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Chapter 6: Problem 139

Construct a table with the headings \(q, w, \Delta U,\) and \(\Delta H .\) For each of the following processes, deduce whether each of the quantities listed is positive \((+)\) negative \((-),\) or zero \((0) .\) (a) Freezing of benzene. (b) Compression of an ideal gas at constant temperature. (c) Reaction of sodium with water. (d) Boiling liquid ammonia. (e) Heating a gas at constant volume. (f) Melting of ice.

Short Answer

Expert verified
The values for each process are: a) Freezing of benzene: q=-, w=0, ΔU=0, ΔH=-. b) Compression of an ideal gas at constant temperature: q=0, w=-, ΔU=0, ΔH=0. c) Reaction of sodium with water: q=+, w=0, ΔU=+, ΔH=+. d) Boiling liquid ammonia: q=+, w=0, ΔU=+, ΔH=+. e) Heating a gas at constant volume: q=+, w=0, ΔU=+, ΔH=+. f) Melting of ice: q=+, w=0, ΔU=+, ΔH=+.

Step by step solution

01

a) Freezing of benzene

For the freezing of benzene: q (heat) is negative, because heat is released during the process. w (work done) is zero, as there is no mention of any volume change or any external work being done on the system. ΔU (change in internal energy) is zero, because for any phase transition at constant temperature and pressure, the change in internal energy is zero. ΔH (change in enthalpy) is also negative, because heat is released in this exothermic process.
02

b) Compression of an ideal gas at constant temperature

For the compression of an ideal gas at constant temperature: q is zero, because the temperature is constant. w is negative, because work is done on the system (the gas is compressed). ΔU is zero, because there is no change in temperature. And, ΔH is also zero because at constant temperature, the enthalpy change is zero.
03

c) Reaction of sodium with water

For the reaction of sodium with water: q is positive, because heat is generated during the reaction. w is zero, as there is no mention of any volume change. ΔU is positive, because the internal energy increases with the release of heat. And, ΔH is also positive, because this is an exothermic reaction and heat is released.
04

d) Boiling liquid ammonia

For boiling of liquid ammonia: q is positive, as heat is absorbed. w is zero, as there is no mention of any volume change. ΔU is positive, as there is an increase in internal energy due to the temperature increase. ΔH is also positive, as heat is gained in this endothermic process.
05

e) Heating a gas at constant volume

For heating a gas at constant volume: q is positive, because heat is added. w is zero, because there is no volume change (thus no work is done). ΔU is positive, because the internal energy increases with the addition of heat. ΔH is also positive, because heat is gained in this endothermic process.
06

f) Melting of ice

For melting of ice: q is positive, because heat is absorbed during melting. w is zero, because there is no change in volume. ΔU is positive, because the internal energy increases with the temperature increase. ΔH is also positive, as the process is endothermic with heat getting absorbed.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Enthalpy Change
Enthalpy change, represented by \( \Delta H \), plays a crucial role in thermochemistry as it denotes the heat exchanged in a chemical reaction at constant pressure. Imagine boiling water for a cup of tea; the water absorbs heat, causing an endothermic reaction with a positive enthalpy change. Conversely, when you watch a candle burn, it releases heat, exemplifying an exothermic reaction where the enthalpy change is negative. During phase transitions like the freezing of benzene or melting of ice, heat is either released or absorbed without changing temperature, implying a change in enthalpy.
Internal Energy
When considering internal energy (\( \Delta U \)), we're talking about the total energy contained within a system, including kinetic and potential energies at the molecular level. In the realm of small-scale classroom experiments, it's tough to detect changes in internal energy. However, when a chemical reaction occurs and heat is exchanged, such as in the reaction of sodium with water, the energy within the system changes – the molecules move more energetically, and thus, the internal energy increases. Similarly, when gas is heated at constant volume, its internal energy rises due to the addition of heat.
Phase Transition
Phase transitions are like magic acts of matter, where substances transform from one state to another – solid, liquid, or gas. Think of melting ice cubes in your drink during a hot day; it transitions from solid to liquid, absorbing heat in the process. This absorption or release of heat during phase transitions like the freezing and melting mentioned in the textbook problem does not change the temperature of the substance, but it does involve a change in enthalpy and internal energy.
Exothermic and Endothermic Processes
These two types of processes are all about the direction of heat flow in a chemical reaction. Exothermic processes, like an impressive fireworks display, release heat, making the surroundings warmer. The enthalpy change is negative here because the system is losing heat. On the flip side, when you’re cold and cozy up under a heat lamp, the lamp's endothermic process absorbs energy from the surroundings to produce the warming effect, resulting in a positive enthalpy change. Whether heat is added or released affects both the internal energy of the system and also dictates whether the process is exothermic or endothermic.

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Most popular questions from this chapter

Explain the meaning of this thermochemical equation: $$\begin{array}{r}4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g) \\\\\Delta H=-904 \mathrm{~kJ} / \mathrm{mol}\end{array}$$

For which of the following reactions does \(\Delta H_{\mathrm{rxn}}^{\circ}=\) \(\Delta H_{\mathrm{f}}^{\circ} ?\) (a) \(\mathrm{H}_{2}(g)+\mathrm{S}(\) rhombic \() \longrightarrow \mathrm{H}_{2} \mathrm{~S}(g)\) (b) C(diamond) \(+\mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)\) (c) \(\mathrm{H}_{2}(g)+\mathrm{CuO}(s) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{Cu}(s)\) (d) \(\mathrm{O}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{3}(g)\)

A piece of silver of mass \(362 \mathrm{~g}\) has a heat capacity of \(85.7 \mathrm{~J} /{ }^{\circ} \mathrm{C}\). What is the specific heat of silver?

Why is it dangerous to add water to a concentrated acid such as sulfuric acid in a dilution process?

The combustion of \(0.4196 \mathrm{~g}\) of a hydrocarbon releases \(17.55 \mathrm{~kJ}\) of heat. The masses of the products are \(\mathrm{CO}_{2}=1.419 \mathrm{~g}\) and \(\mathrm{H}_{2} \mathrm{O}=0.290 \mathrm{~g}\). (a) What is the empirical formula of the compound? (b) If the approximate molar mass of the compound is \(76 \mathrm{~g}\), calculate its standard enthalpy of formation.
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