Problem 111 - Chemistry: An Atoms-Focused Approach

Question

The industrial production of hydrogen chloride gas is most frequently carried out by direct synthesis from hydrogen and chlorine: $$\mathrm{H}_{2}(g)+\mathrm{Cl}_{2}(g) \rightarrow 2 \mathrm{HCl}(g)$$ Smaller quantities of HCl \((g)\) may be produced on the laboratory scale by the reaction of sodium chloride and sulfuric acid: $$2 \mathrm{NaCl}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \rightarrow 2 \mathrm{HCl}(g)+\mathrm{Na}_{2} \mathrm{SO}_{4}(s)$$ Apply concepts discussed in this chapter and data from the appendix to determine if either heating or cooling is required when these reactions are carried out.

Step-by-Step Solution

Answer

Answer: The first reaction is exothermic and does not require heating or cooling. The second reaction is endothermic and requires heating to proceed.

Step 1: Identify the components

From the given exercise, we have two reactions: 1. \(\mathrm{H}_{2}(g) + \mathrm{Cl}_{2}(g) \rightarrow 2 \mathrm{HCl}(g)\) 2. \(2 \mathrm{NaCl}(s) + \mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \rightarrow 2 \mathrm{HCl}(g) + \mathrm{Na}_{2} \mathrm{SO}_{4}(s)\)

Step 2: Calculate the enthalpy change (ΔH) for each reaction

We'll use the following formula to find the enthalpy change for each reaction: $$\Delta H = \sum \Delta H_{f} \: (products) - \sum \Delta H_{f} \: (reactants)$$ The standard enthalpy of formation values (\(\Delta H_{f}\)) for the substances involved in the reactions can be found in the appendix or reference data tables. We'll use these values to calculate the enthalpy change, which indicates if the reaction is endothermic or exothermic.

Step 3: Calculate ΔH for Reaction 1

For the first reaction, \(\mathrm{H}_{2}(g) + \mathrm{Cl}_{2}(g) \rightarrow 2 \mathrm{HCl}(g)\), we have the following enthalpy of formation values: - \(\mathrm{H}_{2}(g)\): 0 kJ/mol (standard state) - \(\mathrm{Cl}_{2}(g)\): 0 kJ/mol (standard state) - \(\mathrm{HCl}(g)\): -92.3 kJ/mol Using these values, let's calculate the enthalpy change: $$\Delta H_1 = [2(-92.3)] - [0 + 0] = -184.6 \: \mathrm{kJ/mol}$$ Since \(\Delta H_1\) is negative, the first reaction is exothermic, which means it releases heat.

Step 4: Calculate ΔH for Reaction 2

For the second reaction, \(2 \mathrm{NaCl}(s) + \mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \rightarrow 2 \mathrm{HCl}(g) + \mathrm{Na}_{2} \mathrm{SO}_{4}(s)\), we have the following enthalpy of formation values: - \(\mathrm{NaCl}(s)\): -411.2 kJ/mol - \(\mathrm{H}_{2} \mathrm{SO}_{4}(\ell)\): -814.1 kJ/mol - \(\mathrm{HCl}(g)\): -92.3 kJ/mol - \(\mathrm{Na}_{2} \mathrm{SO}_{4}(s)\): -1387.5 kJ/mol Using these values, let's calculate the enthalpy change: $$\Delta H_2 = [2(-92.3) + (-1387.5)] - [2(-411.2) + (-814.1)] = 24.7 \: \mathrm{kJ/mol}$$ Since \(\Delta H_2\) is positive, the second reaction is endothermic, which means it absorbs heat.

Step 5: Determine if heating or cooling is required

Based on our calculations, we can conclude the following: - For Reaction 1, it's exothermic, which means it releases heat. Therefore, no heating or cooling is required for this reaction. - For Reaction 2, it's endothermic, which means it absorbs heat. Therefore, heating is required for this reaction to proceed.

Key Concepts

Exothermic Reaction

An exothermic reaction is a type of chemical reaction that releases energy, usually in the form of heat, to its surroundings. This energy release is a result of the formation of new bonds, which requires less energy than was used to break the initial bonds in the reactants. As a result, the products of the reaction are usually at a lower energy state than the reactants.

A classic example of an exothermic reaction is the production of hydrogen chloride gas from hydrogen and chlorine, as shown in the reaction: \[\mathrm{H}_{2}(g) + \mathrm{Cl}_{2}(g) \rightarrow 2 \mathrm{HCl}(g)\]Here, the enthalpy change (\(\Delta H\)) is negative (-184.6 kJ/mol), indicating that the reaction gives off heat.

• Releases heat to surroundings
• Products are at lower energy state than reactants
• Often marked by a temperature rise in the environment

Endothermic Reaction

In contrast to exothermic reactions, endothermic reactions absorb energy from their surroundings. This absorption is necessary to drive the reaction forward, often requiring heat to be added. These reactions are characterized by a positive enthalpy change (\(\Delta H\)), where more energy is needed to break the bonds of the reactants compared to what is released when new bonds are formed in the products.

The reaction of sodium chloride with sulfuric acid to produce hydrogen chloride and sodium sulfate is endothermic:\[2 \mathrm{NaCl}(s) + \mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \rightarrow 2 \mathrm{HCl}(g) + \mathrm{Na}_{2} \mathrm{SO}_{4}(s)\]Here, the enthalpy change (\(\Delta H\)) is +24.7 kJ/mol, meaning heat is absorbed.

• Absorbs heat from surroundings
• Reactants are at lower energy state than products
• Often requires heating to proceed

Chemical Synthesis

Chemical synthesis is the process of creating complex chemical compounds from simpler precursors. This is fundamental in industrial chemistry and research, as it allows for the production of materials and substances necessary for everyday life and scientific advancement.

For instance, synthesizing hydrogen chloride gas can involve a direct reaction between elements like hydrogen and chlorine or through other methods such as the reaction involving sodium chloride and sulfuric acid. Each synthetic route requires a detailed understanding of the reaction conditions, energy requirements, and chemical properties.

• Combines simpler chemicals to form complex compounds
• Applicable in industry, laboratories, and everyday products
• Requires careful consideration of reaction conditions

Thermochemistry

Thermochemistry is a branch of chemistry that deals with the energy and heat associated with chemical reactions and physical transformations. It allows scientists to predict reaction tendencies, energy changes, and reaction feasibility.

Key concepts in thermochemistry include understanding enthalpy changes like \(\Delta H\), which indicates whether a reaction is exothermic or endothermic. This knowledge is crucial in both laboratory-scale and industrial chemical processes, where managing energy exchange is essential for efficiency, safety, and cost-effectiveness.

• Studies energy changes in chemical reactions
• Essential for understanding reaction dynamics
• Applies to both endothermic and exothermic processes