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Chapter 5: Problem 3

Identify whether the following processes are exothermic or endothermic. Is the sign on \(q_{\text {sys }}\) positive or negative? (a) combustion of methane (b) melting of ice (c) raising the temperature of water from \(25^{\circ} \mathrm{C}\) to \(100^{\circ} \mathrm{C}\) (d) heating \(\mathrm{CaCO}_{3}(\mathrm{s})\) to form \(\mathrm{CaO}(\mathrm{s})\) and \(\mathrm{CO}_{2}(\mathrm{g})\)

Short Answer

Expert verified
(a) Exothermic, negative \( q_{\text{sys}} \); (b) Endothermic, positive; (c) Endothermic, positive; (d) Endothermic, positive.

Step by step solution

01

Analyze the Combustion of Methane

Combustion of methane is a chemical reaction where methane reacts with oxygen to produce carbon dioxide and water. This process releases energy, usually in the form of heat and light. Therefore, it is exothermic. The sign on \( q_{\text{sys}} \) is negative because heat is released to the surroundings.
02

Analyze the Melting of Ice

Melting of ice is a phase change from solid to liquid. Energy is required to break the intermolecular forces holding the ice together, so the process absorbs heat from the surroundings. This makes it endothermic. The sign on \( q_{\text{sys}} \) is positive because heat is absorbed by the system.
03

Analyze Raising the Temperature of Water

Raising the temperature of water involves adding heat to increase the kinetic energy of the water molecules. Thus, the process is endothermic as it requires energy input. The sign on \( q_{\text{sys}} \) is positive because heat is absorbed by the water.
04

Analyze the Heating of \( \text{CaCO}_3 \)

Heating \( \text{CaCO}_3 \) to form \( \text{CaO} \) and \( \text{CO}_2 \) involves breaking the chemical bonds in calcium carbonate, which requires energy. Therefore, the reaction is endothermic. The sign on \( q_{\text{sys}} \) is positive as heat is absorbed during this endothermic reaction.

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Key Concepts

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

Combustion of Methane
When methane burns in the presence of oxygen, it undergoes combustion, producing carbon dioxide and water. This reaction is exothermic, meaning it releases energy into the surroundings. This energy often manifests as heat and light, which is why we can see and feel fire. In terms of thermodynamics, the system loses heat, so the sign of the heat energy (\( q_{\text{sys}} \)) is negative.
  • Exothermic Reaction: Energy is released.
  • Sign of \( q_{\text{sys}} \): Negative because heat exits the system.
Melting of Ice
The melting of ice is the transition from solid to liquid state. This phase change absorbs energy, needed to break the bonds between the water molecules. The process is endothermic; energy is absorbed from the surroundings to accomplish this change. As energy is absorbed, the sign of the heat energy (\( q_{\text{sys}} \)) becomes positive.
  • Endothermic Process: Energy is absorbed.
  • Sign of \( q_{\text{sys}} \): Positive because heat enters the system.
Temperature Change in Water
Raising the temperature of water from \(25^{\circ} \text{C}\) to \(100^{\circ} \text{C}\) involves increasing the water's kinetic energy. To achieve this rise in temperature, heat must be added to the system. This process is endothermic, as it requires energy input. As heat is absorbed by the water, the sign of the heat energy (\( q_{\text{sys}} \)) is positive.
  • Endothermic Heating: Energy is input into the system.
  • Sign of \( q_{\text{sys}} \): Positive due to heat absorption.
Thermal Decomposition of Calcium Carbonate
Heating calcium carbonate (\( \text{CaCO}_3 \)) results in its decomposition into calcium oxide (\( \text{CaO} \)) and carbon dioxide (\( \text{CO}_2 \)). This decomposition requires breaking chemical bonds, necessitating an input of energy. Thus, this reaction is endothermic. The absorption of heat in this process means the sign of the heat energy (\( q_{\text{sys}} \)) is positive.
  • Endothermic Reaction: Energy is needed to break bonds.
  • Sign of \( q_{\text{sys}} \): Positive indicating heat uptake.

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Most popular questions from this chapter

Sublimation of \(1.0 \mathrm{g}\) of dry ice, \(\mathrm{CO}_{2}(\mathrm{s}),\) forms \(0.36 \mathrm{L}\) of \(\mathrm{CO}_{2}(\mathrm{g})\left(\mathrm{at}-78^{\circ} \mathrm{C} \text { and } 1.01 \times 10^{5} \mathrm{Pa}\right)\) The expanding gas can do work on the surroundings (Figure 5.8 ). Calculate the amount of work done on the surroundings.

Heat, Work, and Internal Energy As a gas cools, it is compressed from 2.50 L to 1.25 L under a constant pressure of \(1.01 \times 10^{5}\) Pa. Calculate the work (in J) required to compress the gas.

In the lab, you plan to carry out a calorimetry experiment to determine \(\Delta_{r} H\) for the exothermic reaction of \(\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{s})\) and \(\mathrm{HCl}(\mathrm{aq}) .\) Predict how each of the following will affect the calculated value of \(\Delta_{t} H .\) (The value calculated for \(\Delta_{i} H\) for this reaction is a negative value so choose your answer from the following: \(\Delta, H\) will be too low [that is, a larger negative valuel, \(\Delta_{r} H\) will be unaffected, \(\Delta_{r} H\) will be too high [that is, a smaller negative value].) (a) You spill a little bit of the \(\mathrm{Ca}(\mathrm{OH})_{2}\) on the benchtop before adding it to the calorimeter. (b) Because of a miscalculation, you add an excess of HCl to the measured amount of \(\mathrm{Ca}(\mathrm{OH})_{2}\) in the calorimeter. (c) \(\mathrm{Ca}(\mathrm{OH})_{2}\) readily absorbs water from the air. The \(\mathrm{Ca}(\mathrm{OH})_{2}\) sample you weighed had been exposed to the air prior to weighing and had absorbed some water. (d) After weighing out \(\mathrm{Ca}(\mathrm{OH})_{2},\) the sample sat in an open beaker and absorbed water. (e) You delay too long in recording the final temperature. (f) The insulation in your coffee-cup calorimeter was poor, so some energy as heat was lost to the surroundings during the experiment. (g) You have ignored the fact that energy as heat also raised the temperature of the stirrer and the thermometer in your system.

Determine whether energy as heat is evolved or required, and whether work was done on the system or whether the system does work on the surroundings, in the following processes at constant pressure: (a) Liquid water at \(100^{\circ} \mathrm{C}\) is converted to steam at \(100^{\circ} \mathrm{C}\) (b) Dry ice, \(\mathrm{CO}_{2}(\mathrm{s}),\) sublimes to give \(\mathrm{CO}_{2}(\mathrm{g})\)

The enthalpy change for the oxidation of naphthalene, \(\mathrm{C}_{10} \mathrm{H}_{8},\) is measured by calorimetry. $$ \begin{aligned} \mathrm{C}_{10} \mathrm{H}_{g}(\mathrm{s})+12 \mathrm{O}_{2}(\mathrm{g}) \rightarrow & 10 \mathrm{CO}_{2}(\mathrm{g})+4 \mathrm{H}_{2} \mathrm{O}(\ell) \\\ \Delta_{i} H^{\circ} &=-5156.1 \mathrm{kJ} / \mathrm{mol}-\mathrm{rxn} \end{aligned} $$ Use this value, along with the standard enthalpies of formation of \(\mathrm{CO}_{2}(\mathrm{g})\) and \(\mathrm{H}_{2} \mathrm{O}(\ell),\) to calculate the enthalpy of formation of naphthalene, in kJ/mol.
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