Download CP Chemistry - Final Exam Review KEY

Final Review 1
Name ______KEY __________________________ Test Date ______________ Period ______
CP Chemistry – Final Exam Review
Show your work and round using significant figure rules for all the problems.
1.
What is chemistry?
 Chemistry is the study of matter and the changes it undergoes.
2.
3.
4.
5.
Write the following in scientific notation:
a. .005
5 x 10-3
c. 53099000
5.3099 x 107
b. 5050
5.05 x 103
d. 0.00053030
5.3030 x 10-4
Write the following in standard notation:
a. 1.5x103
1500
c. 4x10-5
0.00004
b. 3.75 x10-2
0.0375
d. 9.12 x107
91200000
How many significant figures are in the following:
a. 0.02
1
c. 501
3
b. 501.0
4
d. 0.020
2
Calculate and report in correct significant figures:
a. 52.568 cm + 689.45 cm + 55.5694 cm = 797.59 cm
b. 8426. 200 g ÷ 254 mol =
6.
You measure a Calcium cube to have 2.01 cm sides and a mass of 11.6 grams:
a. What is the density as you calculate it?
v = s3 = (2.01 cm)3 = 8.12 cm3
d=
7.
33.2 g /mol
m
11.6 g
=
= 1.43 g/cm3
v
8.12 cm3
b. If calcium’s accepted density is 1.54 g/cm3, what is your percent error?
|1.54 g/cm3 – 1.43 g/cm3|
%error =
x 100 = 7.14 %
1.54 g/cm3
Perform the following conversions:
a. 53.1 mm to m
1m
53.1 mm x 3
= 0.0531 m
10 mm
b. 18.4 kg to dag
1 x 10-1 dag
18.4 kg x
= 1840 dag
1 x 10-3 kg
c. 93.6 mi to km (1 mi = 1.6 km)
1.6 km
93.6 mi x
= 150. km
1 mi
Final Review 2
8.
9.
10.
11.
What is the difference between precision and accuracy?
 Precision is the reproducibility of data, while accuracy is how close the
measurement is to a known value.
What two properties does all matter have?
 All matter has mass and volume.
What is the difference between an element, compound, and mixture?
 An element is the simplest form of matter, and one type of atom. A compound is
two or more atoms chemically combined. A mixture is two or more substances
physically combined.
Identify the following as either a(n) element, compound, homogeneous mixture, or
heterogeneous mixture:
a. Chlorine
Element
d. Steel
Homogeneous Mixture
b. Water
c. Soil
12.
13.
14.
15.
Compound
e. NaCl
Compound
Heterogeneous Mixture
f. Paint
Heterogeneous Mixture
Classify the following as a chemical or a physical property.
a. Color
Physical
b. Reaction with HCl
Chemical
c. Boiling point
Physical
d. Density
Physical
e. Hardness
Physical
f. Flammablity
Chemical
Compare and contrast chemical and physical changes. List signs of chemical changes.
 A chemical change results in a new, different substance, while a physical change
does not. Chemical changes are shown with bubbling, color change, precipitate
formation, temperature change and a substance “disappearing.”
What is the law of conservation of mass? Give an example to explain it.
 The law of conservation of mass states that matter is neither created nor
destroyed in a chemical reaction. If liquid water is broken into H2 and O2 gases,
the sum of the gases’ mass equals the water’s mass.
What did each of these scientists contribute to chemistry:
a. Dalton
c. Mendeleev
 Dalton’s Atomic Theory.
 Modern periodic table.
Partial Pressure. Law of
Predicted properties of
Multiple Proportions.
undiscovered elements.
b. Rutherford
d. Thompson
 Discovered the nucleus of
 Discovered electron in the
the atom. Nuclear Atomic
cathode ray tube
Model.
experiment. Plum Pudding
Model.
Final Review 3
16.
What happens to an atom when it becomes an anion? A cation?

17.
What is different about 2 isotopes of titanium? What is the same?

18.
Hydrogen-1
b.
45
c.
183
d.
20.
Isotopes of titanium have same number of protons and atomic number. Different
isotopes have different amounts of neutrons and mass numbers.
Complete the following table (assume atoms are neutral):
Atomic
Mass
Number of
Isotope Name
number
Number
Protons
a.
19.
An anion is the result of gaining an electron to become negatively charged. A
cation loses an electron to be positively charged.
Sc
W
Bromine - 81
Number of
Neutrons
Number of
Electrons
1
1
1
0
1
21
45
21
24
21
74
183
74
109
74
35
81
35
46
35
Element X has two naturally occurring isotopes. Isotope 35X has an abundance of 77.60%
and isotope 37X has an abundance of 22.40 %.
a. What is the atomic mass of element X (to 2 decimal places)?
35
X: 35 amu x 0.7760 = 27.16 amu
mass of X = 35.45 amu
37
X: 37 amu x 0.2240 = 8.29 amu
b. What is the identity of element X?
Element X is Chlorine (Cl)
What is the makeup and symbol for an:`
a. Alpha particle?

An alpha particle (α) is a helium nucleus ( 42He , 2 protons, 2 electrons).
b. Beta particle?

A beta particle (β) is an electron ( -10e).
c. Gamma particle?

21.
Write a balanced equation for Cobalt-60 emitting an alpha particle.
60
27
22.
4
Co→ 56
25 Mn+ 2 He
Write a balanced equation for Uranium-236 emitting a beta particle.
236
92
23.
A gamma particle (γ) is high-energy radiation.
U→ 236
Np+ -10 e
93
What is electromagnetic radiation? What is the speed of all electromagnetic radiation?
 Electromagnetic radiation is energy that travels like a wave through space.
All radiation moves at the speed of light (3 x 108 m/s).
Final Review 4
24.
What frequency of light will have a wavelength of 3.12 x 10-2 m?
v=
25.
c
3 x 108 m/s
=
= 9.6 x 109 Hz
λ 3.12 x 10-2 m
What is the energy of a wave with a frequency of 1.76 x 1016 Hz?
E = hv = (6.626 x 10-34 J·s)(1.76 x 1016 Hz) = 1.16 x 10-17 J
Explain in detail how an atom emits light and why it is a different color for each element.

What do quantum numbers tell you about an electron?

30.
a. Al
1s2
2s2
2p6
b. N
1s2
2s2
2p3
c. K
1s2
2s2
2p6
3s2
3p1
3s2
3p6
Write the noble gas shorthand for the following:
a. Mg
[Ne] 3s2
c. Ta
[Xe] 6s2 4f14 5d3
b. Fe
[Ar] 4s2 3d6
d. Br
[Ar] 4s2 3d10 4p5
Determine the number of valence electrons and draw the Lewis dot diagram for:
a. P – 5 valence ec. Ca – 2 valence e-
∙
∙
∙∙
P
∙
b. O – 6 valence e∙∙
∙
∙
O
31.
∙∙
Ca
d. I – 7 valence e-
∙∙
I
∙∙
When nitrogen becomes an ion, its electron configuration is the same as what element?

32.
4s1
∙
29.
What is the COMPLETE electron configuration with orbital diagram for the following:
∙∙
28.
Quantum numbers tell you the most likely location of an electron in an atom.
∙
27.
An excited atom moves up to a higher energy level. On the way down, it releases
the extra energy as light. Each element has its own electron configuration and its
own color released from the electrons.
∙
26.
Nitrogen gains 3 electrons to be like Neon.
Use the periodic table to name the following:
a. Group 1
Alkali Metals
b. Group 2 Alkaline Earth Metals
c. Group 7
Halogens
d. Group 8
Noble Gases
Final Review 5
33.
What is the periodic law?

34.
35.
The periodic law states that properties repeat periodically when elements are
listed in order of increasing atomic number.
Write the symbol (with charge) for the most common ions formed by:
a. Boron
B3+
c. Nitrogen
N3-
b. Hydrogen
H+
d. Oxygen
O2-
Which has the smaller atomic radius?
a. Hg
36.
Bi
b. F
or
O
or
Ga
b. Al
or
Tl
Metals are shiny, ductile, generally solid (except Hg), and cations. Nonmetals are
brittle, not ductile, usually liquid or gas and anions.
Ionic bonds donate electrons between metals and nonmetals. Covalent bonds
share electrons between 2 nonmetals. Metallic bonds have metal cations in a sea
of electrons.
A single bond shares 2 e-, double shares 4 e- and triple shares 6 e-.
Compare ionic and molecular compounds in terms of melting point and conductivity.

42.
or
How many electrons are shared in a single covalent bond? A double bond? A triple bond?

41.
Se
Compare and contrast ionic bonds and covalent bonds.

40.
or
What are the properties of a metal? A nonmetal?

39.
b. Ca
Which has the higher electronegativity?
a. Br
38.
Cd
Which has the higher ionization energy?
a. N
37.
or
Ionic compounds have a higher melting point and conductivity than molecular
compounds.
What type(s) of bonds are in the following:
a. Ca3P2
b. Sn(OH)2
c. H2O
d. BaSO4
Ionic
Ionic and Covalent
Covalent
Ionic and Covalent
e. CsF
Ionic
f. P2O5
Covalent
g. CdO
Ionic
h. AgNO3 Ionic and Covalent
Final Review 6
43.
Write the name or formula for the following:
a. Sulfur Difluoride
SF2
b. Pb(C2H3O2)2
Lead (II) Acetate
c. Potassium Iodide
KI
d. CO2
Carbon Dioxide
e. Fe2O3
Iron (III) Oxide
f. Cobalt (II) Carbonate
CoCO3
g. Tetraphosphorus Pentoxide
P4O5
h. BaBr2
44.
What is the basis of the VSEPR theory?

45.
Formula
Barium Bromide
The VSEPR theory is based on electron pairs in atoms repelling each other as
much as possible.
Complete the following table:
Lewis Structure
Molecular
Shape
Molecule Polarity
Strongest
Intermolecular
Force
a
H2O
Bent
Polar
Hydrogen
Bonds
b
PCl3
Trigonal
Pyramidal
Polar
Dipole-Dipole
Forces
c
CF4
Tetrahedral
Nonpolar
London
Dispersion
Forces
Final Review 7
46.
Name the seven diatomic elements.

Bromine (Br2), Iodine (I2), Nitrogen (N2), Chlorine (Cl2), Hydrogen (H2), Oxygen
(O2), and Fluorine (F2).
47.
An exothermic reaction releases energy as a product.
48.
An endothermic reaction takes in energy as a reactant.
49.
Write the general formula and explain what occurs in the reaction types below:
A + B  AB
a. Synthesis

2 substances combine to make 1 compound.
AB  A+B
b. Decomposition

1 compound breaks down into 2 or more simpler substances.
A + BX  AX + B or
c. Single Replacement

A free metal replaces another metal in a compound or a halogen replaces
a halogen in a compound.
AX + BY  AY + BX
d. Double Replacement

2 compounds switch anions.
CxHy + O2  CO2 + H2O
e. Combustion

50.
A substance (usually a hydrocarbon) burns in oxygen to make carbon
dioxide and water.
Balance and classify the following equations:
a. 4 Fe + 3 O2 
2 Fe2O3
Synthesis
b. 2 Al(NO3)3 + 3 H2SO4 
c. 2 PbO2 
51.
Y + AX  AY + X
2 PbO +
Al2(SO4)3 + 6 HNO3
O2
Double Replacement
Decomposition
d. 2 KBr + Cl2  2 KCl + Br2
Single Replacement
e. CH4 + 2 O2  CO2 + 2 H2O
Combustion
For the reactions below, predict the products, balance, and classify the type:
a. HCl +
NaOH  NaCl + H2O
b. 2 Al + 3 Cl2 
2 AlCl3
Synthesis
c. 2 H2O  2 H2 + O2
d. 2 C2H6 + 7 O2  4 CO2 +
Double Replacement
Decomposition
6 H2O
Combustion
Final Review 8
52.
53.
54.
55.
56.
57.
What do these reaction symbols represent?
a. (g)
Gas
d. 
Reversible Reaction
e. (aq)
Aqueous Solution
b. 
Yields or Reacts to Form
f. (s)
Solid
c. (l)
Liquid
What is the molar mass of CO2?
C: 1 x 12.01 g/mol = 12.01 g/mol, O: 2 x 16.00 g/mol = 32.00 g/mol MM=44.01 g/mole
What is the mass in grams of 8.44 moles of the element tellurium, Te? How many atoms
is that?
127.60 g Te
8.44 mol Te x
= 1080 g
1 mol Te
6.02 x 1023 atoms Te
8.44 mol Te x
= 5.08 x1024 atoms Te
1 mol Te
At STP, 6.81 moles of N2 gas will take up what volume?
22.4 L N2
6.81 mol N2 x
= 153 L N𝟐
1 mol N2
How many moles of Cu(OH)2 are there in 83.5 g?
1 mole Cu(OH)2
83.5 g Cu(OH)2 x
= 0.856 mol Cu(OH)2
97.57 g Cu(OH)2
What is the percent composition (rounded to the hundredths place) of:
a. Mg(NO3)2 – Mass = 148.33 amu
24.31 amu
Mg: 1 x 24.31 amu = 148.33 amu x100 = 16.39 % Mg
N: 2 x 14.01 amu =
28.02 amu
x100 = 18.89 % N
148.33 amu
96.00 amu
O: 6 x 16.00 amu = 148.33 amu x100 = 64.72 % O
b. Al2(SO4)3 – Mass = 342.17 amu
53.96 amu
Al: 2 x 26.98 amu = 342.17 amu x100 = 15.77 % Al
S: 3 x 32.07 amu =
96.21 amu
x100 = 28.12 % S
342.17 amu
192.00 amu
O: 12 x 16.00 amu = 342.17 amu x100 = 56.11 % O
c. H2O – Mass = 18.02 amu
2.02 amu
H: 2 x 1.01 amu = 18.02 amu x100 = 11.21 % H
16.00 amu
58.
59.
Cl: 1 x 16.00 amu = 18.02 amu x100 = 88.79 % O
Write the empirical formula for the following:
a. FeS
FeS
b. Si2O4
A compound is composed of 40.00 % C, 6.72 % H and 53.29 % O.
a. Find its empirical formula.
1 mol C
C: 40.00 g C x 12.01 g C = 3.33 mol C / 3.33 mol = 1
H:
1 mol H
6.72 g H x 1.01 g H = 6.65 mol H / 3.33 mol = 2
SiO2
Empirical Formula – CH2O
1 mol O
O: 53.29 g O x 16.00 g O = 3.33 mol O / 3.33 mol = 1
b. If the molar mass of the compound is 60.06 g/mol, what is the molecular formula?
Molecular Formula
60.06 g/mol
CH2O molar mass=30.03 g/mol, Empirical Formula = 30.03 g/mol = 2, Molecular Formula = C2H4O2
Final Review 9
60.
61.
Write the name or formula for the following hydrates:
Calcium Chloride Hexahydrate
b. Barium hydroxide octahydrate
Ba(OH)2 ∙ 8 H2O
What is a mole ratio? How can it be determined?

62.
a. CaCl2 ∙ 6H2O
The mole ratio is the ratio between substances in a reaction. It is determined
from the coefficients of the balanced reaction.
Using the equation below:
C2H4 (g) + 3 O2 (g) 2 CO2 (g) + 2 H2O (g)
63.
a. How many moles of carbon dioxide are made with 6.33 moles of water?
2 mol CO2
6.33 mol H2O x
= 6.33 mol H2O
2 mol H2O
b. How many grams of water are made from 2.12 moles of ethylene (C2H4)?
2 mol H2O 18.02 g H2O
2.12 mol C2H4 x
x
= 76.4 g H2O
1 mol C2H4
1 mol H2O
c. How many liters of carbón dioxide can be made with 66.2 L of oxygen?
2 L CO2
66.2 L O2 x
= 44.1 L CO2
3 L O2
d. How many grams of oxygen are needed to produce 33.1 grams of ethylene?
1 mol C2H4
3 mol O2
32.00 g O2
33.1 g C2H4 x
x
x
= 113 g O2
28.06 g C2H4
1 mol C2H4
1 mol O2
You collect 582 g of a product in a reaction. You calculated the theoretical yield to be
733 g. What is the percent yield in your reaction?
actual
582 g
x 100 =
x 100 = 79.40 % yield
theoretical
733 g
What state(s) of matter:
Percent Yield =
64.
a. Is most compressible?
b. Has the highest density?
65.
d. Are fluids?
Solid
Gas and Liquid
London dispersions are temporary dipole attractions between nonpolar
substances and the weakest forces, dipole-dipole interactions are between polar
substances, and hydrogen bonds are between polar compounds with N, O, or F
and the strongest forces.
What type of energy changes during a phase change?

67.
Solid
c. Has the lowest energy?
Compare and contrast these intermolecular forces: London dispersion, dipole-dipole, and
hydrogen bonds.

66.
Gas
During a phase change, the potential energy is changing, kinetic stays the same.
What is STP? What are the conditions of STP?
 STP is standard temperature (0 °C) and pressure (1 atm).
Final Review 10
68.
What state change occurs when:
a. A gas becomes a solid?
b. A liquid becomes a gas?
Deposition
Boiling/Evaporation/Vaporization
69.
Answer the following questions based on the phase diagram below:
70.
a. Name points d & e on the phase diagram.
 d – triple point (all 3 states exist).
 e – critical point (liquid and gas indistinguishable – supercritical fluid).
b. What state change occurs as you heat a substance from C to B on the phase
diagram?
 Heating from C to B causes Boiling/Evaporation/Vaporization.
c. What is the normal melting point of the substance above?
 The normal melting point is 60 °C for the substance above.
Using the heating curve to the right:
71.
72.
a. What state of matter is being heated at
leg c?
 At leg c, the liquid is being
heated.
b. On which line is the substance melting?
 The substance is melting during
leg b.
What are the assumptions of the kinetic
molecular theory?
 The kinetic molecular theory assumes
that there are no interactions between
particles, no volume from the particles, there are perfectly elastic collisions, and
the temperature is proportional to kinetic energy.
Perform the following pressure conversions:
a. 685 mm Hg to atm
1 atm
685 mm Hg x
= 0.901 atm
760 mm Hg
b. 112 kPa to torr
760 torr
112 kPa x
= 840. torr
101.3 kPa
Final Review 11
73.
If a gas at a pressure of 1.25 atm is compressed from 311 mL to 154 mL, what is the new
pressure? Which law is this?
P1V1 = P2V2; P2 =
74.
P1 V1 (1.25 atm)(311 mL)
=
= 2.52 atm , Boyle's Law
V2
0154 mL
A tire at a temperature of 25 °C increases in pressure from 740 mmHg to 1125 mmHg.
What is the new temperature? Which law is this?
P1 P2
P2 T1 (1125 mmHg)(298 K)
=
; T2 =
=
= 453 K , Gay-Lussac's Law.
T1 T2
P1
740 mmHg
75.
A gas takes up a volume of 17.3 liters, has a pressure of 2.30 atm, and a temperature of
299 K. If I raise the temperature to 350. K and lower the pressure to 1.5 atm, what is the
new volume of the gas? Which law is this?
P1 V1
P2 V2 (2.30 atm)(17.3 L)
(1.5 atm)V2
=
;
=
; V2 = 31.1 L , Combined Gas Law.
T1
T2
299 K
350. K
76.
If I have an unknown quantity of gas at a pressure of 1.20 atm, a volume of 31.0 liters,
and a temperature of 87 °C, how many moles of gas do I have? Which law is this?
PV
(1.20 atm)(31.0 L)
=
= 1.26 mol , Ideal Gas Law
RT (0.0821 L ∙ atm )(360. K)
mol ∙ K
Which will diffuse faster: He or N2? Why?
PV = nRT ; n =
77.

78.
What temperature is absolute zero? What happens at absolute zero?

79.
80.
He will diffuse faster because it is lighter (smaller molar mass) than N2.
Absolute zero is 0 K. At absolute zero, all matter has no kinetic energy and
freezes.
The substance being dissolved in a solution is the solute and the substance doing the
dissolving is the solvent.
What is the molarity of a 0.875 L solution with 21.2 g of lithium bromide?
21.2 g LiBr x
81.
1 mol LiBr
mol 0.244 mol
= 0.244 mol LiBr ; M =
=
= 0.279 M
86.84 g LiBr
V
0.875 L
How many moles of sodium chloride are needed to make 250 mL of a 2.5 M solution?
mol = MV = (2.5 M)(0.25 L) = 0.625 mol
82.
At 50°C, a solution contains 20 g of KCl in 100 g of water.
According to the solubility curve, is the solution saturated,
unsaturated, or supersaturated?

83.
The solution would be unsaturated.
According to the solubility curve, how much more KClO3
can be added if the solution is heated from 30°C to 70°C?

30 g – 10 g = 20 g
Final Review 12
84.
How much heat is needed to raise the temperature of 29.5 g of iron from 295 K to 365 K
if the specific heat of iron is 0.450 J/g∙K?
Q = m C ΔT = (29.5 g)(0.450 J/g∙K)(365 K – 295 K) = 929 J
85.
Describe the properties of acids and bases below:
Acid
a. Taste
Sour
Bitter
b. Ion released in water
H+
OH-
c. Reactivity with Metal
Reactive
Unreactive
d. pH level
86.
Base
<7
>7
A solution has a pH of 3.
a. What is the [H+]?
[H+] = 1 x 10-3 M
b. What is the pOH?
pOH = 14 – pH = 14 – 3 = 11
c. What is the [OH-]?
[OH-] =
Kw
1 x 10-14
-11
M
+ =
-3 = 1 x 10
[H ]
1 x 10
d. Is the solution more likely to be acetic acid or potassium hydroxide?

87.
The solution is most likely acetic acid.
Write the name or formula for the following acids:
a. HNO3
b. Hydrofluoric acid
c. H2S
d. Sulfurous Acid
Nitric Acid
HF
Hydrosulfuric Acid
H2SO3
Look over your notes, assignment packets and these review questions.
Good luck & Study! Study!! Study!!!