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AP Chemistry Full-length practice exam 3

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Full-length practice exam 3

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Question 1

Rank these liquids in order of increasing viscosity: ethanol ( $C_{2} H_{5} OH$ ), glycerol ( $C_{3} H_{5} (OH)_{3}$ ), and hexane ( $C_{6} H_{14}$ ).

A.Hexane

<

Ethanol

<

GlycerolB.Ethanol

<

Hexane

<

GlycerolC.Glycerol

<

Ethanol

<

HexaneD.Hexane

<

Glycerol

<

Ethanol

Question 2

What is the electron-pair geometry and molecular geometry of $S F_{4}$ ?

A.Tetrahedral; tetrahedralB.Trigonal bipyramidal; trigonal bipyramidalC.Trigonal bipyramidal; see-sawD.Octahedral; square planar

Question 3

Identify the acid-base behavior of $A l^{3 +}$ dissolved in water according to Lewis theory.

A l^{3 +}

is a Lewis base because it donates lone pairs to water.B.

A l^{3 +}

is a Lewis acid because it accepts lone pairs from water molecules.C.

A l^{3 +}

is a Brønsted–Lowry acid because it donates protons to water.D.

A l^{3 +}

has no acid-base behavior in water because it is a metal cation.

Question 4

A mixture of gas A (molar mass $4 g/mol$ ) and gas B (molar mass $64 g/mol$ ) effuses through a small orifice. According to Graham's Law, how much faster does gas A effuse compared to gas B?

A.4 times fasterB.16 times fasterC.

4 = 2

times fasterD.8 times faster

Question 5

The rate of decomposition of $N_{2} O_{5}$ is measured at two temperatures:

Temperature	Rate constant $k$
$25° C$	$3.0 \times 1 0^{- 5} s^{- 1}$
$45° C$	$1.2 \times 1 0^{- 4} s^{- 1}$

Which statement about this data is correct?

A.The reaction rate doubles for every

10° C

increase because that is a chemistry rule.B.The rate constant increases by a factor of 4 over a 20°C range, consistent with Arrhenius behavior.C.The rate constant doubles when temperature is doubled, so the reaction is second order in temperature.D.The rate constant is independent of temperature for a first-order reaction.

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Question 6

Which of the following best explains why ionic compounds have high melting points compared with molecular compounds?

A.Ionic compounds are denser, so more energy is needed to expand the lattice.B.Ionic bonds are covalent in nature and therefore very strong.C.Electrostatic forces between oppositely charged ions in the crystal lattice are very strong.D.Ionic compounds cannot conduct electricity, which raises their melting points.

Question 7

For a first-order reaction with rate constant $k = 0.0231 mi n^{- 1}$ , what is the half-life?

43.3 min

30.0 min

0.0231 min

1.00 min

Question 8

The concentration of a reactant A in a first-order reaction decreases from $0.800 M$ to $0.100 M$ over $90 min$ . What is the half-life of the reaction?

11.25 min

30 min

45 min

90 min

Question 9

According to the kinetic molecular theory of gases, what happens to the average kinetic energy of gas molecules when temperature is doubled (from 300 K to 600 K)?

A.It doubles.B.It quadruples.C.It increases by

2

.D.It remains the same; only pressure changes.

Question 10

Which of the following thermochemical properties can be calculated from a Born-Haber cycle?

A.Rate constant of an ionic reactionB.Lattice energy of an ionic compoundC.Equilibrium constant for a dissolution reactionD.Bond order of an ionic compound

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Question 11

A reaction has the energy profile shown below (described): the reactants are at a higher energy than the products, and there is a single energy maximum (transition state) between them. The activation energy for the forward reaction is $80 kJ/mol$ and for the reverse reaction is $120 kJ/mol$ . What is $Δ H$ for the forward reaction?

- 40 kJ/mol

+ 40 kJ/mol

+ 200 kJ/mol

- 120 kJ/mol

Question 12

Given:

(1) $C (s) + O_{2} (g) \to C O_{2} (g)$ , $Δ H_{1} = - 393.5 kJ$

(2) $2 CO (g) + O_{2} (g) \to 2 C O_{2} (g)$ , $Δ H_{2} = - 566.0 kJ$

Use Hess's Law to find $Δ H$ for: $2 C (s) + O_{2} (g) \to 2 CO (g)$

- 959.5 kJ

- 221.0 kJ

+ 221.0 kJ

- 172.5 kJ

Question 13

Which of the following best describes an endothermic reaction?

A.Products have lower potential energy than reactants;

Δ H < 0

.B.The system absorbs heat from the surroundings;

Δ H > 0

.C.The reaction releases energy as light.D.Temperature of the surroundings increases during the reaction.

Question 14

What is the formal charge on the oxygen atom in $C O_{2}$ (a carbon–oxygen double bond)?

- 2

- 1

0

+ 1

Question 15

A solution is prepared by dissolving $5.85 g$ of NaCl (molar mass $58.5 g/mol$ ) in enough water to make $250 mL$ of solution. What is the molarity of NaCl?

0.100 M

0.400 M

2.50 M

0.0400 M

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Question 16

Which of the following aqueous solutions is the strongest base?

0.10 M

N H_{3}

(

K_{b} = 1.8 \times 1 0^{- 5}

)B.

0.10 M

NaOH

0.10 M

C H_{3} N H_{2}

(

K_{b} = 4.4 \times 1 0^{- 4}

)D.

0.10 M

N a_{2} C O_{3}

Question 17

For the equilibrium $N_{2} (g) + 3 H_{2} (g) ⇌ 2 N H_{3} (g)$ , which expression correctly represents $K_{c}$ ?

K_{c} = \frac{[ N H _{3} ] ^{2}}{[ N _{2} ] [ H _{2} ] ^{3}}

K_{c} = \frac{[ N _{2} ] [ H _{2} ] ^{3}}{[ N H _{3} ] ^{2}}

K_{c} = \frac{[ N H _{3} ]}{[ N _{2} ] [ H _{2} ]}

K_{c} = \frac{[ N H _{3} ] ^{2}}{[ N _{2} ] ^{2} [ H _{2} ] ^{3}}

Question 18

In a bomb calorimeter experiment, $1.00 g$ of glucose ( $M = 180 g/mol$ ) causes a temperature rise of $2.00° C$ in a calorimeter with heat capacity $C_{cal} = 7.78 kJ/°C$ . What is the molar heat of combustion of glucose?

- 15.6 kJ/g

- 2800 kJ/mol

- 5330 kJ/mol

- 156 kJ/mol

Question 19

Which type of intermolecular force is present in ALL molecular substances?

A.Hydrogen bondingB.Dipole–dipole forcesC.London dispersion forcesD.Ion–dipole forces

Question 20

For the dissolution of calcium fluoride, $Ca F_{2} (s) ⇌ C a^{2 +} (aq) + 2 F^{-} (aq)$ , the $K_{s p} = 3.9 \times 1 0^{- 11}$ . What is the molar solubility of $Ca F_{2}$ ?

2.1 \times 1 0^{- 4} M

3.9 \times 1 0^{- 11} M

6.2 \times 1 0^{- 6} M

1.3 \times 1 0^{- 4} M

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Question 21

When excess $NaOH (aq)$ is added to a solution of $FeC l_{3} (aq)$ , a reddish-brown precipitate forms. The full ionic equation includes $F e^{3 +} (aq) + 3 O H^{-} (aq) \to Fe (OH)_{3} (s)$ . What are the spectator ions?

N a^{+}

and

C l^{-}

F e^{3 +}

and

O H^{-}

N a^{+}

and

F e^{3 +}

C l^{-}

and

O H^{-}

Question 22

A proposed mechanism for a reaction is:

Step 1 (slow): $N O_{2} + N O_{2} \to N O_{3} + NO$

Step 2 (fast): $N O_{3} + CO \to N O_{2} + C O_{2}$

What is the overall reaction and the predicted rate law?

A.Overall:

N O_{2} + CO \to NO + C O_{2}

; rate

= k [N O_{2}] [CO]

B.Overall:

N O_{2} + CO \to NO + C O_{2}

; rate

= k [N O_{2}]^{2}

C.Overall:

2 N O_{2} + CO \to N O_{3} + NO + C O_{2}

; rate

= k [N O_{2}]^{2} [CO]

D.Overall:

N O_{2} + CO \to NO + C O_{2}

; rate

= k [N O_{3}] [CO]

Question 23

Which of the following processes has a positive entropy change ( $Δ S > 0$ )?

H_{2} O (l) \to H_{2} O (s)

(freezing)B.

N_{2} (g) + 3 H_{2} (g) \to 2 N H_{3} (g)

NaCl (s) \to N a^{+} (aq) + C l^{-} (aq)

C O_{2} (g) \to C O_{2} (s)

(dry ice)

Question 24

Which of the following is the conjugate base of $H_{2} P O_{4}^{-}$ ?

H_{3} P O_{4}

HP O_{4}^{2 -}

P O_{4}^{3 -}

H_{2} O

Question 25

According to Le Châtelier's principle, what happens to the equilibrium position of $N_{2} (g) + 3 H_{2} (g) ⇌ 2 N H_{3} (g)$ ( $Δ H = - 92 kJ$ ) when the temperature is increased?

A.Shifts forward; more NH

_{3}

is produced.B.Shifts reverse;

K_{c}

decreases.C.No effect; temperature does not affect

K_{c}

.D.Shifts forward;

K_{c}

increases.

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Question 26

Which of the following correctly ranks the entropy of three states of water at the same temperature?

S (ice) > S (water) > S (steam)

S (steam) > S (water) > S (ice)

S (water) > S (steam) > S (ice)

D.All three have equal entropy at the same temperature.

Question 27

How many grams of $H_{2} O$ are produced when $4.00 g$ of $H_{2}$ reacts completely with excess $O_{2}$ ?

$2 H_{2} (g) + O_{2} (g) \to 2 H_{2} O (g)$

(Molar masses: H $_{2}$ = 2.02 g/mol, H $_{2}$ O = 18.02 g/mol)

36.0 g

18.0 g

9.00 g

72.1 g

Question 28

Using average bond enthalpies, estimate $Δ H$ for the reaction $H_{2} (g) + C l_{2} (g) \to 2 HCl (g)$ , given: $D (H - H) = 436 kJ/mol$ , $D (Cl - Cl) = 243 kJ/mol$ , $D (H - Cl) = 432 kJ/mol$ .

+ 185 kJ/mol

- 185 kJ/mol

- 679 kJ/mol

+ 247 kJ/mol

Question 29

At constant temperature, the pressure of a fixed amount of gas is increased from $1.00 atm$ to $4.00 atm$ . By what factor does the volume change?

A.Increases by a factor of 4B.Decreases by a factor of 4C.Decreases by a factor of 2D.Remains unchanged

Question 30

A buffer solution is prepared containing $0.200 M$ acetic acid ( $C H_{3} COOH$ , $p K_{a} = 4.74$ ) and $0.200 M$ sodium acetate ( $C H_{3} COONa$ ). What is the pH of this buffer?

2.87

4.74

7.00

9.26

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Question 31

A gas occupies $3.00 L$ at $2.00 atm$ and $300 K$ . What volume will it occupy at $1.00 atm$ and $600 K$ , assuming ideal behavior?

3.00 L

6.00 L

12.0 L

1.50 L

Question 32

Spectral lines in the visible emission spectrum of hydrogen are produced when an electron transitions from a higher energy level to:

Select the statement that is both correct and explains the underlying mechanism.

n = 1

, releasing ultraviolet photons equal in energy to the transition.B.

n = 2

, releasing a photon whose energy equals the difference in quantized energy levels.C.

n = 3

, releasing infrared radiation.D.

n = 2

, absorbing a photon to move to a lower energy state.

Question 33

Which of the following species has a trigonal planar molecular geometry?

N H_{3}

S O_{3}

PC l_{3}

Cl F_{3}

Question 34

Which hybridization is consistent with a carbon atom that forms two double bonds and no single bonds (as in $C O_{2}$ )?

s p^{3}

s p^{2}

s p

s p^{3} d

Question 35

When a hot metal block is placed in cool water and the system reaches thermal equilibrium, which statement is correct?

A.Heat flows from the water to the metal until their temperatures equalize.B.The heat lost by the metal equals the heat gained by the water (assuming no heat loss to surroundings).C.The metal loses more heat than the water gains because metal is denser.D.The temperature changes of the metal and water must be equal.

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Question 36

The relationship between standard cell potential and Gibbs free energy is $Δ G^{\circ} = - n F E_{cell}^{\circ}$ , where $F = 96, 485 C/mol$ . For a cell with $E^{\circ} = + 0.50 V$ and $n = 2$ , what is $Δ G^{\circ}$ ?

+ 96.5 kJ/mol

- 96.5 kJ/mol

- 48.2 kJ/mol

- 1.00 kJ/mol

Question 37

Which of the following correctly orders the elements by increasing first ionization energy?

$Li, Na, K$

Li < Na < K

K < Na < Li

Na < K < Li

Li < K < Na

Question 38

In the reaction $N_{2} (g) + 3 H_{2} (g) \to 2 N H_{3} (g)$ , if $28.0 g$ of $N_{2}$ and $9.00 g$ of $H_{2}$ are mixed, which is the limiting reagent? (Molar masses: N $_{2}$ = 28.0 g/mol, H $_{2}$ = 2.02 g/mol)

N_{2}

, because it has a higher molar mass.B.

H_{2}

, because more moles of it are required per mole of product.C.

N_{2}

, because comparing moles available to the stoichiometric requirement shows N

_{2}

runs out first.D.Neither; both are present in stoichiometric amounts.

Question 39

The Bohr model correctly predicts the emission spectrum of hydrogen. However, it fails for multi-electron atoms primarily because it:

A.Assumes circular orbits and ignores electron–electron repulsion.B.Does not quantize electron energy levels.C.Ignores the wave nature of light entirely.D.Places electrons in the nucleus.

Question 40

The solubility product of silver chloride is $K_{s p} (AgCl) = 1.8 \times 1 0^{- 10}$ at $25° C$ . What is the molar solubility (in mol/L) of AgCl in pure water?

1.8 \times 1 0^{- 10} M

1.3 \times 1 0^{- 5} M

9.0 \times 1 0^{- 21} M

3.6 \times 1 0^{- 10} M

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Question 41

A student burns $1.00 g$ of ethanol in a calorimeter containing $500 g$ of water. The temperature rises from $22.5° C$ to $35.7° C$ . What is the heat absorbed by the water? ( $c_{water} = 4.18 J/ (g \cdot °C)$ )

2.76 kJ

27.6 kJ

55.2 kJ

0.276 kJ

Question 42

In which of the following does the reaction become spontaneous as temperature increases from low to high?

Δ H = - 80 kJ

Δ S = - 200 J/K

Δ H = + 60 kJ

Δ S = + 150 J/K

Δ H = - 60 kJ

Δ S = + 150 J/K

Δ H = + 80 kJ

Δ S = - 200 J/K

Question 43

Acetic acid has $K_{a} = 1.8 \times 1 0^{- 5}$ . What is the pH of a $0.10 M$ $C H_{3} COOH$ solution?

1.0

2.87

4.74

3.37

Question 44

Which of the following correctly explains why water has an unusually high surface tension?

A.Water molecules at the surface are attracted only by gravity, not by IMFs.B.Surface water molecules experience a net inward force due to extensive hydrogen bonding with interior molecules, resisting surface expansion.C.Water has a low molar mass, so surface molecules move slowly and resist disruption.D.London dispersion forces in water are exceptionally strong because of water's high electronegativity.

Question 45

A buffer contains $0.300 M$ $N H_{3}$ and $0.200 M$ $N H_{4} Cl$ . What is the pH? ( $K_{b} (N H_{3}) = 1.8 \times 1 0^{- 5}$ ; $p K_{a} (N H_{4}^{+}) = 9.26$ )

9.44

9.08

4.56

9.26

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Question 46

According to Coulomb's law, the energy of attraction between an electron and the nucleus is proportional to $\frac{Z _{eff}}{r}$ . Which isoelectronic pair would have the electron held most tightly (most negative energy)?

O^{2 -}

F^{-}

Ne

M g^{2 +}

Question 47

An Arrhenius plot of $ln k$ vs $1/ T$ gives a straight line with slope $- 8000 K$ . What is the activation energy $E_{a}$ ? ( $R = 8.314 J/ (mol \cdot K)$ )

9.63 kJ/mol

66.5 kJ/mol

8000 kJ/mol

960 kJ/mol

Question 48

An element's fifth ionization energy is dramatically larger than its fourth. What does this indicate about the element's electron configuration?

A.The element is in Group 4 of the periodic table.B.The element is in Group 5 of the periodic table.C.The element has five valence electrons.D.The fourth and fifth electrons occupy the same orbital.

Question 49

At equilibrium, increasing the pressure on a gas-phase reaction by reducing volume will shift the equilibrium in the direction that:

A.Decreases the total number of moles of gas.B.Increases the total number of moles of gas.C.Has no effect, since gases are compressible.D.Increases temperature.

Question 50

Which of the following describes a precipitation reaction?

HCl (aq) + NaOH (aq) \to NaCl (aq) + H_{2} O (l)

CaC l_{2} (aq) + N a_{2} S O_{4} (aq) \to CaS O_{4} (s) + 2 NaCl (aq)

Zn (s) + CuS O_{4} (aq) \to ZnS O_{4} (aq) + Cu (s)

2 H_{2} O_{2} (l) \to 2 H_{2} O (l) + O_{2} (g)

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Question 51

In a balanced redox equation in acidic solution, $Mn O_{4}^{-}$ is reduced to $M n^{2 +}$ while $F e^{2 +}$ is oxidized to $F e^{3 +}$ . How many $F e^{2 +}$ ions are required to reduce one $Mn O_{4}^{-}$ ?

2

3

5

7

Question 52

In a Lewis structure, three resonance structures can be drawn for $S O_{3}$ . What is the best interpretation of these structures?

A.The molecule rapidly switches between the three structures.B.The true structure is a hybrid in which the electron density is delocalized equally over all S–O bonds.C.The three structures represent three different

S O_{3}

isomers.D.Only the structure with the fewest formal charges is correct; the others are wrong.

Question 53

At $25° C$ , the enthalpy of combustion of hydrogen is $- 241.8 kJ/mol$ (as gas) when $H_{2} O$ is formed as a vapor. The enthalpy of vaporization of water is $44.0 kJ/mol$ . What is the enthalpy of combustion when liquid water is formed?

- 197.8 kJ/mol

- 285.8 kJ/mol

- 241.8 kJ/mol

+ 44.0 kJ/mol

Question 54

Adding an inert gas to a gaseous equilibrium system at constant volume has what effect on the equilibrium position?

A.Shifts toward the side with more moles of gas.B.Shifts toward the side with fewer moles of gas.C.No shift; the equilibrium position is unchanged.D.Shifts forward to consume the extra pressure.

Question 55

The specific heat of aluminum is $0.900 J/ (g \cdot °C)$ and of water is $4.18 J/ (g \cdot °C)$ . Equal masses of aluminum and water absorb equal amounts of heat. Which statement is correct?

A.Aluminum undergoes a greater temperature increase.B.Water undergoes a greater temperature increase.C.Both undergo the same temperature increase.D.Water absorbs more total heat because it has a higher specific heat.

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Question 56

For the reaction $A + B \to C$ , tripling $[A]$ while keeping $[B]$ constant triples the reaction rate. Tripling $[B]$ while keeping $[A]$ constant has no effect on the rate. What is the rate law?

rate = k [A] [B]

rate = k [A]

rate = k [A]^{2}

rate = k [B]

Question 57

Which statement correctly describes how a catalyst increases reaction rate?

A.A catalyst increases the temperature of the reaction, giving molecules more kinetic energy.B.A catalyst is consumed in the reaction, producing more product.C.A catalyst provides an alternative pathway with a lower activation energy, increasing the fraction of successful collisions.D.A catalyst shifts the equilibrium toward products by raising

Δ G

Question 58

A PES spectrum shows two peaks for an element: a very high binding-energy peak with relative area 1, and a low binding-energy peak with relative area 2. Which element is this?

A.HeliumB.LithiumC.BerylliumD.Boron

Question 59

The standard Gibbs free energy is related to the equilibrium constant by $Δ G^{\circ} = - R T ln K$ . If $Δ G^{\circ} = 0$ , what is $K$ ?

K = 0

K = 1

K = \infty

K = e

Question 60

When a small piece of sodium metal is added to water, bubbles form vigorously and the solution becomes basic. What type of reaction is occurring?

A.Double-displacement (metathesis) reactionB.Combustion reactionC.Single-displacement (redox) reactionD.Acid–base neutralization

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