Reactions in which a substance decomposes by losing CO are called decarbonylation reactions. The decarbonylation of acetic acid proceeds according to: CH₃COOH(l) → CH₃OH(g) + CO(g) By using data from Appendix C, calculate the minimum temperature at which this process will be spontaneous under standard conditions. Assume that ΔH° and ΔS° do not vary with temperature.

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Identify the reaction: CH_3COOH(l) → CH_3OH(g) + CO(g).

Use the Gibbs free energy equation: ΔG° = ΔH° - TΔS°.

For spontaneity, set ΔG° < 0, which implies ΔH° < TΔS°.

Calculate ΔH° for the reaction using standard enthalpies of formation: ΔH° = ΣΔH°_f(products) - ΣΔH°_f(reactants).

Calculate ΔS° for the reaction using standard entropies: ΔS° = ΣS°(products) - ΣS°(reactants).

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Gibbs Free Energy

Gibbs Free Energy (ΔG) is a thermodynamic potential that helps predict the spontaneity of a reaction at constant temperature and pressure. A reaction is spontaneous when ΔG is negative, which can be calculated using the equation ΔG = ΔH - TΔS, where ΔH is the change in enthalpy, T is the temperature in Kelvin, and ΔS is the change in entropy.

Enthalpy and Entropy

Enthalpy (ΔH) is a measure of the total heat content of a system, while entropy (ΔS) quantifies the degree of disorder or randomness in a system. In the context of decarbonylation reactions, understanding how these two thermodynamic properties change during the reaction is crucial for determining the conditions under which the reaction becomes spontaneous.

Standard Conditions

Standard conditions refer to a set of specific conditions used as a reference point in thermodynamics, typically defined as 1 bar of pressure and a specified temperature, often 25°C (298 K). When calculating thermodynamic properties like ΔH and ΔS, it is essential to ensure that the values used are consistent with these standard conditions to accurately assess the spontaneity of the reaction.

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Textbook Question

Consider the following reaction between oxides of nitrogen: NO₂(g) + N₂O(g) → 3 NO(g) (b) Calculate ΔG at 800 K, assuming that ΔH° and ΔS° do not change with temperature. Under standard conditions is the reaction spontaneous at 800 K?

Textbook Question

Consider the following reaction between oxides of nitrogen: NO₂(g) + N₂O(g) → 3 NO(g) (c) Calculate ΔG at 1000 K. Is the reaction spontaneous under standard conditions at this temperature?

Textbook Question

Consider the following reaction between oxides of nitrogen: NO₂(g) + N₂O(g) → 3 NO(g) (a) Use data in Appendix C to predict how ΔG for the reaction varies with increasing temperature.

Textbook Question

A certain constant-pressure reaction is barely nonspontaneous at 45 °C. The entropy change for the reaction is 72 J/K. Estimate ΔH.