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Chapter 1: Problem 28

Campers carry hot packs with them, especially when camping during the winter months or at high altitudes. The design is similar to that described in Problem 27, except that calcium chloride is used in place of the ammonium nitrate. The equation for the dissolution of calcium chloride in water is shown below. How does the hot pack work? $$ \mathrm{CaCl}_{2}(s) \stackrel{\mathrm{H}_{2} \mathrm{O}}{\longrightarrow} \mathrm{Ca}^{2+}(a q)+2 \mathrm{Cl}^{-}(a q) \quad \Delta H=-81 \mathrm{~kJ} \cdot \mathrm{mol}^{-1} $$

Short Answer

Expert verified
The hot pack works by releasing heat due to the exothermic dissolution of calcium chloride in water.

Step by step solution

01

Understand the Chemical Equation

The provided equation shows the dissolution of solid calcium chloride (\( \mathrm{CaCl}_2(s) \)) into water, where it dissociates into calcium ions (\( \mathrm{Ca}^{2+}(aq) \)) and chloride ions (\( 2\mathrm{Cl}^-(aq) \)).
02

Analyze the Enthalpy Change

The equation includes \( \Delta H = -81 \, \mathrm{kJ/mol} \), which indicates the reaction is exothermic. In an exothermic process, energy is released as heat when the reaction occurs.
03

Relate to the Function of a Hot Pack

Since the dissolution of calcium chloride is exothermic (as indicated by the negative \( \Delta H \)), the process releases heat when \( \mathrm{CaCl}_2 \) dissolves in water. This heat release increases the temperature of the surroundings, which is the primary function of a hot pack.

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Key Concepts

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

Understanding Calcium Chloride
Calcium Chloride, or \( \mathrm{CaCl}_2 \), is a common salt with diverse applications. It is a white, crystalline solid that dissolves readily in water. When dissolved, it forms calcium ions \( \mathrm{Ca}^{2+} \) and chloride ions \( \mathrm{Cl}^- \). This compound is utilized in many fields:
  • De-icing roads and sidewalks due to its ability to depress the freezing point of water.
  • As a dust suppressant on construction sites to reduce airborne particles.
  • Within the food industry as a firming agent for canned vegetables.
  • In hot packs to generate heat through an exothermic reaction.
This versatile use is due to the chemical properties of calcium chloride when it interacts with water, which is essential in situations where warmth or specific chemical reactions are needed.
Exploring the Dissolution Process
The dissolution process of calcium chloride involves the breaking of ionic bonds in the solid state and the formation of ions in solution. When you add \( \mathrm{CaCl}_2 \) to water, it dissociates into two types of ions:
  • Calcium ions \( \mathrm{Ca}^{2+} \)
  • Chloride ions \( \mathrm{Cl}^- \)
Once dissolved, these ions are surrounded by water molecules, a process known as hydration. During hydration, water molecules, which are polar, surround and stabilize the ions. This energy-intensive process results in the release of heat. The efficient dissolution of \( \mathrm{CaCl}_2 \) makes it effective in generating heat, essential for hot packs, by utilizing the energy released when the crystalline solid dissolves in water.
The Role of Enthalpy Change in Exothermic Reactions
Enthalpy change, denoted as \( \Delta H \), measures the heat absorbed or released during a chemical reaction. In the case of an exothermic reaction involving calcium chloride, \( \Delta H \) is negative, specifically \( -81 \, \mathrm{kJ/mol} \). This indicates that the reaction releases energy to its surroundings.Exothermic reactions like this are characterized by:
  • A release of heat, increasing the temperature of the surroundings.
  • A negative \( \Delta H \) value, signifying the release of thermal energy.
  • Practical applications such as heating devices or hot packs.
Understanding \( \Delta H \) is crucial because it directly explains why certain chemical processes raise the temperature of nearby environments鈥攎aking this property useful, for instance, in keeping warm during cold weather.

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Most popular questions from this chapter

A monoclonal antibody binds to the protein cytochrome \(c\). The \(\Delta H\) value for binding at \(25^{\circ} \mathrm{C}\) is \(-87.9 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}\) and the \(\Delta S\) is \(-118 \mathrm{~J} \cdot \mathrm{K}^{-1} \cdot \mathrm{mol}^{-1}\). a. Does entropy increase or decrease when the antibody binds to the protein? b. Calculate \(\Delta G\) for the formation of the antibody-protein complex. Does the complex form spontaneously? c. The \(\Delta G\) value for the binding of a second monoclonal antibody to cytochrome \(c\) is \(-58.2 \mathrm{~kJ} \cdot \mathrm{mol}^{-1}\). Which antibody binds more readily to the protein?

Urea \(\left(\mathrm{NH}_{2} \mathrm{CONH}_{2}\right)\) dissolves readily in water; i.e., this is a spontaneous process. The beaker containing the dissolved compound is cold to the touch. What conclusions can you make about the sign of the a. enthalpy change and b. entropy change for this process?

The nutritive quality of food can be analyzed by measuring the amounts of the chemical elements it contains. Most foods are mixtures of the three major types of molecules: a. fats (lipids), b. carbohydrates, and c. proteins. What elements are present in each of these types of molecules?

Which of the following processes are spontaneous? a. A reaction that occurs with any size decrease in enthalpy and any size increase in entropy. b. A reaction that occurs with a small increase in enthalpy and a large increase in entropy. c. A reaction that occurs with a large decrease in enthalpy and a small decrease in entropy. d. A reaction that occurs with any size increase in enthalpy and any size decrease in entropy.

For a given reaction, the value of \(\Delta H\) is \(15 \mathrm{~kJ}^{-1} \mathrm{~mol}^{-1}\) and the value of \(\Delta S\) is \(51 \mathrm{~J} \cdot \mathrm{K}^{-1} \cdot \mathrm{mol}^{-1}\). Above what temperature will this reaction be spontaneous?
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