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AP Unit 4 Review

Total questions: 171

Worksheet time: 29hrs 30mins

Name
Class
Date
1.

Which of the following is an intramolecular force?

a)

hydrogen bonding

b)

polar covalent bonding

c)

London dispersion forces

d)

dipole-dipole interaction

2.

On a potential energy graph, what is graphed on the x-axis?

a)

internuclear distance

b)

potential energy

c)

IMF distances

d)

Coulombic forces

3.

What is the relationship between potential energy and bond energy?

a)

low PE, high bond energy

b)

low PE, low bond energy

c)

high PE, high bond energy

d)

there is no relation

4.

Based on the lattice energy for the decomposition of NaF, how would the lattice energy for MgO compare?

a)

greater than because the ions in MgO have a greater charge

b)

less than because Mg is a smaller ion than Na

c)

equal to because the difference in ion sizes is similar between both compounds

d)

less than because the molar mass of MgO is less than the molar mass of NaF

5.

The graph shows the amount of potential energy between two hydrogen atoms as the distance between them changes. At which point in the graph would a molecule of H2 be the most stable?

a)

A

b)

B

c)

C

d)

D

6.

What is the correct Lewis Dot Structure for ammonia NH3

a)
b)
c)
d)
7.

Which is the correct molecular structure for carbon dioxide?

a)
b)
c)
d)
8.
Which of the following is the correct Lewis dot structure for the molecule fluorine (F2)?
a)
A
b)
B
c)
C
d)
D
9.
A molecule of methane has what shape?
a)
Tetrahedral
b)
Triangular
c)
Pyramidal
d)
Octahedral
10.

Type of bond between Br & Br

a)

Ionic

b)

Polar Covalent

c)

Nonpolar Covalent

11.
What molecule could this be? 
a)
BF3
b)
CH4
c)
H2O
d)
CO2
12.
Calculate the Formal Charge for the Oxygen Labeled 1 
a)
-1
b)
+1
c)
0
d)
-2
13.
Calculate the Formal Charge for the Nitrogen
a)
-1
b)
+1
c)
0
d)
-2
14.
What is the formal charge for Xenon
a)
-1
b)
+1
c)
0
d)
+2
15.
How many resonance structures for NO3- ion?
a)
1
b)
2
c)
3
d)
4
16.
What is the bond order for NO3- ion?
a)
1
b)
1.5
c)
2
d)
1.3
17.

What is the bond order between S and O?

a)

1

b)

1.5

c)

2

d)

3

18.
What is the VSEPR theory used to predict?
a)
Bond Strength
b)
Polarity
c)
Molecular Shape
d)
Electronegativity
19.

What is the the shape of this molecule according to VSEPR theory?

a)

Linear

b)

Tetrahedral

c)

Trigonal Planar

d)

Trigonal pyramidal

20.

What could this be?

a)

CO2

b)

NH3

c)

H2S

d)

CH4

21.

This carbon atom went through ....

a)

sp hybridization

b)

sp2 hybridization

c)

sp3 hybridization

d)

dsp3 hybridization

22.

Which of the following images represents a Trigonal Planar molecular geometry?

a)
b)
c)
d)
23.

What is the hybridization of this molecule shown above

a)

sp

b)

sp2

c)

sp3

d)

sp4

24.

What is the bond angle for the CH4 molecule?

a)

120°

b)

180°

c)

109.5°

d)

90°

25.

Two carbon atoms form a linear shaped ethene, C2H4, molecule when bonded with four other hydrogen atoms. Prior to bonding, the carbon atom

a)

does not hybridize

b)

undergoes sp hybridization

c)

undergoes sp2 hybridization

d)

undergoes sp3 hybridization

26.
How many sigma bonds does this have? 
a)
1
b)
2
c)
3
d)
4
27.
How many pi bonds does this have? 
a)
1
b)
2
c)
3
d)
4
28.
How many sigma and pi bonds does this have?
a)
1 sigma and 1 pi
b)
2 sigma and 1 pi
c)
1 sigma and 2 pi
d)
2 sigma and 2 pi
29.

Compounds that are composed of the same number and type of atoms but have them arranged in different ways are ___________.

a)

isomers

b)

isotopes

c)

polymers

d)

alkanes

30.

In a stereoisomer, ________ means "on the same side" and ______ means "across from."

a)

cis; trans

b)

cis; L

c)

D; L

d)

trans; cis

31.

Given the formula representing a compound:


Which formula represents an isomer of this compound?

a)
b)
c)
d)
32.
In a polar bond, the more electronegative element will assume a partial ________ charge.
a)
positive
b)
negative
33.
What is the shape of H2O?
a)
linear
b)
tetrahedral
c)
bent
d)
trigonal pyramidal
34.
What is the first thing you need to do when drawing Lewis Structures?
a)
Draw your bonds
b)
Count the number of valence electrons
c)
Subtract to make the valence number even
d)
Add to make the valence number even
35.
After drawing in your bonds, what do you do if you don't have enough electrons to get each atom to its octet?
a)
Place dots until everything has eight
b)
Place dots only around the terminal atoms
c)
Add a multiple bond
d)
Have eight only around the central atom
36.
In a Lewis structure, what do you do  with the total number of valence if your ion is 2- charged?
a)
Add 2 to the valence number
b)
Subtract 2 from the valence number
c)
Multiply the valence number by 2
d)
Divide the valence number by 2
37.

Which of the following is a valid Lewis structure for hydrogen cyanide, HCN?

a)
b)
c)
d)
38.

Which of the following forces are found between two NONPOLAR molecules?

a)

London dispersion force

b)

Dipole-dipole

c)

Hydrogen bond

39.

What intermolecular force is present between two molecules of HCl?

a)

Dipole-Dipole

b)

London

c)

Hydrogen bonds

d)

Ionic

40.

What type of intermolecular force is present between two molecules of HF?

a)

Dipole-Dipole

b)

London

c)

Hydrogen bonds

d)

ionic

41.

Rank these in order of strength (from strongest to weakest):

London forces

hydrogen bond

dipole-dipole attraction

a)

dipole-dipole>hydrogen bond>London

b)

London>dipole-dipole>hydrogen bond

c)

hydrogen bond>dipole-dipole>London

d)

London>hydrogen bond>dipole-dipole

42.
In general, substances with stronger intermolecular forces have ________ boiling points than those with weaker forces
a)
Higher
b)
Lower
c)
Forces and boiling point are not related
43.
In this Lewis structure, the symbol above F means...
a)
electrons are being transferred to Fluorine
b)
electrons are less attracted to F than H
c)
electrons are more attracted to F than H
d)
Fluorine has formed a cation
44.

A diatomic molecule like O2 is always_______ because electrons are shared ________.

a)
nonpolar; equally
b)
polar; equally
c)
nonpolar; unequally
d)
nonpolar; unequally
45.

Is the following molecule polar or nonpolar?

a)

polar

b)

nonpolar

46.

Is the following molecule polar or nonpolar?

a)

polar

b)

nonpolar

47.

Is the following molecule polar or nonpolar?

a)

polar

b)

nonpolar

48.

Is the following molecule polar or nonpolar?

a)

polar

b)

nonpolar

49.
What explains the very high melting and boiling point of water
a)
Strong dipole-dipole bonds between water molecules
b)
Strong hydrogen bonds between water molecules
c)
Dispersion forces which are present in all molecules
d)
Asymmetrical shape of the polar bonds.
50.
Liquids with weak intermolecular forces
a)
contain a lot of kinetic energy
b)
easily evaporate
c)
cannot diffuse
d)
freeze easily
51.

Which model represents a LOWER Vapor Pressure?

a)

A

b)

B

52.

Which model represents STRONGER Intermolecular Forces?

a)

A

b)

B

53.

Intermolecular forces for: NH3

a)

Dispersion Force

b)

Dipole dipole

c)

Hydrogen bonding

54.

The amount of heat needed to change 1 gram of a substance from the solid phase to the liquid phase is ___

a)

Heat of Vaporization

b)

Boiling point

c)

Heat Energy

d)

Heat of fusion

55.

Consider the following temperature-time graph when a solid is heated. which part of the graph involves heat of vaporization?

a)

Part I

b)

Part II

c)

Part III

d)

Part IV

56.

A substance has a melting point of 25 oC and a boiling point of 235 oC. What state is it at 5oC?

a)

Solid

b)

Liquid

c)

Gas

57.
When the temperature of matter increases the particles...
a)
speed up and move closer
b)
speed up and move farther apart
c)
slow down and move closer together
d)
slow down and move farther apart
58.

Which type of IMF is responsible for the attraction pictured above?

a)

Dipole-Dipole Interaction

b)

Ion-Dipole Interaction

c)

Hydrogen Bonds

d)

Covalent Bond

e)

Ionic Bond

59.

Viscosity is

a)

the resistance to flow

b)

the resistance to spread out

c)

the ability to make sheets of metal

d)

the ability to make wire

60.

A weak IMF will result in a

a)

high melting point

b)

high viscosity

c)

high surface tension

d)

low surface tension

61.

When a covalent solid such as dry ice changes phase, what bonds are broken?

a)

covalent bonds between the atoms

b)

intramolecular forces in the molecule

c)

intermolecular forces between the molecules

d)

bonds both within and between the molecules

62.

Vapor pressure increases with increasing ______.

a)

temperature

b)

pressure

c)

molecular weight

d)

size of the container

63.

As the IMF strength increases, ____ decreases.

a)

boiling point

b)

melting point

c)

surface tension

d)

vapor pressure

e)

viscosity

64.

As the temperature of a liquid increases, the viscosity _____.

a)

increases

b)

decreases

c)

is unaffected

65.

Sugar is added to tea to make sweet tea. In chemical terms the sugar is the ___________ and the tea (mostly water) is the _____.

a)

endothermic, exothermic.

b)

chemical, dissolution.

c)

reactant, product.

d)

solute, solvent.

66.

The seam between atoms in a crystalline solid is called a ___

a)

division line

b)

fracture point

c)

discontinuity boundary

d)

cleavage plane

67.
Molecular solids are made from ______. 
a)
metals 
b)

metals + nonmetals

c)
sodium
d)
non-metals 
68.
Ionic bonds _____ dissolve in polar solvents. 
a)
will
b)
will not 
c)
might 
d)
I don't know
69.

Which term means "can be hammered or shaped?"

a)
malleable 
b)
ductile 
c)
hard 
d)
soft
70.
Cgraphite
a)
metallic
b)
atomic/molecular
c)

network covalent

d)
ionic
71.
Cdiamond
a)
metallic
b)
atomic/molecular
c)

network covalent

d)
ionic
72.
This type of solid will conduct electricity when dissolved.
a)
ionic
b)
molecular
c)
network covalent
d)
metallic
73.
This type of solid will conduct electricity when a solid.
a)
ionic
b)
molecular 
c)
metallic
d)
network covalent
74.
This type of solid will have the lowest boiling point.
a)
ionic
b)
molecular
c)
network
d)
metallic
75.
What type of solid is both ductile and malleable?
a)
ionic
b)
metallic
c)
network
d)
molecular
76.
What type of solid is described as hard but brittle.
a)
ionic
b)
metallic
c)
network
d)
covalent
77.

This type of network covalent solid is soft and can conduct electricity because it is 2D and has free electrons.

a)
diamond
b)
sand
c)
graphite
78.

This type of solid forms a regular repeating

three-dimensional structure called a crystal lattice. What type of solid is this?

a)

amorphous

b)

crystalline

c)

glass

d)

diamond

79.

In amorphous solids, the atoms or molecules are held together in a completely random formation.

a)

true

b)

false

80.

One major difference between crystalline and amorphous solids is that

a)

crystalline solids have a precise melting point.

b)

amorphous solids have a lattice structure.

c)

amorphous solids always behave consistently and uniformly.

d)

crystalline solids break unpredictably and can produce curved fragments.

81.

A sample of a hard, solid binary compound at room temperature did not conduct electricity as a pure solid but became highly conductive when dissolved in water. Which of the following types of interactions is most likely found between the particles in the substance?

a)

Ionic bonds

b)

Metallic bonds

c)

Covalent bonds

d)

Hydrogen bonds

82.

A certain crystalline substance that has a low melting point does not conduct electricity in solution or when melted. This substance is likely to be

a)

a network covalent solid

b)

a metallic solid

c)

an ionic solid

d)

a molecular solid

83.

A student is given a sample of a pure, white crystalline substance. Which of the following would be most useful in providing data to determine if the substance is an ionic compound?

a)

Examining the crystals of the substance under a microscope

b)

Determining the density of the substance

c)

Testing the electrical conductivity of the crystals

d)

Testing the electrical conductivity of an aqueous solution of the substance

84.

Which of the following most likely describes the solid represented in the diagram above?

a)

It is soft and is a poor thermal and electrical conductor.

b)

It has a low melting point and is a poor thermal and electrical conductor as a solid.

c)

It is a brittle, water-soluble electrolyte that is a poor thermal and electrical conductor as a solid.

d)

It is malleable and is an excellent thermal and electrical conductor as a solid.

85.

Three substances were studied in the laboratory, and the data in the table above were collected. Based on the data, which of the following shows the type of bonding in each substance?

a)

Substance X: Network covalent; Substance Y: Ionic; Substance Z: Metallic

b)

Substance X: Ionic; Substance Y: Molecular; Substance Z: Network covalent

c)

Substance X: Molecular; Substance Y: Network covalent; Substance Z: Metallic

d)

Substance X: Network covalent; Substance Y: Metallic; Substance Z: Ionic

86.

Which type of solid has very strong covalent bonds between neighboring atoms?

a)

Ionic Crystal

b)

Covalent Molecular Crystal

c)

Metallic Crystal

d)

Network Covalent Crystal

87.

Very hard, very high melting point, nonconductors of electricity in solution

a)

Ionic

b)

Metallic

c)

Network Covalent

d)

Covalent Molecular

88.

Identify the property that gives metallic solids the ability to conduct electricity well, but not network covalent solids.

a)

Electrons in metallic solids are strongly bound in place

b)

Electrons in network covalent solids are freely moving around

c)

Network covalent solids are not bound to other atoms

d)

Metallic solids have a sea of electrons that are not bound and can move about the solid freely.

89.

Metallic solids are composed of metal cations held together by a delocalized "sea" of valence electrons.

a)

True

b)

False

90.

Which of the following are examples (3) of covalent network solids?

a)

diamond

b)

silicon dioxide

c)

water

d)

graphite

e)

carbon tetrachloride

91.
Molar volume of any gas at STP
a)
22.4 L
b)
2.24 L
c)
6.02 x 1023 L
d)
224 L
92.

Standard temperature and pressure (STP) mean

a)

A temperature of 0°C and a pressure of 1 atmosphere (atm).

b)

A temperature of 100°C and a pressure of 1 atmosphere (atm).

c)

A temperature of 0°K and a pressure of 1 atmosphere (atm).

93.
PV=nRT
a)
Charles Law
b)
Boyle's Law
c)
Combined Gas Law
d)
Ideal Gas Law
94.
Calculate the volume that a 0.323-mol sample of a gas will occupy at 265 K and a pressure of 0.900 atm.
a)
7.18 L 
b)
7.81 L
c)
4.63 L
d)
4.36 L
95.
Determine the Kelvin temperature required for 0.0470 mol of gas to fill a balloon to 1.20 L under .998 atm pressure. 
a)
0 K 
b)
107 K 
c)
207 K 
d)
307 K 
96.

What does R stand for

a)

Ideal gas constant

b)

Real gas constant

c)

Temperature constant

d)

Ideal gas law

97.
Dalton's Law states...
a)
that at a constant temperature the volume of a confined ideal gas varies inversely with its pressure.
b)
that the total pressure exerted by the mixture of non-reactive gases is equal to the sum of the partial pressures of individual gases
c)
how gases tend to expand when heated
d)
that the density of an ideal gas at constant pressure varies inversely with the absolute temperature of the gas.
98.
Dalton's law of...
a)
Partial Pressure
b)
Thermodynamics
c)
Ideal Gas
d)
Physics
99.

Three gases, Ar, N2 and H2 are mixed in a 500 L container. Ar has a pressure of 255 torr, N2 has a pressure of 228 torr, and H2 has pressure of 752 torr. What is the total pressure in the container?

a)

483 torr

b)

270 torr

c)

1235 torr

100.

A 250 mL sample of oxygen is collected over water at 25° C and 760.0 torr pressure. What is the pressure of the gas alone? (Vapor pressure of water at 25° C = 23.8 torr)

a)

783.8 torr

b)

205.38 torr

c)

1250 torr

d)

736.2 torr

101.

Total number of the mole fractions for all component in the solution is equal to

a)
1.0
b)
0.95
c)
0.50
d)
2.0
102.

According the KMT, the particles are in

a)

different states

b)

organized patterns

c)

constant motion

d)

various containers

103.

What does increasing the temperature of a solid lead to?

a)

increased mass of particles

b)

decreased space between particles

c)

decreased velocity of particles

d)

increased velocity of particles

104.

How do the particles in a gas behave?

a)

Neatly organized without any movement.

b)

Move around trapped in a small volume and with the same surrounding particles.

c)

Move fast enough to break attractions and pushing other particles out of the way.

d)

Move freely in a straight line until they collide with another particle.

105.

If the particles are moving fast enough to complete break the attraction pulling it back towards other particles, what process is occuring?

a)

sublimation

b)

melting

c)

freezing

d)

evaporation

106.

When particles in a gas collide, they change

a)

direction and velocity

b)

velocity and temperature

c)

direction and volume

d)

volume and temperature

107.
Gas pressure is caused by
a)
gas molecules colliding with surfaces
b)
gas molecules condensing to a liquid
c)
gas molecules hitting other gas molecule
d)
barometers
108.
a)
Flask 1 contain the most molecules
b)
Flask 2 contain the most molecules
c)
Flask 3 contain the most molecules
d)
Flask 4 contain the most molecules
109.
a)
Since all flask have same amount of gas, flask 1 has the least pressure
b)
Since all flask have same amount of gas, flask 2 has the least pressure
c)
Since all flask have same amount of gas, flask 3 has the least pressure
d)
Since all flask have same amount of gas, flask 4 has the least pressure
110.
a)
Gas particles are elastic and do not attract each other.
b)
The kinetic energy of gas is dependent on temperature 
c)
Energy is not lost when gas particles collide with each other or with the walls of the container.
d)
All of the above
111.

If He atoms (mass 4), Ne atoms (10), Ar atoms (40) and Kr atoms (84) at the same temperature, what is true about their average kinetic energy?

a)

He has the highest kinetic energy

b)

Kr has the highest kinetic energy

c)

Ar has the highest kinetic energy

d)

They all have the same kinetic energy

112.

At a given temperature , all gases have the same

a)

Volume

b)

Mass

c)

Average Kinetic Energy

d)

Velocity

113.

According to the kinetic molecular theory, which statement describes the particles in a sample of an ideal gas?

a)

The particles have attractive forces between them.

b)

The particles are considered to have negligible volume.

c)

The particles are constantly moving in circular paths.

d)

The particles collide, decreasing the total energy of the system.

114.

Which phrase describes the motion and attractive forces of ideal gas particles?

a)

random curved-line motion and no attractive forces

b)

random straight-line motion and strong attractive forces

c)

random straight-line motion and no attractive forces

d)

random curved-line motion and strong attractive forces

115.

A real gas behaves most like an ideal gas at

a)

average potential energy of its particles

b)

ionization energy of its particles

c)

low pressure and high temperature

d)

activation energy of its particles

116.

The concept of an ideal gas is used to explain

a)

the behavior of a gas sample

b)

why some gases are monatomic

c)

why some gases are diatomic

d)

the mass of a gas sample

117.

Which graph best represents the pressure-volume relationship for an ideal gas at constant temperature?

a)
b)
c)
d)
118.

Which graph shows the relationship between pressure and Kelvin temperature for an ideal gas at constant volume?

a)
b)
c)
d)
119.

As the temperature of a gas increases with the volume remaining constant, the pressure of the gas

a)

decreases

b)

increases

c)

remains the same

120.

The variable that determines average kinetic energy is the _____.

a)

Temperature

b)

Volume of the container

c)

Mass of the particles

121.

Which of the following is true based on the distribution, assuming all gases are at the same temperature?

a)

Helium has the highest average kinetic energy

b)

Ne has a greater average velocity than Xe

c)

warmer gases have less variability in terms of speed

d)

larger particles move faster than smaller particles on average

122.

The diagram shows the Maxwell−Boltzmann distribution of kinetic energy in a gas at two different temperatures.

Which letter represents the most probable energy of the molecules at the higher temperature? (which axis is kinetic energy or speed graphed on)

a)

A

b)

B

c)

C

d)

D

123.

This question is about the Maxwell–Boltzmann distribution of molecular energies in a sample of a gas shown in the figure below.


What does the area under the curve represent?

a)

The total energy of the particles.

b)

The total number of particles.

c)

The number of particles that can react with each other.

d)

The total number of particles that have activation energy.

124.

For objects that have the same kinetic energy, bigger masses move ___

a)

more slowly

b)

faster

c)

at the same rate

d)

sideways

125.

When gases are allowed to move freely, they tend to move from regions of ___

a)

low concentration to regions of high concentration

b)

high concentration to regions of low concentration

c)

low temperature to regions of high concentration

d)

high concentration to regions of low temperature

126.
Choose the gas that would demonstrate the GREATEST deviation from ideal gas behavior.
a)
He
b)
F2
c)
SF2
d)
CH4
127.

The partial pressures of CH4, N2, and O2 in a sample of gas were found to be 143 mmHg, 469 mmHg, and 563 mmHg, respectively. Calculate the mole fraction of oxygen.

a)

20.1

b)

0.399

c)

0.479

d)

0.741

128.
What element is being Oxidized?
a)
N
b)
H
c)
O
d)
S
129.
What element is being Reduced?
a)
N
b)
H
c)
O
d)
S
130.
What is oxidation number of H in H2O?
a)
0
b)
-2
c)
-1
d)
+1
131.
What is oxidation number of Cr in Cr2O72-?
a)
-2
b)
+2
c)
+6
d)
+12
132.
What is oxidation number of H in CaH2?
a)
+1
b)
-1
c)
0
d)
+2
133.
What is oxidation number of O in O2 ?
a)
0
b)
-2
c)
+1
d)
+2
134.
If in a reaction, copper is reduced; its number of electrons has:
a)
Increased
b)
Decreased
c)
Remained Constant
d)
Varies Randomly
135.
In the reaction
2Ca(s) + O2(g) --> 2CaO(s), calcium is __________
a)
Reduced
b)
Synthesized
c)
Oxidized
d)
None of the above
136.

Defined as the loss of electrons.

a)

oxidation

b)

reduction

c)

redox

d)

OIL RIG

137.

Cl2 + 2e- --> 2Cl - is an example of:

a)

a chemical reaction

b)

redox

c)

oxidation

d)

reduction

138.

Substance that oxidizes another substance by accepting its electrons.

a)

reducing agent

b)

oxidation number

c)

combustion

d)

oxidizing agent

139.

What is a half-reaction of this equation, Na + Cl --> NaCl ?

a)

Cl + e- --> Cl-

b)

Cl - e- --> Cl-

c)

Na --> Na + e-

d)

Na° + Cl°--> Na+1Cl-1

140.
What is the conjugate acid in the following equation?
a)
PO43- 
b)
HNO3 
c)
NO3- 
d)
HPO42-
141.
What is the conjugate base in the following reaction?
a)

HCO3- 

b)

HCl

c)

 H2CO3 

d)

Cl-

142.

Could C2H3O2- act as a base when placed in water?

a)

Yes, because it can accept a H+

b)

Yes, because it can donate a H+

c)

No, because it cannot donate a H+

d)

No, because it cannot accept a H+

143.

What type of reaction is this?

AgNO3(aq) + NaCl(aq) --> AgCl(s) + NaNO3(aq)

a)

Synthesis

b)

Single Displacement

c)

Precipitation

d)

Acid/Base

144.

What type of reaction is this?

HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O(l)

a)

Combustion

b)

Acid/Base

c)

Precipitation

d)

Decomposition

145.

How do you know a REDOX rxn has taken place?

a)

You will see a red glow

b)

You will be happy

c)

Oxidation #'s will change for at least 2 elements

d)

A gas will be produced

146.

Which of the following is an example of a weak acid/strong base reactions?

a)

NaOH + HCl

b)

NH3 + HC2H3O2

c)

HCl + NH3

d)

HC2H3O2 + NaOH

147.

Which of the following represents the net ionic equation for the precipitation reaction between sodium carbonate and magnesium sulfate?

a)

Na2CO3 + MgSO4 --> MgCO3 + Na2SO4

b)

2Na+ + CO3-2 + Mg+2 + SO4-2 --> Mg+2 + CO3-2 + 2Na+ + SO4-2

c)

2Na+ + SO4-2 --> Na2SO4

d)

CO3-2 + Mg+2 --> MgCO3

148.

Which of the following represent a Brønsted-Lowry conjugate acid-base pair?

a)

HF and H2O

b)

NaOH and H2O

c)

HF and NaOH

d)

HF and F-

149.

Given the following: NH3(aq) + HCl(aq) ⇄ NH4+(aq) + Cl-(aq)

The Brønsted-Lowry acids in the reaction represented above are

a)

NH3(aq) and NH4+(aq)

b)

HCl(aq) and NH4+(aq)

c)

HCl(aq) and Cl-(aq)

d)

NH3(aq) and HCl(aq)

150.
Which of the following half reactions correctly represents a reduction half reaction?
a)
Fe →Fe2+ + 2e-
b)
Pb4+ + 2e- →Pb2+
c)
2O-2 →O2 + 4e-
d)
Fe + 3e- →Fe3+
151.

Which of the following represents oxidation?

a)

C → CH4

b)

Fe3+ → Fe2+

c)

Cl2 → 2Cl-

d)

Zn → ZnO

152.

What is the purpose of an "indicator"?

a)

changes color if a chemical reaction occurs

b)

changes color if a physical change occurs

c)

changes color in the presence of an acid or base

d)

changes color when a salt dissolves in water

153.

The table shows the pH of several solutions. Which solution has the greatest concentration (amount) of H+ ions?

a)

vinegar

b)

milk

c)

water

d)

bleach

154.

You have a solution that is a base. You want to lower the pH so you should ___.

a)

add a salt

b)

add an acid

c)

add more base

d)

wait until it evaporates

155.

Molarity is a measure of the ___.

a)

moles of solute dissolved in a solvent.

b)

moles of solvent dissolved in a solute.

c)

Liters of solute dissolved in a solvent.

d)

Liters of solvent dissolved in a solute.

156.
Ca(OH)2 is an example of a(n)
a)
Acid
b)
Base
c)
Neutral
157.
What might happen if you mixed a strong acid with an equally strong base?
a)
You would see an explosive chemical reaction.
b)
The acid would destroy the base.
c)
The base would destroy the acid.
d)
You'd wind up with a pH-neutral substance.
158.
A hydrogen ion, H+, is the same as a(n):
a)
neutron
b)
electron
c)
proton
d)
hydroxide ion
159.
I am titrating 1M HCl with 1M NaOH.  I have 25mL of HCl. How much NaOH will I need?
a)
2.5mL
b)
5mL
c)
25mL
d)
50mL
160.
what is the reading on this burette?
a)

4.4 mL

b)

3.50 mL

c)

3.60 mL

d)

3.6 mL

161.

Why do we perform a titration?

a)

To determine the concentration of an unknown solution

b)

To see if a reaction will occur between an acid and base

c)

To find the mass of an unknown acid

d)

To find the molar mass of an unknown solution

e)

To calculate the viscosity of the solution

162.

Which point on the titration curve corresponds to the point at which the moles of the added strong base are equal to the moles of the weak acid initially present?

a)

Q

b)

R

c)

S

d)

T

163.

A titration works by using a ___________________________ reaction.

a)

Single Replacement

b)

Neutralization

c)

Combustion

d)

Synthesis

164.

In a titration, the solution changes color at the

a)

end point

b)

equivalence point

165.

A titration can be used to determine the _______________ of a solution

a)

molarity

b)

density

c)

boiling point

d)

reactivity

166.

A 10.00 mL sample of H2SO4 was neutralized by 13.5 mL of 1.00 M KOH by the following reaction: H2SO4 +  2KOH  →  K2SO4  +  2H2OH_2SO_4\ +\ \ 2KOH\ \ \rightarrow\ \ K_2SO_4\ \ +\ \ 2H_2O  
What is the molarity of H2SO4?

a)

0.675 M

b)

0.910 M

c)

1.05 M

d)

1.20 M

167.

Which of the following is a strong acid?

a)

HF

b)

KOH

c)

HClO4

d)

HClO

168.

Which of the following is paired incorrectly?

a)

H2SO4 - strong acid

b)

HNO3 - weak acid

c)

Ba(OH)2 - strong base

d)

HCl - strong acid

e)

NH3 - weak base

169.

The following reactions are examples of:

Pb2+ + 2I- --> PbI2

2Ce4+ + 2I- --> I2 + 2Ce3+

HOAc + NH3 --> NH4+ + OAc-

a)

acid-base reactions

b)

precipitation, acid-base, and redox reactions, respectively

c)

redox, acid-base, and precipitation reactions, respectively

d)

precipitation, redox, and acid-base reactions, respectively

170.

The following reactions are examples of:

2K(s) + Br2(l) --> 2KBr(s)

AgNO3(aq) + NaCl(aq) --> AgCl(s) + NaNO3(aq)

HCl(aq) + KOH(aq) --> H2O(l) + KCl(aq)

a)

precipitation reactions

b)

redox, precipitation, and acid-base reactions, respectively

c)

precipitation (two) and acid-base reactions, respectively

d)

redox reactions

171.

ALL of the following reactions can be classified as:

2Al(s) + 3Br2(l) --> 2AlBr3(s)

2Ag2O(s) --> 4Ag(s) + O2(g)

CH4(l) + 2O2(g) --> CO2(g) + 2H2O(g)

a)

oxidation-reduction reactions

b)

combustion reactions

c)

precipitation reactions

d)

oxidation-reduction reactions AND combustion reactions

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