# Worksheet: Spontaneous Reactions

In this worksheet, we will practice using the second law of thermodynamics to predict the spontaneity of physical processes and chemical reactions.

Q1:

The combustion of benzene produces carbon dioxide and water as the only products.

Write a balanced chemical equation for this reaction.

• A
• B
• C
• D
• E

Predict the signs of the enthalpy change and entropy change for this reaction.

• A is negative and the sign of is zero.
• B is positive and is negative.
• CBoth and are positive.
• D is negative and is positive.
• EBoth and are negative.

Why is heating necessary for this reaction to take place?

• ATo produce a positive value of
• BTo produce a negative value of
• CTo overcome the activation energy
• DTo shift the position of equilibrium toward products
• ETo vaporize the liquid reactant

Q2:

Which of the following statements is false?

• AIn a spontaneous process, the enthalpy of the system must decrease.
• BIn a spontaneous process, the entropy of the universe must increase.
• CA spontaneous process may become nonspontaneous above a threshold temperature.
• DIn a spontaneous process, the Gibbs free energy of the system must decrease.
• EA spontaneous process occurs without a continual input of energy from an external source.

Q3:

Which of the following determines whether molecules react spontaneously?

• AFree energy change being greater in magnitude than activation energy
• BProduct-free energy being more negative than reactant-free energy
• CEntropy change being positive
• DEnthalpy change being negative
• EKinetic energy being greater than activation energy

Q4:

Why is the conversion of diamond to graphite not observed at room temperature and pressure?

• AThe activation energy for the process is too high.
• BThe melting point of diamond is too high.
• CThe process involves a decrease in entropy.
• DThe process is nonspontaneous.
• EThe process is endothermic.

Q5:

Which of the following processes are spontaneous?

1. Water freezing above its freezing point
2. Combustion of gasoline
3. Iron rusting in a humid atmosphere
• A1 and 3
• B1 and 2
• C3 only
• D2 and 3
• E2 only

Q6:

Consider the following reaction for the decomposition of sodium bicarbonate. The and for this reaction have values of 85.2 kJ/mol and 215 J/K⋅mol respectively. What is the minimum temperature required for an sample to spontaneously decompose into the products shown above (under 1 bar pressure conditions)?

Q7:

For the process , occurring at a temperature of , the has a value of J/K⋅mol. Which of the following would reflect the values of and associated with the freezing of water at a temperature of ?

• A,
• B,
• C,
• D,

Q8:

For a given reaction, the enthalpy and entropy changes are negative. The reaction is .

• Anonspontaneous at all temperatures
• Bspontaneous only at high temperatures
• Cspontaneous at all temperatures
• Dspontaneous only at low temperatures

Q9:

A given reaction is non-spontaneous at all temperatures. Which of the following best accounts for this behavior?

• A
• B
• C
• D
• E

Q10:

A given reaction is spontaneous at all temperatures. Which of the following best accounts for this behavior?

• A,
• B,
• C,
• D,
• E,

Q11:

A given reaction is spontaneous at high temperatures but not at low temperatures. Which of the following best accounts for this behavior?

• A,
• B,
• C,
• D,
• E,

Q12:

A given reaction is spontaneous at low temperatures but not at high temperatures. Which of the following best accounts for this behavior?

• A,
• B,
• C,
• D,
• E,

Q13:

For a given reaction, the enthalpy change is positive and the entropy change is negative. The reaction is .

• Aspontaneous at all temperatures
• Bspontaneous only at high temperatures
• Cspontaneous only at low temperatures
• Dnonspontaneous at all temperatures

Q14:

For a given reaction, the enthalpy change is negative and the entropy change is positive. The reaction is .

• Aspontaneous only at low temperatures
• Bnonspontaneous at all temperatures
• Cspontaneous only at high temperatures
• Dspontaneous at all temperatures

Q15:

For a given reaction, the enthalpy and entropy changes are positive. The reaction is .

• Anonspontaneous at all temperatures
• Bspontaneous only at high temperatures
• Cspontaneous only at low temperatures
• Dspontaneous at all temperatures

Q16:

What are the values of Gibbs free energy, , and the change in Gibbs free energy, , when a system reaches equilibrium?

• ABoth and are zero.
• B is maximized and is zero.
• C is zero and is minimized.
• D is zero and is maximized.
• E is minimized and is zero.

Q17:

The standard enthalpy change for a reaction, , is 100 kJ/mol. The standard entropy change for the reaction is . It may be assumed that both quantities remain constant with varying temperature.

Calculate the minimum necessary value of for the reaction to be spontaneous at 298 K.

If , calculate the minimum temperature at which the reaction is spontaneous.

Q18:

Zinc reacts with solid copper sulfate to form copper and zinc sulfate. The standard entropies of these and other materials are shown in the table.

 Material Standard Molar Entropy 𝑆⦵(J/K·mol) Cu()g Cu()s CuSO()4s Zn()s Zn()g ZnSO()4s 166.38 33.15 109.20 41.60 160.98 110.50

Calculate, to 3 significant figures, the standard entropy change for this reaction, expressed per mole of zinc.

Calculate, to 3 significant figures, the maximum value of , the standard molar enthalpy change per mole of zinc, at which this reaction would occur spontaneously at 298 K.

Q19:

The standard entropies for two phases of sodium chloride are shown below:

Phase Standard Molar Entropy 𝑆⦵ (J/K⋅mol) NaCl()s NaCl()l 72.11 95.06

The standard enthalpy of fusion of sodium chloride, , is 27.95 kJ/mol.

Assuming that these thermodynamic parameters do not vary with temperature, estimate the melting point of sodium chloride to the nearest degree Celsius.

The observed melting point of sodium chloride is 1,074 K. Which of the following is not a potential explanation for the difference between the observed and calculated melting temperatures?

• A and
• B and
• C and
• D and
• E and

Q20:

Calcium carbonate decomposes into calcium oxide and carbon dioxide gas. The standard entropies and enthalpies of formation for calcium carbonate and other materials are shown in the table.

MaterialStandard Molar Entropy (J/K⋅mol)Standard Enthalpy of Formation (kJ/mol)
110.0
184.1
38.1
83.4
197.7
213.8

All thermodynamic parameters are measured under a standard pressure of 1.000 bar. It may be assumed that the parameters do not vary significantly with temperature.

A sample of calcium carbonate is stored in an atmosphere of carbon dioxide at standard pressure. Calculate, to 3 significant figures, the minimum temperature in kelvin at which the sample would spontaneously decompose.

• A K
• B K
• C K
• D K
• E K

Calculate, to 1 significant figure, the equilibrium partial pressure of carbon dioxide when a sample of calcium carbonate is heated to .

• A bar
• B bar
• C bar
• D bar
• E bar