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AP Chemistry Full-length practice exam 2

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Full-length practice exam 2

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Question 1

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 2

The common-ion effect predicts that the molar solubility of AgCl ( $K_{s p} = 1.8 \times 1 0^{- 10}$ ) in a $0.10 M$ NaCl solution is:

A.Greater than

1.3 \times 1 0^{- 5} M

1.3 \times 1 0^{- 5} M

(unchanged)C.

1.8 \times 1 0^{- 9} M

1.8 \times 1 0^{- 10} M

Question 3

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 4

A student reacts $10.0 g$ of $CaC O_{3}$ with excess HCl and collects $C O_{2}$ gas. The molar mass of $CaC O_{3}$ is $100.1 g/mol$ . How many moles of $C O_{2}$ are produced?

$CaC O_{3} (s) + 2 HCl (aq) \to CaC l_{2} (aq) + H_{2} O (l) + C O_{2} (g)$

0.100 mol

0.200 mol

1.00 mol

0.0500 mol

Question 5

For the reaction $2 N O_{2} (g) \to 2 NO (g) + O_{2} (g)$ , experimental data shows that doubling $[N O_{2}]$ increases the initial rate by a factor of 4. What is the rate law and overall reaction order?

rate = k [N O_{2}]

; first order overallB.

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

; second order overallC.

rate = k [N O_{2}]^{2} [NO] [O_{2}]

; fourth order overallD.

rate = k [N O_{2}]^{1/2}

; half order overall

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Question 6

An element has two naturally occurring isotopes: one with mass 10.013 u (19.9% abundance) and one with mass 11.009 u (80.1% abundance). What is the weighted average atomic mass?

10.51 u

10.81 u

11.01 u

10.21 u

Question 7

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 8

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

Question 9

Which electron configuration correctly represents a ground-state sulfur ( $Z = 16$ ) atom?

[Ne] 3 s^{2} 3 p^{4}

[Ne] 3 s^{2} 3 p^{6}

[Ne] 3 s^{2} 3 p^{2}

[Ne] 3 s^{1} 3 p^{5}

Question 10

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

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Question 11

A sample of ideal gas has a volume of $10.0 L$ at $27° C$ . At constant pressure, to what temperature (in $° C$ ) must the gas be cooled to reduce its volume to $7.50 L$ ?

- 48° C

0° C

20.25° C

- 75° C

Question 12

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 13

A $25.0 mL$ sample of $0.100 M$ HCl is titrated with $0.100 M$ NaOH. What is the pH at the equivalence point?

1.0

7.0

13.0

4.74

Question 14

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 15

For the reaction $PC l_{5} (g) ⇌ PC l_{3} (g) + C l_{2} (g)$ , $K_{c} = 1.8 \times 1 0^{- 2}$ at a certain temperature. If a reaction mixture has $[PC l_{5}] = 0.50 M$ , $[PC l_{3}] = 0.15 M$ , and $[C l_{2}] = 0.15 M$ , in which direction will the reaction proceed?

A.Forward, toward products, because

Q < K_{c}

.B.Reverse, toward reactants, because

Q > K_{c}

.C.Neither; the system is at equilibrium because all concentrations are nonzero.D.Forward, because the products have lower concentrations than the reactant.

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Question 16

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 17

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 18

Rank the following in order of increasing bond length: C–C single bond, C=C double bond, $C \equiv C$ triple bond.

C \equiv C < C = C < C - C

C - C < C = C < C \equiv C

C = C < C - C < C \equiv C

D.All are equal because the atoms are the same.

Question 19

Which of the following best describes the orientation and energy requirements for an effective collision in collision theory?

A.Any collision produces a reaction; increasing collision frequency always increases rate.B.Only collisions with energy

\geq E_{a}

AND correct geometric orientation lead to reaction.C.Orientation does not matter; only the kinetic energy of the collision is important.D.Energy is irrelevant; rate depends only on the frequency of correctly oriented collisions.

Question 20

At $0° C$ and $1.00 atm$ , $1 mol$ of a real gas occupies a slightly smaller volume than predicted by the ideal gas law. What does this suggest about the dominant deviation at this condition?

A.Molecular volume is dominant, pushing measured volume above the ideal.B.Intermolecular attractions are dominant, pulling molecules together and reducing volume below the ideal.C.The gas molecules are moving faster than predicted by the kinetic molecular theory.D.The gas is undergoing a phase change at this condition.

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Question 21

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 22

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 23

A $100.0 mL$ sample of $0.100 M$ acetic acid ( $p K_{a} = 4.74$ ) is titrated with $0.100 M$ NaOH. After adding $50.0 mL$ of NaOH, what is the pH?

7.00

4.74

2.87

8.72

Question 24

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 25

How many lone pairs of electrons are on the central nitrogen atom in $N H_{3}$ ?

0

1

2

3

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Question 26

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

Question 27

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 28

Using the following standard enthalpies of formation, calculate $Δ H^{\circ}$ for $C H_{4} (g) + 2 O_{2} (g) \to C O_{2} (g) + 2 H_{2} O (l)$ :

$Δ H_{f}^{\circ} [C H_{4} (g)] = - 74.8 kJ/mol$

$Δ H_{f}^{\circ} [C O_{2} (g)] = - 393.5 kJ/mol$

$Δ H_{f}^{\circ} [H_{2} O (l)] = - 285.8 kJ/mol$

- 604.5 kJ/mol

- 890.3 kJ/mol

+ 890.3 kJ/mol

- 965.1 kJ/mol

Question 29

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

Question 30

For a reaction with $K_{c} = 5.0 \times 1 0^{- 8}$ , which statement best characterizes the equilibrium?

A.At equilibrium, the mixture contains mostly products.B.At equilibrium, reactants and products are present in roughly equal amounts.C.At equilibrium, the mixture is predominantly reactants.D.The reaction does not reach equilibrium because

K_{c}

is too small.

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Question 31

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 32

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 33

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.

Question 34

Which of the following correctly ranks the species by increasing atomic/ionic radius?

$C l^{-}, Cl, N a^{+}, Na$

N a^{+} < C l^{-} < Na < Cl

N a^{+} < Na < C l^{-} < Cl

N a^{+} < Cl < Na < C l^{-}

Cl < N a^{+} < Na < C l^{-}

Question 35

Which of the following molecules is polar?

B F_{3}

CC l_{4}

C O_{2}

H_{2} S

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Question 36

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 37

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

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

Question 38

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 39

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 40

What is the pH of a $0.010 M$ solution of $HCl$ ?

1.0

2.0

3.0

12.0

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Question 41

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 42

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 43

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 44

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 45

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

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Question 46

Calculate the standard cell potential $E_{cell}^{\circ}$ for the reaction:

$Zn (s) + C u^{2 +} (aq) \to Z n^{2 +} (aq) + Cu (s)$

Given: $E_{red}^{\circ} (C u^{2 +} /Cu) = + 0.34 V$ ; $E_{red}^{\circ} (Z n^{2 +} /Zn) = - 0.76 V$

- 1.10 V

+ 1.10 V

+ 0.42 V

+ 0.34 V

Question 47

Explain why $HF$ has a much higher boiling point ( $19.5° C$ ) than $HCl$ ( $- 85° C$ ), even though HCl has the larger molar mass.

A.HF has stronger London dispersion forces because fluorine is highly electronegative.B.HF molecules engage in hydrogen bonding, a much stronger intermolecular attraction than the dipole–dipole forces in HCl.C.The H–F bond is covalent while H–Cl is ionic.D.HF has a higher molar mass than HCl.

Question 48

A reaction at constant pressure releases $150 kJ$ of heat and does $30 kJ$ of work on the surroundings. What is $Δ H$ for this process?

- 120 kJ

- 180 kJ

- 150 kJ

+ 150 kJ

Question 49

Which of the following is the correct net ionic equation for the reaction between aqueous $Pb (N O_{3})_{2}$ and aqueous $KI$ , which forms a yellow precipitate of $Pb I_{2}$ ?

Pb (N O_{3})_{2} (aq) + 2 KI (aq) \to Pb I_{2} (s) + 2 KN O_{3} (aq)

P b^{2 +} (aq) + 2 I^{-} (aq) \to Pb I_{2} (s)

K^{+} (aq) + N O_{3}^{-} (aq) \to KN O_{3} (aq)

P b^{2 +} (aq) + I^{-} (aq) \to Pb I_{2} (s)

Question 50

For the equilibrium $A (g) ⇌ 2 B (g)$ , an ICE table gives the following initial and change rows:

	A	B
Initial	$0.600 M$	$0$
Change	$- x$	$+ 2 x$
Equilibrium	$0.600 - x$	$2 x$

If $K_{c} = 0.040$ , what is the equilibrium concentration of A?

0.550 M

0.600 M

0.200 M

0.0500 M

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Question 51

The average bond enthalpy of an O=O double bond is $498 kJ/mol$ and that of an O–H bond is $464 kJ/mol$ . Using bond enthalpies, estimate $Δ H$ for:

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

( $D (H - H) = 436 kJ/mol$ )

- 478 kJ

- 484 kJ

+ 484 kJ

- 928 kJ

Question 52

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 53

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 54

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)

Question 55

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)

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Question 56

In a mass spectrometry experiment, a sample of neon gas produces peaks at $m / z = 20, 21,$ and $22$ . What do the different $m / z$ values represent?

A.Different ions of neon with varying chargesB.Neon atoms that lost different numbers of electronsC.Isotopes of neon with different numbers of neutronsD.Impurities in the neon sample

Question 57

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

- 2

- 1

0

+ 1

Question 58

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

Question 59

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 60

When $HF$ ionizes in water, fluoride ion is formed. Why is $F^{-}$ a stronger base than $C l^{-}$ ?

A.Fluorine has a higher electronegativity, which makes the F–H bond stronger and F

^{-}

more reluctant to give it up.B.

F^{-}

is smaller and holds its charge density over a smaller volume, making it a better proton acceptor (stronger base).C.

F^{-}

has a higher negative charge than

C l^{-}

.D.F

^{-}

has more lone pairs available for proton donation.

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