Description

In this collection there will be a bunch of general chemistry Multiple Choice Questions, these questions include Electronic Structure and Periodic Table, Bonding, Phases and Phase Equilibria, Stoichiometry, Thermodynamics and Thermochemistry, Solution Chemistry, Acids and Bases, Electrochemistry

It will be useful for all medical, chemistry and pharmacy students or any student interested in chemistry science

Study Set Content:
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General Chemistry Questions 

 
 
Electronic Structure and Periodic Table 

1. What value or values of 

m

l

 are allowable for an orbital with 

l

 = 2?  

a.

 

0  

b.

 

2  

c.

 

-1  

d.

 

none of the above  

e.

 

all of the above  

2. According to Bohr Theory, which of the following transitions in the hydrogen atom 
will give rise to the 

least energetic

 photon?  

Use the equation: 

E

n

 = (-2.18 x 10

-18

 J)(1/

n

2

)  

a.

 

n

 = 5 to 

n

 = 3  

b.

 

n

 = 6 to 

n

 = 1  

c.

 

n

 = 4 to 

n

 = 3  

d.

 

n

 = 5 to 

n

 = 4  

e.

 

n

 = 6 to 

n

 = 5  

3. Consider a 3

d

xz

 orbital. Which of the following statements is 

incorrect

?  

a.

 

The 

xz

 plane is a nodal surface.  

b.

 

The 

xz

 plane divides the electron probability distribution into two identical 

mirror-image halves.  

c.

 

The 

xy

 plane divides the electron probability distribution into two identical 

mirror-image halves.  

d.

 

The 

yz

 plane divides the electron probability distribution into two identical 

mirror-image halves.  

e.

 

The nucleus is located at a node.  

4. The electronic configuration of the element whose atomic number is 26 is:  

a.

 

1

s

2

 2

s

2

 2

p

6

 3

s

2

 3

p

6

 4

s

0

 3

d

8

  

b.

 

1

s

2

 2

s

2

 2

p

6

 3

s

2

 3

p

6

 3

d

6

 4

s

2

  

c.

 

1

s

2

 2

s

2

 2

p

6

 3

s

2

 3

p

6

 4

s

2

 3

d

6

  

d.

 

1

s

2

 2

s

2

 2

p

6

 3

s

2

 3

p

6

 4

s

2

 3

d

4

 4

p

2

  

e.

 

none of the above  

5. Which of the following has the largest radius?  

a.

 

F  

b.

 

N  

c.

 

C  

d.

 

O  

e.

 

Ne  

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6. Which of the following elements has the largest ionization energy?  

a.

 

Na  

b.

 

Ne  

c.

 

F  

d.

 

K  

e.

 

Rb  

7. Which of the following has the greatest electron affinity (most negative value)?  

a.

 

Cl  

b.

 

K  

c.

 

He  

d.

 

Na  

e.

 

Rb  

8. Which of the following species is not isolectronic with any of the others?  

a.

 

V

3+

  

b.

 

Ca

2+

  

c.

 

Ar  

d.

 

Cl

-

  

e.

 

S

2-

  

9. In Bohr's model of the hydrogen atom, the radius of an orbit  

a.

 

is proportional to 

n

2

.  

b.

 

is smallest for the highest energy state.  

c.

 

increases when a photon of light is emitted from an excited atom.  

d.

 

can have any value that is larger than the ground-state radius.  

e.

 

none of the above  

10. Which of the following atoms is 

not

 a one-electron system?  

a.

 

H  

b.

 

He

+

  

c.

 

Li

2+

  

d.

 

Be

2+

  

e.

 

O

7+

 

 

11. Which of the following statements about periodic properties is incorrect?  

a.  Both electron affinity and ionization energy decrease down a group.  
b.  Atomic size increases to the right across a period.  
c.  Ionization energy increases to the right across a period.  
d.  Atomic size increases down a group.  
e.  Electron affinity increases to the right across a period.  
 

 

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Bonding 

1. Which one of the following is most likely to be an ionic compound?  

a.

 

HNF

2

  

b.

 

H

2

CO  

c.

 

N

2

H

4

  

d.

 

CaCl

2

  

e.

 

CH

3

Cl  

2. In which of the following processes does the enthalpy change (

Δ

H

) directly represent 

the magnitude of the lattice energy of KCl(

s

)?  

a.

 

Cl

2

(

g

) + 2K(

s

 2KCl(

s

)  

b.

 

KCl(

s

 K

+

(

aq

) + Cl

-

(

aq

)  

c.

 

KCl(

s

 K

+

(

g

) + Cl

-

(

g

)  

d.

 

KCl(

s

 K(

s

) + Cl

-

(

g

)  

e.

 

KCl(

s

 K(

s

) + Cl(

g

)  

3. Order the following by increasing bond strength: N

N, N=N, N-N  

a.

 

N

N, N=N, N-N  

b.

 

N

N, N-N, N=N  

c.

 

N-N, N=N, N

N  

d.

 

N=N, N-N, N

N  

e.

 

N=N, N

N, N-N  

4. Which of the following compounds has the greatest bond polarity?  

a.

 

PH

3

  

b.

 

NH

3

  

c.

 

HF  

d.

 

H

2

S  

e.

 

CH

4

  

5. Which of the following is not planar?  

a.

 

BCl

3

  

b.

 

ClF

3

  

c.

 

PCl

3

  

d.

 

XeF

4

  

e.

 

C

2

H

4

  

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6. Use VSEPR theory to predict the ideal bond angles around the two carbon atoms in 
acetaldehyde, CH

3

CHO. (The first carbon has single bonds to three H atoms and one C 

atom; the second carbon has single bonds to C and H, and a double bond to O.)  

a.

 

109°, 109°  

b.

 

109°, 120°  

c.

 

120°, 109°  

d.

 

120°, 90°  

e.

 

105°, 105°  

7. In a carbon-carbon triple bond, what is the nature of the bonding between the carbons?  

a.

 

two 2

s

 orbitals overlapping  

b.

 

two 2

p

 orbitals overlapping  

c.

 

two 

sp

 orbitals overlapping, two 2

p

y

 overlapping and two 2

p

z

 overlapping  

d.

 

an 

sp

 and 

sp

2

 overlapping and 2

p

 orbitals overlapping  

e.

 

an 

sp

2

 and 

sp

2

 overlapping and 2

p

 orbitals overlapping   

8. Which of the following molecules has 

sp

3

 hybridization and a dipole moment?  

a.

 

SiH

4

  

b.

 

BF

3

  

c.

 

NH

3

  

d.

 

BrF

3

  

e.

 

PCl

5

 

 

9. In the molecular orbital description of bonding in benzene (C

6

H

6

), how many electrons 

occupy delocalized  MOs?  

a. 2 

 

b. 3  
c. 4 

 

d. 5  
e. 6 

 

10. In which of the following species is the octet rule violated by the central atom?  

a. CH

4

  

b. SF

4

  

c. PCl

4

+

  

d. SO

2

  

e. NH

3

  

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11. The number of electron dots in the Lewis symbol for an element equals the  

a.

 

number of outermost 

p

 electrons.  

b.

 

number of electrons needed to fill the outermost 

p

 orbital.  

c.

 

period number that contains the element.  

d.

 

number of outermost 

s

 and 

p

 electrons.  

e.

 

number of outermost 

s

 electrons. 

 

 
Phases and Phase Equilibria  

1. Calculate the pressure of 0.55 mol of NH

3

 gas in a 2.00 L vessel at 25 °C, using the 

ideal gas law.  

a.

 

2.5 atm  

b.

 

6.7 atm  

c.

 

0.6 atm  

d.

 

7.5 atm  

e.

 

3.4 atm  

2. A steel tank contains carbon dioxide at 34 °C and is at a pressure of 13.0 atm. 
Determine the internal gas pressure when the tank and its contents are heated to 100 °C.  

a.

 

10.7 atm  

b.

 

9.4 atm  

c.

 

38.2 atm  

d.

 

1.9 atm  

e.

 

15.8 atm  

3. Deviations from the ideal gas law are less at:  

a.

 

high temperatures and high pressures  

b.

 

high temperatures and low pressures  

c.

 

low temperatures and high pressures  

d.

 

low temperatures and low pressures  

e.

 

high volumes and low temperatures  

4. A mixture of three gases has a pressure of 1380 mmHg at at 298 K. The mixture is 
analyzed and is found to contain 1.27 mol CO

2

, 3.04 mol CO, and 1.50 mol Ar. What is 

the partial pressure of Ar?  

a.  238 mm Hg  
b.  302 mm Hg  
c.  356 mm Hg  
d.  1753 mm Hg  
e.  8018 mm Hg  

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5. Which of the following exhibits the most hydrogen bonding?  

a.

 

LiH  

b.

 

CH

4

  

c.

 

NH

3

  

d.

 

H

2

S  

e.

 

CH

2

F

2

  

6. Which of the following carbon compounds has the highest melting point?  

a.

 

CF

4

  

b.

 

CCl

4

  

c.

 

CBr

4

  

d.

 

CI

4

  

e.

 

CH

4

  

7. Water has such a high specific heat because  

a.

 

it has such a low molecular weight.  

b.

 

it is rather dense.  

c.

 

the O-H single bond has a high bond energy.  

d.

 

it has many relatively strong hydrogen bonds.  

e.

 

it dissolves both ionic and covalent compounds.  

8. The triple point is  

a.  an end to the liquid-gas line in a phase diagram.  
b.  the relationship between the boiling point, melting point and vapor pressure of a 

substance.  

c.  the point on a phase diagram where solid, liquid, and gas are in equilibrium.  
d.  the three pieces of data needed to solve the Clausius-Clapeyron equation.  
e.  the (P,V,T) coordinate of a point on a phase diagram.  

9. The main forces responsible for the structure of DNA are  

a.  ionic bonds and covalent bonds.  
b.  covalent bonds and ionic bonds.  
c.  hydrogen bonds and dipole-dipole interactions.  
d.  covalent bonds and hydrogen bonds.  
e.  covalent bonds and dipole-dipole interactions.  

10. Which of the following is not likely to exhibit hydrogen bonding?  

a. CH

3

CH

2

OH  

b. CH

3

NH

2

  

c. H

2

O  

d. NH

2

OH  

e. (CH

3

)

 3

N  

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Stoichiometry 

 

1. What is the mass of one mole of acetylsalicylic acid (aspirin), C

9

H

8

O

4

?  

a.

 

29 g  

b.

 

108 g  

c.

 

196 g  

d.

 

180. g  

e.

 

none of the above  

2. Determine the number of moles of aluminum in 2.154 x 10

-1

 kg of Al.  

a.

 

5816 mol  

b.

 

7.984 mol  

c.

 

6.02 x 10

23

 mol  

d.

 

4.801 mol  

e.

 

8.783 mol  

3. How many grams of zinc are there in 22.7 g of ZnCl

2

?  

a.

 

0.35 g  

b.

 

0.17 g  

c.

 

10.9 g  

d.

 

1476 g  

e.

 

0.32 g  

4. A compound with a composition of 87.5 % N and 12.5 % H was recently discovered. 
What is the empirical formula for this compound?  

a.

 

NH

2

  

b.

 

N

2

H

3

  

c.

 

NH  

d.

 

N

2

H

2

  

e.

 

N

2

H  

5. This equation is unbalanced: PCl

3

 + H

2

 H

3

PO

3

 + HCl When it is correctly 

balanced, the coefficients are, respectively  

a.

 

1,3,1,1  

b.

 

1,1,1,3  

c.

 

1,3,1,3  

d.

 

2,3,2,3  

e.

 

none of the above  

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6. Given 6 mol of each reactant, which one would be limiting in the following reaction? 
      4Au + 8NaCN + O

2

 + 2H

2

 4NaAu(CN)

2

 + 4NaOH  

a.

 

Au  

b.

 

NaCN  

c.

 

O

2

  

d.

 

H

2

O  

e.

 

There is no limiting reactant.  

7. In the direct reaction of silicon with Cl

2

 the yield of SiCl

4

 is 50. %. How many grams 

of silicon must be reacted with excess chlorine in order to obtain 17 g SiCl

4

?  

a.

 

1.4 g  

b.

 

2.8 g  

c.

 

5.6 g  

d.

 

17 g  

e.

 

28 g  

8. In the reaction of Fe

3

O

4

 with carbon to form carbon dioxide and iron, the number of 

moles of carbon required to convert 23 g of Fe

3

O

4

 to products is  

a. 0.05 

 

b. 0.1  
c. 0.2 

 

d. 0.3  
e. 0.4 

 

9. A 20.0 mL sample of an element with a density of 3.0 g/mL contains 4 x 10

23

 atoms. 

What is the atomic weight of this element?  

a. 300 

 

b. 40  
c. 60 

 

d. 90  
e.  none of the above  

10. How many moles of oxygen gas will react with 12.4 mol aluminum? 
Equation: 4Al + 3O

2

 

 2Al

2

O

3

  

a. 0.24 mol 

 

b. 0.42 mol  
c. 4.8 mol 

 

d. 9.3 mol  
e. 16.8 mol 

 

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11. Balance the following redox equation occurring in aqueous solution:  
KMnO

4

 + KCl + H

2

SO

4

 

 MnSO

4

 + K

2

SO

4

 + H

2

O + Cl

2

   

What is the stoichiometric coefficient for chlorine (Cl

2

) when the equation is balanced 

with smallest whole number coefficients? 

a. 1 

 

b. 3  
c. 5 

 

d. 8  
e. 10 

 

 
 
Thermodynamics and Thermochemistry 

 

1. Data: 
(1) H

2

(

g

) + ½O

2

(

g

 H

2

O(

g

Δ

H

 = -241.8 kJ  

(2) H

2

(

g

) + ½O

2

(

g

 H

2

O(

l

Δ

H

 = -285.8 kJ  

On the basis of the above data, which of the following statements is 

false

?  

a.

 

Reaction (1) is exothermic.  

b.

 

Reaction (2) is the formation reaction for H

2

O(

l

).  

c.

 

The reverse of reaction (2) is endothermic.  

d.

 

The energy content of H

2

O(

g

) is lower than H

2

O(

l

).  

e.

 

Δ

H

 for the reaction: H

2

O(

l

 H

2

O(

g

) is + 44 kJ/mol.  

2. What is the amount of heat necessary to raise the temperature of 8.5 kg of water from 
12.5 °C to 84 °C?  

a.

 

3.0 x 10

3

 kJ  

b.

 

36 J  

c.

 

2.5 x 10

3

 kJ  

d.

 

2.5 x 10

6

 kJ  

e.

 

25 kJ  

3. Data:   

Δ

H

°

f

 values: CH

4

(

g

), -74.8 kJ; CO

2

(

g

), -393.5 kJ; H

2

O(

l

), -285.8 kJ. 

Using the 

Δ

H

°

f

 data above, calculate 

Δ

H

°

rxn

 for the reaction below. 

Reaction: CH

4

(

g

) + 2O

2

(

g

 CO

2

(

g

) + 2H

2

O(

l

)  

a.

 

-604.2 kJ  

b.

 

890.3 kJ  

c.

 

-997.7 kJ  

d.

 

-890.3 kJ  

e.

 

none of the above  

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4. Data: 
2Ba(

s

) + O

2

(

g

 2BaO(

s

)    

Δ

H

° = -1107.0 kJ  

Ba(s) + CO

2

(

g

) + ½O

2

(

g

 BaCO

3

(

s

)   

Δ

H

° = -822.5 kJ  

Given the data above, calculate 

Δ

H

° for the reaction below. 

Reaction: BaCO

3

(

s

 BaO(

s

) + CO

2

(

g

)   

a.

 

-1929.5 kJ  

b.

 

-1376.0 kJ  

c.

 

-284.5 kJ  

d.

 

269.0 kJ  

e.

 

537 kJ  

5. Which of the following is 

not

 a state function?  

a.

 

Δ

E

  

b.

 

Δ

H

  

c.

 

q

  

d.

 

P

  

e.

 

V

  

6. Two solutions (the system), each of 25.0 mL volume and at 25.0 °C, are mixed in a 
beaker. A reaction occurs between them, causing the temperature to drop to 20.0 °C. 
After the products have equilibrated with the surroundings, the temperature is again 25.0 
°C and the total volume is 50.0 mL. No gases are involved in the reaction. Which one of 
the following relationships concerning the change from initial to final states (both at 25.0 
°C) is correct?  

a. 

Δ

E

 = 0  

b. 

Δ

H

 = 0  

c. 

Δ

E

 < 0  

d. 

w

 = 0  

e. 

q

 = 0  

7. Which one of the following processes is exothermic?  

a. H

2

(

l

 H

2

(

g

)  

b. CO

2

(

s

 CO

2

(

g

)  

c. H

2

O(

g

 H

2

O(

l

)  

d. 16CO

2

(

g

) + 18H

2

O(

l

 2C

8

H

18

(

l

) + 25O

2

(

g

)  

e. H

2

(

g

 2H(

g

)  

8. Predict the signs of 

Δ

H°, 

Δ

S°, and 

Δ

G° for the vaporization of liquid water at 150°C.  

a.

 

Δ

H° > 0, 

Δ

S° > 0, 

Δ

G° > 0  

b.

 

Δ

H° < 0, 

Δ

S° < 0, 

Δ

G° < 0  

c.

 

Δ

H° > 0, 

Δ

S° < 0, 

Δ

G° > 0  

d.

 

Δ

H° > 0, 

Δ

S° > 0, 

Δ

G° < 0  

e.

 

none of the above  

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9. Which of the following substances has the lowest standard molar entropy (

S

°) at 25°C?  

a.

 

CH

3

OH(

l

)  

b.

 

CO(

g

)  

c.

 

MgO(

s

)  

d.

 

H

2

O

(l)

  

e.

 

CaCO

3

(

s

)  

10. When crystalline solid barium hydroxide octahydrate and crystalline solid ammonium 
nitrate are mixed in a beaker at room temperature, a spontaneous reaction occurs. The 
temperature of the beaker contents rapidly falls to below 0º C. Use this information to 
decide whether the reaction is exothermic or endothermic and what the signs of 

Δ

H and 

Δ

S are.  

a.

 

endothermic; 

Δ

H > 0;

Δ

 S > 0  

b.

 

exothermic; 

Δ

H < 0;

Δ

 S > 0  

c.

 

endothermic; 

Δ

H < 0;

Δ

 S < 0  

d.

 

endothermic; 

Δ

H < 0;

Δ

 S > 0  

e.

 

exothermic; 

Δ

H > 0;

Δ

 S < 0  

11. Sodium carbonate can be made by heating sodium hydrogen carbonate: 
2NaHCO

3

(

s

 Na

2

CO

3

(

s

) + CO

2

(

g

) + H

2

O(

g

For this reaction, 

Δ

H

° = 128.9 kJ and 

Δ

S

° = 321 J/K. At approximately what temperature 

will 

K

 = 1? 

a.

 

401.6 K  

b.

 

401.6º C  

c.

 

33.1 K  

d.

 

33.1º C  

e.

 

none of the above  

 
Rate Processes in Chemical Reactions—Kinetics and Equilibrium 

 

1. For the overall hypothetical reaction A + 5B 

 4C, the rate of appearance of C given 

by 

Δ

[C]/

Δ

t is the same as  

a.

 

Δ

[A]/

Δ

t  

b.

 

-(5/4)(

Δ

[B]/

Δ

t)  

c.

 

-(4/5)(

Δ

[B]/

Δ

t)  

d.

 

-(1/4)(

Δ

[A]/

Δ

t)  

e.

 

none of the above.  

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2. The initial rate of the reaction 
PCl

5

 

 PCl

3

 + Cl

2

 

is increased a factor of four when the concentration of PCl

is doubled. Therefore, the rate  

a.

 

depends on the concentrations of PCl

3

 and Cl

2

.  

b.

 

is first order with respect to PCl

5

.  

c.

 

is second order with respect to PCl

5

.  

d.

 

is fourth order with respect to PCl

5

.  

e.

 

is first order with respect to PCl

3

.  

3. Consider the reaction A 

 products. Which of the following plots is consistent with a 

zero-order reaction?  

a.

 

[A] plotted against time gives a horizontal, straight line.  

b.

 

In [A] plotted against time gives a straight line of negative slope.  

c.

 

1/[A] plotted against time gives a straight line of positive slope.  

d.

 

[A] plotted against time gives a straight line of negative slope.  

e.

 

[A] plotted against time gives a curved line of negative slope, decreasing in 
magnitude as time increases 

4. The rate constant of a first-order reaction is 3.68 x 10

-2

 s

-1

 at 150°C, and the activation 

energy is 71 kJ/mol. What is the value of the rate constant at 170°C?  

a.

 

9.2 x 10

-2

 s

-1

  

b.

 

3.7 x 10

-2

 s

-1

  

c.

 

2.49 s

-1

  

d.

 

4.0 x 10

-2

 s

-1

  

e.

 

none of the above  

5. The reaction 
3ClO

-

(

aq

 ClO

3

-

(

aq

+ 2Cl

-

(

aq

) has been proposed to occur by the following mechanism. 

ClO

-

(

aq

) + ClO

-

(

 aq

)

 ClO

2

-

(

aq

) + Cl

-

(

 aq

)  (slow) 

ClO

2

-

(

aq

) + ClO

-

(

aq

 ClO

3

-

(

aq

) + Cl

-

(

aq

)  (fast) 

Which rate law is consistent with this mechanism? 

a.

 

rate = 

k

[ClO

-

]  

b.

 

rate = 

k

 [ClO

-

]

3

  

c.

 

rate = 

k

 [ClO

2

-

][ClO

-

]  

d.

 

rate = 

k

 [ClO

-

]

2

  

e.

 

rate = 

k

 [Cl

-

][ClO

-

]

2

  

6. A catalyst speeds up a reaction by  

a.

 

increasing the number of high-energy molecules.  

b.

 

increasing the temperature of the molecules in the reaction.  

c.

 

increasing the number of collisions between molecules.  

d.

 

increasing the activation energy for the reaction.  

e.

 

providing a new reaction pathway for molecules.  

13- Page
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7. Consider the following gas-phase equilibrium:  
H

2

(g) + I

2

(g) 

 2HI(g)  

At a certain temperature, the equilibrium constant 

K

c

 is 4.0. Starting with equimolar 

quantities of H

2

 and I

and no HI, when equilibrium was established, 0.20 moles of HI 

was present. How much H

2

 was used to start the reaction?  

a.

 

0.10 mol  

b.

 

0.23 mol  

c.

 

0.20 mol  

d.

 

4.0 mol  

e.

 

Need to know the volume of the reaction vessel.  

8. At a certain temperature the equilibrium constant 

K

p

 = 0.132 for the reaction:  

PCl

5

(g) 

 PCl

3

(g) + Cl

2

(g)  

At equilibrium, the partial pressures of both PCl

5

 and PCl

3

 are 100. mmHg. What is the 

total pressure of the equilibrium system, in mmHg?  

a.

 

100. mmHg  

b.

 

200. mmHg  

c.

 

300. mmHg  

d.

 

400. mmHg  

e.

 

332 mmHg  

9. Ammonium iodide dissociates reversibly to ammonia and hydrogen iodide:  
NH

4

I(s) 

 NH

3

(g) + HI(g)  

At 400ºC, 

K

p

 = 0.215. If 150 g of ammonium iodide is placed into a 3.00-L vessel and 

heated to 400º C, calculate the partial pressure of ammonia when equilibrium is reached.  

a.

 

0.22 atm  

b.

 

0.46 atm  

c.

 

0.11 atm  

d.

 

0.88 atm  

e.

 

1.2 atm  

10. Consider the equilibrium reaction:  
3CIO

-

(aq) 

 CIO

3

-

(aq) + 2CI

-

(aq)  

The equilibrium constant 

K

c

 = 3.2 X 10

3

. The following concentrations are present:  

[Cl

-

] = 0.50 mol/L; [ClO

3

-

] = 0.32 mol/L; [ClO

-

] = 0.24 mol/L. Is the mixture at 

equilibrium and, if not, in which direction will reaction proceed?  

a.

 

The system is at equilibrium.  

b.

 

The system is not at equilibrium; reaction will proceed left to right.  

c.

 

The system is not at equilibrium; reaction will proceed right to left.  

d.

 

The system cannot reach equilibrium since the ClO

3

-

 and Cl

-

 concentrations are 

not in the stoichiometric ratio.  

e.

 

There is not enough information to tell.  

14- Page
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11. Consider the following reaction in the gas phase:  
H

2

 + I

2

 

 2HI  

If the pressure increased by reducing the the volume of the flask,  

a.

 

more HI will be produced.  

b.

 

more H

2

 and I

2

 will be produced.  

c.

 

the results will depend on what the amounts of each are.  

d.

 

the amount of HI will remain the same.  

e.

 

the equilibrium constant will change.  

 
Solution Chemistry

 

1. Which of the following ions has an 

incorrect

 charge?  

a.

 

N

3-

  

b.

 

Al

3+

  

c.

 

S

2-

  

d.

 

Cl

-

  

e.

 

Mg

2-

  

2. Which of the following pairs of elements would be most likely to form an ionic 
compound?  

a.  P and Br  
b.  Zn and K  
c.  C and O  
d.  Al and Rb  
e.  F and Ca 

 3. What is the name of NaI?  

a.  sodium iodide  
b.  sodium(I) iodide  
c. sodium monoiodide 

 

d.  sodious iodide  
e. sodium 

iodine 

 4. Which of the following combinations of names and formulas is incorrect?  

a. H

3

PO

4

 phosphoric acid  

b. HNO

3

 nitric acid  

c. NaHCO

3

 sodium carbonate  

d. H

2

CO

3

 carbonic acid  

e.  KOH potassium hydroxide  

15- Page
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5. Calculate the concentration of calcium ions in a saturated calcium phosphate solution. 
(

K

sp

 for Ca

3

(PO

4

)

2

 = 1.3 X 10

-26

)  

a.

 

1.2 x 10

-5

 mol/L  

b.

 

2.0 x 10

-5

 mol/L  

c.

 

6.6 x 10

-6

 mol/L  

d.

 

7.8 x 10

-6

 mol/L  

e.

 

8.3 x 10

-6

 mol/L  

6. Calculate the molar solubility of silver carbonate in 1.0 

M

 sodium carbonate solution.  

(

K

sp

 for Ag

2

CO

3

 = 8.1 x 10

-12

)  

a.

 

8.1 x 10

-12

 

M

  

b.

 

2.8 x 10

-6

 

M

  

c.

 

1.4 x 10

-6

 

M

  

d.

 

1.4 x 10

-8

 

M

 

e.

 

2.0 x 10

-4

 

M

 

7. Calculate the pH of a solution necessary to just begin the precipitation of Mg(OH)

2

 

when [Mg

2+

] = 0.001 

M

. (

K

sp

 for Mg(OH)

2

 = 1.2 x 10

-11

a. 11 

 

b. 10  
c. 9 

 

d. 8  
e. 4 

 8. In qualitative analysis, the metals of Ion Group 1 can be separated from other ions by 
precipitating them as chloride salts. A solution initially contains Ag

+

 and Pb

2+

 at a 

concentration of 0.10 

M

. Aqueous HCl is added to this solution until the Cl

-

 concentration 

is 0.10 

M

. What will the concentrations of Ag

+

 and Pb

2+

 be at equilibrium? 

(

K

sp 

for AgCl = 1.8 x 10

-10

K

sp

 for PbCl

2

 = 1.7 x 10

-5

)  

a. [Ag

+

] = 1.8 x 10

-11 

M

; [Pb

2+

] = 1.7 x 10

-6 

M

  

b. [Ag

+

] = 1.8 x 10

-7 

M

; [Pb

2+

] = 1.7 x 10

-4 

M

  

c. [Ag

+

] = 1.8 x 10

-11 

M

; [Pb

2+

] = 8.5 x 10

-5 

M

  

d. [Ag

+

] = 1.8 x 10

-9 

M

; [Pb

2+

] = 1.7 x 10

-3 

M

  

e. [Ag

+

] = 1.8 x 10

-9 

M

; [Pb

2+

] = 8.5 x 10

-6 

M

  

9. Silver chloride is relatively insoluble in water (

K

sp

 for AgCl = 1.8 x 10

-10

) but it is 

soluble in aqueous ammonia, due to the formation of the complex ion Ag(NH

3

)

2

+

. How 

many moles of AgCl will dissolve in 1.00 L of solution containing 6.0 moles of free 
NH

3

? (

K

f

 for Ag(NH

3

)

2

+

 = 1.7 x 10

7

)  

a.  9.1 x 10

-6

 mol  

b.  2.9 x 10

-4

 mol  

c. 0.0091 mol 

 

d. 0.084 mol  
e. 0.33 mol 

 

16- Page
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10. What is the mass of C

12

H

22

O

11

 in 60.0 mL of 0.0880 

M

 solution?  

a.

 

0.181 g  

b.

 

1.81 g  

c.

 

5.02 g  

d.

 

5.28 g  

e.

 

none of the above  

11. The freezing point of pure camphor is 178.4 °C, and its molal freezing-point constant, 
K

f

 is 40.0 °C/

m

. Find the freezing point of a solution containing 3.00 g of a compound of 

molar mass 125 g/mol in 45.0 g of camphor.  

a.  174.1 °C  
b.  157.1 °C  
c.  135.2 °C  
d.  140.4 °C  
e.  11.6 °C  

 
 
Acids and Bases  

1. Calculate the 

hydroxide

 ion concentration of a solution if its pH is 6.389.  

a.

 

1.00 x 10

-14

 mol/L  

b.

 

4.08 x 10

-7

 mol/L  

c.

 

9.92 x 10

-7

 mol/L  

d.

 

2.45 x 10

-8

 mol/L  

e.

 

none of the above  

2. Which of the following is a correct description of the natural direction of a Brønsted-
Lowry acid-base reaction?  

a.

 

weaker acid + weaker base 

 stronger acid + stronger base  

b.

 

weaker acid + stronger base 

 stronger acid + weaker base  

c.

 

stronger acid + weaker base 

 weaker acid + stronger base  

d.

 

stronger acid + stronger base 

 weaker acid + weaker base  

e.

 

None of the above statements is always correct.  

3. In a 0.100 

M

 HF solution, the percent dissociation is determined to be 9.5%. Calculate 

the 

K

a

 for HF based on this data.  

a.

 

9.5 x 10

-2

  

b.

 

1.0 x 10

-3 

 

c.

 

3.1 x 10

-3

  

d.

 

7.6 x 10

-4

  

e.

 

9.5 x 10

-4

  

17- Page
background image

4. What is the pH of a solution prepared from 0.250 mol of NH

3

 dissolved in sufficient 

water to make 1.00 L of solution? (

K

b

 = 1.8 x 10

-5

)  

a.

 

2.12  

b.

 

2.67  

c.

 

8.92  

d.

 

11.33  

e.

 

13.40  

5. Which of the following reactions illustrate Al(OH)

3

 acting as a Lewis acid?  

a.

 

Al(OH)

3

 

 Al

3+

 + 3OH

-

  

b.

 

Al(OH)

3

 + OH

-

 

 Al(OH)

2

O

-

 + H

2

O  

c.

 

Al(OH)

3

 + OH

-

 

 Al(OH)

4

-

  

d.

 

Al(OH)

3

 + 3H

+

 

 Al

3+

 + 3H

2

O  

e.

 

Al

3+

 + 3OH

-

 

 Al(OH)

3

  

6. Which of the following pairs of species is 

not

 a conjugate acid-base pair?  

a.  HCl and H

+

  

b. HSO

4

-

 and SO

4

2-

  

c. H

2

SO

4

 and HSO

4

-

  

d. H

2

O and OH

-

  

      e.  NH

3

 and NH

2

-

  

7. Consider each of the following pairs of acids. Which statement is correct?  

a. HClO

2

 is a stronger acid than HClO

4

.  

b. H

2

SO

4

 is a stronger acid than H

2

SeO

4

.  

c. H

2

O is a stronger acid than HF.  

d. H

2

S is a stronger acid than H

2

Se.  

      e.  HS

-

 is a stronger acid than H

2

S.  

8. Consider the reaction  
CH

3

NH

2

 + H

2

 CH

3

NH

3

+

 + OH

-

 

where CH

3

NH

is methylamine and CH

3

NH

3

+

 is the methylammonium ion. Select the 

correct description of this reaction in terms of Lewis acid-base theory.  

a.  Methylamine serves as a Lewis acid in the forward reaction and 

methylammonium ion serves as a Lewis base in the reverse reaction.  

b.  Water serves as a Lewis base in the forward reaction and the hydroxide ion serves 

as a Lewis base in the reverse reaction.  

c.  Methylamine serves as a Lewis base in the forward reaction and hydroxide ion 

serves as a Lewis acid in the reverse reaction.  

d.  Water serves as a Lewis acid in the forward reaction and methylammonium ion 

serves as a Lewis base in the reverse reaction.  

e.  Methylamine serves as a Lewis base in the forward reaction and hydroxide ion 

serves as a Lewis base in the reverse reaction.  

18- Page
background image

9. What is the pH of a buffer prepared by adding 180 mL of 0.100 

M

 NaOH to 200 mL of 

0.100 

M

 acetic acid?  

(

K

a

 for CH

3

COOH = 1.8 x 10

-5

)  

a. 3.79 

 

b. 4.34  
c. 4.74 

 

d. 5.04  

      e.  5.70  

10. Consider the titration of 50.00 mL of 0.1000 

M

 HBr with 0.1000 

M

 KOH. Calculate 

the pH after 49.00 mL of the base has been added to the 50.00 mL of HBr.  

a. 2.0 

 

b. 3.0  
c. 4.0 

 

d. 6.0  

      e.  7.0  

11. An aqueous solution of a weak acid, HA, is titrated with NaOH solution. The pH at 
the midpoint of the buffer region is 4.5. What is the 

K

a

 of the acid?  

a.  3.2 x 10

-5

  

b.  3.2 x 10

-10

  

c.  1.8 x 10

-3

  

d.  7.0 x 10

-7

  

e. 4.5 

 

 
 
Electrochemistry 

1. Which of the following statements is incorrect?  

a.

 

In an electrolytic cell, reduction occurs at the anode.  

b.

 

Aluminum metal would form at the cathode during the electrolysis of molten 
AlBr

3

.  

c.

 

The cathode is labeled "+" in a voltaic cell.  

d.

 

Oxidation occurs at the anode in a voltaic cell.  

e.

 

Electrons flow from the anode to the cathode in all electrochemical cells.  

2. Consider the following notation for an electrochemical cell 
ZnlZn

2+

 (1

M

)llFe

2+

 (1

M

), Fe

3+

 (1

M

)lPt 

What is the balanced equation for the cell reaction?  

a.

 

Zn(

s

) + 2Fe

3+

(

aq

 2Fe

2+

(

aq

) + Zn

2+

(

aq

)  

b.

 

Zn

2+

(

aq

) + 2Fe

2+

(

aq

 Zn(

s

) + 2Fe

3+

(

aq

)  

c.

 

Zn(

s

) + 2Fe

2+

(

aq

 2Fe

3+

(

aq

) + Zn

2+

(

aq

)  

d.

 

Zn(

s

) + Fe

3+

(

aq

 Fe

2+

(

aq

) + Zn

2+

(

aq

)  

e.

 

Zn(

s

) + Fe

2+

(

aq

 Fe(

s

) + Zn

2+

(

aq

)  

19- Page
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3. Consider the following half-reactions and select the strongest oxidizing agent present: 
Sr

2+

(

aq

) + 2e

-

 

 Sr(

s

)  

E

° = -2.89 V 

Cr

2+

(

aq

) + 2e

-

 

 Cr(

s

)   

E

° = -0.913 V 

Co

2+

(

aq

) + 2e

 Co(

s

)   

E

° = -0.28 V   

a.

 

Cr

2+

(

aq

)  

b.

 

Sr

2+

(

aq

)  

c.

 

Co

2+

(

aq

)  

d.

 

Sr(

s

)  

e.

 

Co(

s

4. In an electrolytic cell, how many grams of Cu could be plated out of a CuSO

4

 solution 

at a current of 5.00 A for 2.00 min? (

F

 = 96500 C/mol)  

a.

 

318 g  

b.

 

0.395 g  

c.

 

0.329 x 10

-3

 g  

d.

 

0.198 g  

e.

 

5.31 g  

5. A voltaic cell is based on the following two half-reactions: 
Ni

+2

(

aq

) + 2e

-

 

 Ni(

s

E

° = -0.25 V  

Cr

+3

(

aq

) + 3e

-

 

 Cr(

s

E

° = -0.74 V  

Sketch the cell and then select the correct statement about it.  

a.  Cr serves as the cathode.  
b.  The direction of electron flow through the external wire is from the Ni to the Cr 

electrode.  

c.  Anions in solution will migrate 

toward

 the Ni

+2

/Ni electrode.  

d.  The net cell reaction is 3Ni

+2

(

aq

) + 2Cr(

s

 3Ni(

s

) + 2Cr

+3

(

aq

)  

e. 

E

°

cell

 = 0.99 V  

6. Consider the following two electrode reactions and their standard electrode potentials:  
Al

+3

(

aq

) + 3e

-

 

 Al (

s

E

° = -1.66 V  

Cd

+2

(

aq

) + 2e

-

 

 Cd(

s

E

° = -0.40 V  

Write the cell reaction for a voltaic cell based on these two electrodes, and calculate the 
standard cell potential, 

E

°

cell

a. 2Al

+3

(

aq

) + 3Cd

+2

(

aq

 2Al(

s

) + 3Cd(

s

E

°

cell

 = 2.10 V 

b. 2Al(

s

) + 3Cd

+2

(

aq

 2Al

+3

(

aq

) + 3Cd(

s

E

°

cell

 = 1.26 V  

c. 2Al(

s

) + 3Cd

+2

(

aq

 2Al

+3

(

aq

) + 3Cd(

s

E

°

cell

 = 3.78V  

d. 2Al

+3

(

aq

) + 3Cd(

s

 2Al(

s

) + 3Cd

+2

(

aq

E

°

cell

 = 1.26 V  

e. 2Al

+3

(

aq

) + 3Cd(

s

 2Al(

s

) + 3Cd

+2

(

aq

)

 E

°

cell

 = 2.10 V  

20- Page
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7. Use the following standard electrode potentials to predict the species formed at the 
electrodes in the electrolysis of aqueous CuSO

4

O

2

(

g

) + 4H

+

(

aq

) + 4e

-

 

 2H

2

O(

l

E

° = +1.23 V  

Cu

2+

(

aq

) + 2e

-

 

 Cu(

s

E

° = +0.34 V  

SO

4

2-

(

aq

) + 4H

+

(

aq

) + 2e

-

 

 H

2

SO

3

(

aq

) + H

2

O(

l

E

° = 0.20 V 

2H

+

(

aq

) + 2e

-

 

 H

2

(

g

E

° = 0.00 V   

a. H

2

, O

2

, H

+

  

b. Cu, O

2

, H

+

  

c. Cu, 

H

2

  

d. H

2

, H

2

SO

3

, H

2

O  

e. H

2

SO

3

, H

2

O, O

2

, H

+

  

  
8. A constant current was passed through a solution of KAuCl

between gold electrodes. 

Over a period of 20.00 min, the cathode increased in mass by 2.664 g. What was the 
current in amperes?  
(

F

 = 96500 C/mol) 

Cathode half-reaction: AuCl

4

-

(

aq

) + 3e

-

 

 Au(

s

) + 4Cl

-

(

aq

)  

 

a.  1.08 A  
b.  3.26 A  
c.  2.17 A  
d.  6.52 A  
e.  3.48 A  

9. A voltaic cell is constructed from the following half-cells, linked by a KCl salt bridge: 
(a) an Fe electrode in 1.0 

M

 FeCl

2

 solution 

(b) a Ni electrode in 1.0 

M

 Ni(NO

3

)

2

 solution 

Use the table of standard electrode potentials in your textbook to decide which one of the 
following statements is correct.  

a.

 

The Ni electrode is the anode.  

b.

 

Electrons flow from the iron electrode to the nickel electrode.  

c.

 

The iron electrode is positively charged.  

d.

 

The iron electrode will gain mass when current flows.  

e.

 

The salt bridge conducts electrons through solution.  

10. Which one of the following reactions must be carried out in an electrolytic cell, rather 
than a voltaic cell?  

a.

 

Zn + Cd

2+

 

 Cd + Zn

2+

  

b.

 

Al + 3/2Br

2

 

 Al

3+

 + 3Br

-

  

c.

 

2Al

3+

 + 3Fe 

 2Al + 3Fe

2+

  

d.

 

H

2

 + I

2

 

 2H

+

 + 2I

-

  

e.

 

2H

2

 + O

2

 

 2H

2

 

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