Question

The major industrial source of hydrogen gas is by the following reaction: $$ \mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g) $$ Use bond energies to predict r this reaction.

Short answer

The enthalpy change for the given reaction \(\mathrm{CH}_{4}(g) + \mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{CO}(g) + 3 \mathrm{H}_{2}(g)\) can be calculated using bond energies. The total bond energies for the reactants are 2590 kJ/mol, and for the products, 2107 kJ/mol. The enthalpy change for the reaction is ∆H = -483 kJ/mol, indicating an exothermic process.

Step-by-step solution

Understand the reaction and identify the bonds involved

The given reaction is: \[ \mathrm{CH}_{4}(g) + \mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{CO}(g) + 3 \mathrm{H}_{2}(g) \] In order to calculate the enthalpy change, we need to understand the bonds involved. The reactants have the following bonds: - CH₄: 4 C-H bonds - H₂O: 2 O-H bonds The products have the following bonds: - CO: 1 C=O bond (double bond) - 3H₂: 3 H-H bonds

Look up the bond energies for the bonds involved

Using a bond energies table, we can find the average bond energies for the bonds involved in the reaction: - C-H: 414 kJ/mol - O-H: 467 kJ/mol - C=O: 799 kJ/mol - H-H: 436 kJ/mol

Calculate the total bond energies for reactants and products

Now we calculate the total bond energies for both reactants and products: Reactants total bond energy: \(4 \times 414 \, \text{kJ/mol (C-H)} + 2 \times 467 \, \text{kJ/mol (O-H)} = 1656 \, \text{kJ/mol} + 934\, \text{kJ/mol} = 2590 \, \text{kJ/mol}\) Products total bond energy: \(1 \times 799 \, \text{kJ/mol (C=O)} + 3 \times 436 \, \text{kJ/mol (H-H)} = 799 \, \text{kJ/mol} + 1308 \, \text{kJ/mol} = 2107 \, \text{kJ/mol}\)

Calculate the enthalpy change for the reaction

The enthalpy change (∆H) can be calculated by subtracting the total bond energy of the reactants from the total bond energy of the products: ∆H = Total bond energy (Products) - Total bond energy (Reactants) ∆H = 2107 kJ/mol - 2590 kJ/mol = -483 kJ/mol The negative value of ∆H indicates that the reaction is exothermic, and it releases 483 kJ/mol of energy.

Key concepts

Chemical Bond Energy

Understanding chemical bond energy is fundamental to calculating the enthalpy change of a chemical reaction. Bond energy refers to the amount of energy required to break one mole of a specific chemical bond in a gaseous state. Similarly, this energy is released when the bond is formed.

Worth noting is that these values are averages because the energy may vary depending on the molecule's environment. In a practical scenario, during a chemical reaction, bonds in the reactants must break for the reactants to transform into products, which involves absorbing energy. Consequently, new bonds form among the products, releasing energy. The bond energies provided, such as for C-H bonds (414 kJ/mol), O-H bonds (467 kJ/mol), and C=O bonds (799 kJ/mol), enable us to perform these calculations.

The total energy absorbed or released in a reaction is the sum of all the energies absorbed in breaking the original bonds minus the energy released in forming new bonds. This understanding of bond energies helps students to predict whether a reaction will require external energy to proceed or release energy.

Enthalpy Change

The enthalpy change, denoted as \(\Delta H\), signifies the heat change in a reaction at constant pressure. It's an essential concept in thermodynamics, indicating whether a reaction releases energy (exothermic) or absorbs energy (endothermic). The calculation of enthalpy change involves considering the bond energies of both reactants and products, as seen in the step-by-step solution.

For example, if the total bond energy of the reactants is higher than that of the products, the reaction releases excess energy to its surroundings, and \(\Delta H\) is negative, signifying an exothermic process. On the contrary, if more energy is required to break bonds in the reactants than is released when forming new bonds in the products, \(\Delta H\) is positive, indicating an endothermic reaction.

To enhance students' understanding, it's beneficial to highlight the physical interpretation of enthalpy change. A negative \(\Delta H\) means the surroundings are getting warmer, while a positive one implies that the surroundings are cooling down due to the absorption of heat.

Exothermic Reaction

An exothermic reaction is characterized by the release of energy, usually in the form of heat, to the surroundings, resulting in a temperature increase in the surrounding environment. This type of reaction occurs when the total bond energy of the products is less than that of the reactants, leading to a negative enthalpy change \(\Delta H\).

Typical examples include combustion reactions like burning fuels, or the simple act of rust forming on iron. When students confront a problem involving exothermic reactions, they can actually feel the heat released, which makes the concept more tangible. Additionally, when the heat is released during these reactions, it often manifests as light or thermal energy, the latter of which can be leveraged for heating in industrial and domestic settings.

Understanding that an exothermic reaction, such as the one forming hydrogen gas from methane and water, can self-sustain once initiated, helps explain phenomena like the constant burning of a candle or a perpetual industrial furnace.