Reagent
b) Actual Yield
c) Percent Yield
d) Theoretical Yield
in a reaction
The experimental yield, what you actually
get.
(Actual yield/Theoretical yield) x 100.
Percent yield measures how efficient the
reaction is under certain conditions.
The amount of product that could possibly
be produced in a given reaction.
Answer: _________
73) Sodium chloride can be prepared by the reaction of sodium
metal with chlorine gas. 2Na (s) + Cl2 (g)  2 NaCl (s)
a) When 6.70 mol of Cl2 reacts with 3.20 mole of Na, what is
the limiting reactant?
Answer: __________
b) How many moles of NaCl are produced?
Problems:

Make sure that your chemical equation is balanced.

Set up a proportion under the substances in the equation that you are
interested in.
Divide by the coefficient to cancel out moles, the (coefficient * molar
mass) to cancel out grams, or the (coefficient * 22.4) to cancel out
liters

69) How many moles of aluminum oxide are needed to react
completely with 5 moles of SnO?
2Al +
3SnO  3Sn + Al2O3
5 mol SnO 1 mol Al2O3 = 1.7 mol Al2O3
3mol SnO
70) How many moles of sodium are needed to react with 18.3 g of
Sodium chloride?
2 Na + Cl2  2 NaCl
18.3g NaCl 1 mol NaCl 2 mol Na = 0.313 mol Na
58.44 g NaCl 2mol NaCl
Answer: _________
c) How much of the excess reagent remains unreacted?
Answer: _________
74) Calcium carbonate is decomposed by heating.
CaCO3  CaO + CO2
What is the percent yield of this reaction if 27.8 g of CaCO3 is
heated to actually produce 13.6 g of CaO?
Answer: _________
71) If 48.3 g of Fe2O3 react with Aluminum in a single
displacement reaction, how many grams of aluminum are needed?
2 Al + Fe2O3  Al2O3 + 2 Fe
48.3g Fe2O3 1 mol Fe2O3 2 mol Al = 0.605 mol Al
159.7 g Fe2O3 1mol Fe2O3
75) 2 Al + 2 H3PO4  2AlPO4 + 3H2
What is the percent yield if 6.6 g of Al is reacted to produce 1.85 g
of H2?
72) If 40 L of oxygen combine with oxygen to synthesize water,
how many grams of water are produced? 2H2 + O2  2 H2O
Unit 9: GASES (Ch. 14 &Ch 13)
(don’t forget to change C to K)
76) What do moles tell you? What is the unit for n?
Answer: _________
77) What is volume? What are the units of volume?
78) What are the units for pressure and how are they related?
1 atm = 101.325 kPa = 760 mmHg = 760 torrs
79) Fill in the following table
GayAvogadro’s Charles’ Boyle’s
Lussac’s
Formula
V1 = V2 P1V1=P2V2
P1 = P2 V1 = V2
T1
T2 n1 n2
T1
T2
Inverse or Direct Directly
Directly
Directly
Inversely
Graph
R= 8.31 L*kPa
mol*K
; 0.0821 L*atm
mol*K
;
62.4 L*mmHg
mol*K
80) Many gases are available for use in the laboratory in
compressed gas cylinders, in which they are stored at high
pressures. Calculate the number of moles of O2 that could be
stored at 24.0 C and 0.935 atm, in a cylinder with a volume of
16.5 liters?
PV=nRT
Answer: _______________
81) A balloon contains 2.5 mol of helium at a pressure of 1.06 atm
and a temperature of 274. K. What is the volume of the balloon?
PV=nRT
Anwer: _______________
82) A helium balloon with a volume of 22.5 L, a pressure of
101.3.kPa, and a temperature of 328 K is put into an environment
where the pressure is 25.0 kPa and the temperature is 25.0C.
What is the new volume of the balloon?
Answer: _______________
83) The gas balloon has a volume of 4.8 L at 330 K. At what
temperature will the balloon expand to 8.50 L?
Answer: _______________
84) A 2.43 L balloon contains 0.35 moles of He. How much
Helium must be added in order for the balloon to be 5.2 L?
Assume constant temperature and pressure.
Answer: _______________
85) Calculate the temperature (C) of 2.65 g of I2 if it occupies 50.
mL at 676 mm Hg.
PV=nRT
676 mmHg 0.050 L = (2.56g/253.8 g/mol) (62.364 L mmHg mol-1 K-1) T
53 K or -219 C
Answer: _______________
Unit 10: Liquids (Ch. 15&16)
86) The solubility of most liquids increases as temperature
increases. (increases/decreases)
87) The solubility of gases in liquids increases as temperature
decreases.
88) Three colligative properties are:
1. Vapor pressure depression
2. Boiling point elevation
3. Melting point depression
4. Osmotic pressure
89) Four factors which affect solubility.
a)
b)
c)
d)
Temperature
Pressure
Solvent – Solute interaction (like and unlike)
Size
90) What is the molarity of a 1500 mL solution made with 7.5
moles of KCl?
7.5 mol KCl = 1.6 M
1.5 L
91) How many mL of a 0.01 M solution of NaCl would you need
in order to have 1.4 moles?
1.4 mol = 0.01 M NaCl
x = 140,000 mL
xL
92) How could you make a 2.5 M NaOH dilute solution in a 250
mL volumetric flask, from a 3 M concentrated solution of NaOH.
M1V1 = M2 V2 2.5 M 250 ml = 3 M x x = 208 mL
So place 208 mL of the 3M solution in a flask and add 42 mL of H2O
for 250 mL of 2.5 M solution.
102) Fill in the following chart
Lewis Dot
Diagram
CO
# of lone # of Molecular Bond
VSEPR
pairs
bonds Geometry Angles Structure
(name)
skip
skip
H2Se
93) If NaCl dissolves in water? Is NaCl POLAR or nonpolar?
Unit 11: Bonding (Ch. 8 & 7.3)
94) Write electron dot structures for the following ATOMS
a) silicon
b) S
95) Write electron dot structures for the IONS of the following
atoms
a) S-2
GeH4
AlF3
b) Mg+2
96) Use the electronegativity table to determine what type of bond
(IONIC, VERY POLAR COVALENT, MODERATELY
POLAR COVALENT, OR NONPOLAR COVALENT) will
form between the atoms of the following elements.
IONIC
97) Mg and F
MODERATELY POLAR COVALENT
98) C & H
POLAR COVALENT
99) C and O
IONIC
100) Na & F
MODERATELY POLAR COVALENT
101) H and Br
Unit 12: RATES OF REACTION & CHEMICAL EQUILIBRIUM (Ch.18)
103) Using the following equations for complex ions, draw an
arrow to show how each system would change under the
following conditions. (In which direction will the equilibrium
shift?)
[Zn(H2O)4]2+ + H2O ↔ [Zn(H2O)3(OH)]+ + H3O+
a. Removing water _________
b. Adding H3O+ _________
104) For the following reaction, equilibrium is established at a
certain temperature when the following concentrations are present:
[CO] = 0.10 mol/L, [H2O] = 0.80 mol/L, [CO2] = 0.012 mol/L,
and [H2] = 0.012 mol/L. Calculate the Keq value for this reaction.
HINT: Write the equilibrium expression for the reactions, then fill
in the values and solve for Keq
CO (g) + H2O (g) ↔ CO2 (g) + H2 (g)
Answer: ________________
Unit13: Acids, Bases, & Neutralization (Ch. 14)
105) In a neutralization reaction the products are always _______ and _______.
106) An acid turns litmus paper ______ and a base turns litmus paper _______.
107) Water is amphoteric because it can _____________________
108) pH = ___________[H+]
109) Label the Bronsted-Lowry acid, base, conjugate acid, and
conjugate base in the following reactions.
a) H2O + H3PO4  H2PO4-1 + H3O+
+
b) H2O + H2SO4 H3O + HSO4
-1
110) Name each of the following acids and their anions
Formula
Name of anion
Name of acid
a) HBr
b) H2CO3
c) H3PO3
d) H2SO3
111) Write the formulas for the following acids.
Name
Cation
Anion
Formula
a) Chloric acid
b) Phosphorous acid
c) hydronitric acid
112) Complete and balance the following neutralization reactions.
a) H2SO4 + Ca(OH)2  ______ + ________
b) HI + NH4OH  _______ + ________
113) Determine the pH for the aqueous solutions with the following
[H+]. Is the solution acidic, basic or neutral?
[H+]
pH
Acid, base, or
neutral
a) 1 X 10-6 M
b) 1 X 10-7 M
114) Determine the [H+] for aqueous solutions that have the
following pH values. Is the solution acidic, basic, or neutral?
pH
[H+]
Acid, base, or
neutral
a) 11
b)
8
115) Determine the pH of the following solutions that have the
following [OH-]. Is the solution acidic, basic, or neutral?
[OH-]
pH
Acid, base, or
neutral
a) 3.8 X 10-3
b)
0.000001
116) Would a stronger acid have a large or a small Ka?
117) In a titration, a 15 mL solution of hydrochloric acid is
neutralized by 6 mL of 2.5 M sodium hydroxide, using
phenophthalein as an indicator. What is the concentration of the
hydrochloric acid solution?
Answer: _____________
Unit 14: REDOX
(Ch. 20)
O.I.L.
R.I.G.
118) Oxidation is the loss / gain of electrons & increase/decrease of oxidation #.
119) Reduction is the loss / gain of electrons & increase/decrease of oxidation #.
120) Show the oxidation numbers for all of the atoms in the
following equation.
Al(OH)3 + H2SO4 ---> Al2(SO4)3 + HOH
121) In the following equation, which chemical is oxidized?
Reduced? Oxidizing agent? Reducing agent?
Zn + HCl ---> ZnCl2 + H2
Oxidizing Agent: __________
Reducing Agent: __________
Oxidized: _________
Reduced: _________
129) Write the half reactions for the oxidized and reduced
chemicals in previous problem.
Oxidation ½ reaction : ____________________________
Reduction ½ reaction: ____________________________
Unit 15: Thermochemistry (Phase Changes)(Ch. 17)
Heat energy = mcT (WITHIN A PHASE)
Heat Energy = m H (TO CHANGE PHASE)
4.18 J = 1 Cal & specific heat of liquid water = 1 cal/gC
130) A 16.3 gram water sample in a calorimeter has its temperature
raised 8.0 C while an exothermic chemical reaction is taking
place. How much heat is generated in the calorimeter? (solve for
calories and for joules)
Answer: ____________
Answer: ____________
131) How many joules would be needed to raise 20.0 grams of water
forms 10.0C to 98C?
Answer: ____________
132) If 5.0 grams of water at 45C gains 5600.0 Joules of heat,
what is the final temperature of the water? (Hint - solve for T)
Answer: ____________
133) 98.7. grams of a material is melted at it’s freezing point.
Heat of fusion = 472 J/g
Answer: ____________
134) 16.3 grams of a material is heated from 100oC to 200oC
Specific heat = 2.34 J/goC
How much heat is absorbed?
___ 122) the energy required to melt of mole of a
substance
___ 123) the energy required to change one mole of a
substance from a liquid to a gas.
___ 124) a measure of the randomness or disorder of a
system
___ 125) The total energy content of a system
___ 126) A substance that increases the rate of a
chemical reaction and that is not changed by the
reaction.
___ 127) The system gains energy
___ 128) The system releases energy
a. entropy
b. heat of
vaporization
c. heat of fusion
d. endothermic
e. exothermic
f. enthalpy
g. catalyst
Answer: ____________
Match the following:
135) Decide which of these reactions are endothermic and
exothermic,Your DONE!!
_____ a. BCl3 + 3 H2O  H3BO3 + 3 HCl H = -112kJ
_____ b. 2 Fe + 3 CO2 + 26.8 kJ  Fe2O3 + 3CO
_____ c. C2H4  2 H2 + 52.3kJ
_____ d. S2Cl2 + CCl4  CS2 + 3 Cl2 H = +112kJ