Question

$$ \begin{aligned} &\text { Use these bond enthalpy values to answer Question } { . }\\\ &\begin{array}{lclc} \hline \text { Bond } & \begin{array}{c} \text { Bond Enthalpy } \\ (\mathrm{k}\rfloor / \mathrm{mol}) \end{array} & \text { Bond } & \begin{array}{c} \text { Bond Enthalpy } \\ (\mathrm{k} / / \mathrm{mol}) \end{array} \\ \hline \mathrm{H}-\mathrm{F} & 566 & \mathrm{~F}-\mathrm{F} & 158 \\ \mathrm{H}-\mathrm{Cl} & 431 & \mathrm{Cl}-\mathrm{Cl} & 242 \\ \mathrm{H}-\mathrm{Br} & 366 & \mathrm{Br}-\mathrm{Br} & 193 \\ \mathrm{H}-\mathrm{I} & 299 & \mathrm{I}-\mathrm{I} & 151 \\ \mathrm{H}-\mathrm{H} & 436 & & \\ \hline \end{array} \end{aligned} $$ For the reactions of molecular hydrogen with fluorine and with chlorine: (a) Calculate the enthalpy change for breaking all the bonds in the reactants. (b) Calculate the enthalpy change for forming all the bonds in the products. (c) From the results in parts (a) and (b), calculate the enthalpy change for the reaction. (d) Which reaction is most exothermic?

Short answer

The reaction of \( \text{H}_2 \) with \( \text{F}_2 \) is more exothermic with \( \Delta H = -538 \, \text{kJ/mol} \).

Step-by-step solution

Understand the Reactions

The reactions we are analyzing are: \( \text{H}_2 + \text{F}_2 \rightarrow 2\text{HF} \) and \( \text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl} \). For each reaction, identify the bonds broken in the reactants and bonds formed in the products.

Calculate Bonds Broken in Reactants

For \( \text{H}_2 + \text{F}_2 \), bonds broken are one \( \text{H-H} \) and one \( \text{F-F} \). The energy required is: \[ 436 + 158 = 594 \, \text{kJ/mol} \]For \( \text{H}_2 + \text{Cl}_2 \), bonds broken are one \( \text{H-H} \) and one \( \text{Cl-Cl} \). The energy required is: \[ 436 + 242 = 678 \, \text{kJ/mol} \]

Calculate Bonds Formed in Products

For \( \text{H}_2 + \text{F}_2 \), bonds formed are two \( \text{H-F} \). The energy released is:\[ 2 \times 566 = 1132 \, \text{kJ/mol} \]For \( \text{H}_2 + \text{Cl}_2 \), bonds formed are two \( \text{H-Cl} \). The energy released is: \[ 2 \times 431 = 862 \, \text{kJ/mol} \]

Calculate Enthalpy Change for Each Reaction

Use the formula: \( \Delta H = \text{Energy Required to Break Bonds} - \text{Energy Released Forming Bonds} \).For \( \text{H}_2 + \text{F}_2 \):\[ \Delta H = 594 - 1132 = -538 \, \text{kJ/mol} \]For \( \text{H}_2 + \text{Cl}_2 \):\[ \Delta H = 678 - 862 = -184 \, \text{kJ/mol} \]

Determine the More Exothermic Reaction

The reaction with the more negative enthalpy change is more exothermic. In this case, \( \text{H}_2 + \text{F}_2 \) with \( \Delta H = -538 \, \text{kJ/mol} \) is more exothermic than \( \text{H}_2 + \text{Cl}_2 \).

Key concepts

Enthalpy Change

Enthalpy change is a vital concept in understanding chemical reactions. It represents the heat energy change within a system at constant pressure when a chemical reaction occurs. This change can be calculated by determining the difference in enthalpy between the reactants and the products.
In the context of chemical reactions, we often express enthalpy change using the Greek letter delta (Δ), followed by the symbol H: \[\Delta H = \text{Enthalpy of Products} - \text{Enthalpy of Reactants} \]This value tells us whether the reaction absorbs or releases energy. Calculating enthalpy changes helps in predicting the energy requirements and outcomes of chemical processes, which is crucial for scientific and industrial applications.

Exothermic Reaction

Exothermic reactions are chemical reactions that release energy to their surroundings, usually in the form of heat or light. When the enthalpy change \( \Delta H \) is negative, it indicates that the reaction is exothermic. This means that the energy required to break the bonds in the reactants is less than the energy released when new bonds form in the products.
Exothermic reactions are common; they occur in processes like combustion, respiration, and many industrial applications. During an exothermic reaction, the temperature of the surroundings usually increases. People often experience this increase in warmth when touching an object where an exothermic reaction is taking place.

Chemical Reactions

Chemical reactions occur when substances, known as reactants, transform into different substances, called products. These transformations involve the rearrangement of atoms and are facilitated by various factors like energy, temperature, and catalysts. Understanding chemical reactions is essential, as they are the foundation of all chemical studies and industrial chemical processes around us.
Every chemical reaction involves an energy change, which is calculated through bond enthalpies. The sum of the bond energies of reactants subtracted by the sum of bond energies of the products gives us the enthalpy change for the reaction, allowing insights into whether a reaction is endothermic or exothermic.

Bond Formation and Breaking

Bond formation and breaking are the core of chemical reactions. When bonds between atoms in reactants break, it requires energy. Meanwhile, forming new bonds between atoms in the products releases energy. The balance between these two processes determines the overall energy change of the reaction.

• Bond Breaking: This is an endothermic process. Energy is absorbed to break the bonds within the reactants.
• Bond Formation: This is exothermic. Energy is released when new bonds are formed in the products.

The key to understanding reactions lies in comparing the energies of these two processes. Calculating the energies helps predict whether a reaction releases energy (exothermic) or consumes energy (endothermic). This knowledge is not only theoretical but also intensely practical in laboratories and industries.