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Ch.6 - Thermochemistry
All textbooksTro 4th EditionCh.6 - ThermochemistryProblem 101a
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Chapter 6, Problem 101a

Use standard enthalpies of formation to calculate the standard change in enthalpy for the melting of ice. (The ΔH°f for H2O(s) is –291.8 kJ/mol.)

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Key Concepts

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

Enthalpy of Formation

The standard enthalpy of formation (ΔH°f) is the change in enthalpy when one mole of a compound is formed from its elements in their standard states. It provides a reference point for calculating the enthalpy changes in chemical reactions, allowing chemists to predict the energy changes associated with the formation or transformation of substances.
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Enthalpy of Formation

Phase Changes and Enthalpy

Phase changes, such as melting, involve energy changes that can be quantified using enthalpy. The enthalpy change for melting (ΔH_melting) is the energy required to convert a solid into a liquid at its melting point. Understanding this concept is crucial for calculating the overall enthalpy change when ice melts into water.
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Entropy in Phase Changes

Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps, regardless of the pathway taken. This principle allows for the calculation of enthalpy changes in complex reactions by using known enthalpy values, such as the standard enthalpy of formation, to derive the enthalpy change for processes like melting ice.
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Related Practice
Open Question
A 100-W lightbulb is placed in a cylinder equipped with a moveable piston. The lightbulb is turned on for 0.015 hour, and the assembly expands from an initial volume of 0.85 L to a final volume of 5.88 L against an external pressure of 1.0 atm. Use the wattage of the lightbulb and the time it is on to calculate ΔE in joules (assume that the cylinder and lightbulb assembly is the system and assume two significant figures). Calculate w. Calculate q.
Open Question
Evaporating sweat cools the body because evaporation is an endothermic process: H2O(l) → H2O(g) ΔH°rxn = +44.01 kJ. Estimate the mass of water that must evaporate from the skin to cool the body by 0.50°C. Assume a body mass of 95 kg and assume that the specific heat capacity of the body is 4.0 J/g°C.
Textbook Question

LP gas burns according to the exothermic reaction: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) ΔH°rxn = –2044 kJ What mass of LP gas is necessary to heat 1.5 L of water from room temperature (25.0 °C) to boiling (100.0 °C)? Assume that during heating, 15% of the heat emitted by the LP gas combustion goes to heat the water. The rest is lost as heat to the surroundings.

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Textbook Question

Use standard enthalpies of formation to calculate the standard change in enthalpy for the melting of ice. (The ΔH°f for H2O(s) is –291.8 kJ/mol.) Use this value to calculate the mass of ice required to cool 355 mL of a beverage from room temperature (25.0 °C) to 0.0 °C. Assume that the specific heat capacity and density of the beverage are the same as those of water.

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Textbook Question

Dry ice is solid carbon dioxide. Instead of melting, solid carbon dioxide sublimes according to the equation: CO2(s) → CO2(g) ◀ When carbon dioxide sublimes, the gaseous CO2 is cold enough to cause water vapor in the air to condense, forming fog. When dry ice is added to warm water, heat from the water causes the dry ice to sublime more quickly. The evaporating carbon dioxide produces a dense fog often used to create special effects. In a simple dry ice fog machine, dry ice is added to warm water in a Styrofoam cooler. The dry ice produces fog until it evaporates away, or until the water gets too cold to sublime the dry ice quickly enough. Suppose that a small Styrofoam cooler holds 15.0 L of water heated to 85 °C. Use standard enthalpies of formation to calculate the change in enthalpy for dry ice sublimation, and calculate the mass of dry ice that should be added to the water so that the dry ice completely sublimes away when the water reaches 25 °C. Assume no heat loss to the surroundings. (The ΔH°f for CO2(s) is –427.4 kJ/mol.)

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Textbook Question

A 25.5-g aluminum block is warmed to 65.4 °C and plunged into an insulated beaker containing 55.2 g water initially at 22.2 °C. The aluminum and the water are allowed to come to thermal equilibrium. Assuming that no heat is lost, what is the final temperature of the water and aluminum?

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