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Ch.18 - Thermodynamics: Entropy, Free Energy & Equilibrium
McMurry - Chemistry 8th Edition
McMurry8th EditionChemistryISBN: 9781292336145Not the one you use?Change textbook
All textbooksMcMurry 8th EditionCh.18 - Thermodynamics: Entropy, Free Energy & EquilibriumProblem 139
Chapter 18, Problem 139

Chloroform has ΔHvaporization = 29.2 kJ>mol and boils at 61.2 °C. What is the value of ΔSvaporization for chloroform?

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1
Identify the given values: ΔHvaporization = 29.2 kJ/mol, boiling point = 61.2 °C.
Convert the boiling point from Celsius to Kelvin by adding 273.15 to the Celsius temperature. T(K) = 61.2 °C + 273.15.
Use the formula for entropy change during vaporization, ΔSvaporization = ΔHvaporization / T, where T is the temperature in Kelvin.
Substitute the values of ΔHvaporization and T into the formula to find ΔSvaporization.
Ensure the units of ΔSvaporization are in J/(mol·K) by converting kJ to J if necessary.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Enthalpy of Vaporization (ΔHvaporization)

The enthalpy of vaporization is the amount of energy required to convert a unit mass of a liquid into vapor without a change in temperature. It is typically expressed in kJ/mol and is a crucial factor in understanding phase changes, particularly when calculating the energy changes associated with boiling or evaporating substances.
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Enthalpy of Vaporization Example

Entropy of Vaporization (ΔSvaporization)

The entropy of vaporization is a measure of the disorder or randomness that occurs when a liquid transitions to a gas. It is calculated using the relationship ΔSvaporization = ΔHvaporization / T, where T is the temperature in Kelvin. This concept is essential for understanding the thermodynamic favorability of phase changes.

Temperature Conversion

In thermodynamics, it is important to use the absolute temperature scale (Kelvin) when performing calculations involving energy changes. To convert Celsius to Kelvin, add 273.15 to the Celsius temperature. This conversion is necessary for accurately applying formulas that involve temperature, such as those relating to enthalpy and entropy.
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Related Practice
Textbook Question

Consider the unbalanced equation: (b) Use the data in Appendix B and ΔG°f for IO3-(aq)= -128.0 kJ/mol to calculate ΔG° for the reaction at 25 °C.

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Textbook Question

Consider the unbalanced equation: I2(s) → I-(aq) + IO3-(aq) (d) What pH is required for the reaction to be at equilibrium at 25°C when [I-] = 0.10M and [IO3-] = 0.50 M?

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Textbook Question

The lead storage battery uses the reaction: (b) Calculate ∆G for this reaction on a cold winter's day (10 °F) in a battery that has run down to the point where the sulfuric acid concentration is only 0.100 M.

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Textbook Question
A mixture of 14.0 g of N2 and 3.024 g of H2 in a 5.00 L container is heated to 400 °C. Use the data in Appendix B to calculate the molar concentrations of N2, H2, and NH3 at equilibrium. Assume that ∆H° and ∆S° are independent of temperature, and remember that the standard state of a gas is defined in terms of pressure.

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Textbook Question

Methanol (CH3OH) is made industrially in two steps from CO and H2. It is so cheap to make that it is being considered for use as a precursor to hydrocarbon fuels, such as methane (CH4):

Step 1. CO(g) + 2 H2(g) S CH3OH(l) ΔS° = - 332 J/K

Step 2. CH3OH1l2 → CH4(g) + 1/2 O2(g) ΔS° = 162 J/K

(k) Calculate an overall ΔG°, ΔH°, and ΔS° for the formation of CH4 from CO and H2.

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Textbook Question
Suppose that a reaction has ΔH = - 33 kJ and ΔS = - 58 J>K. At what temperature will it change from spontaneous to nonspontaneous?
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