Predict the effect of decreasing the temperature on the amount(s) of reactant(s) in the following reactions: (a) \(\mathrm{C}_{2} \mathrm{H}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightleftharpoons \mathrm{CH}_{3} \mathrm{CHO}(g) \quad \Delta H_{\mathrm{rxn}}^{\circ}=-151 \mathrm{~kJ}\) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}(l)+\mathrm{O}_{2}(g) \rightleftharpoons \mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(l)+\mathrm{H}_{2} \mathrm{O}(g)\) $$ \Delta H_{\mathrm{rxn}}^{\circ}=-451 \mathrm{~kJ} $$ (c) \(2 \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{CH}_{3} \mathrm{CHO}(g)\) (exothermic) (d) \(\mathrm{N}_{2} \mathrm{O}_{4}(g) \rightleftharpoons 2 \mathrm{NO}_{2}(g)\) (endothermic)

One of the critical concepts in chemistry is the idea of an equilibrium shift. This occurs when a system in equilibrium responds to a change in temperature, pressure, or concentration of reactants or products. According to Le Chatelier's Principle, if such a change is introduced, the system will adjust to oppose the effect of the change.
For instance, if the concentration of a reactant is increased, the system shifts to produce more products to balance the added reactant. Conversely, if the temperature is changed, the system will shift in a direction that counteracts the temperature change.
To understand this better, consider a simple reversible reaction:
\(\text{A + B} \rightleftharpoons \text{C + D} \)
If the temperature of this exothermic reaction decreases, the equilibrium will shift to the right, favoring the production of C and D. By shifting the balance, the system compensates for the lost heat, illustrating an equilibrium shift.

An exothermic reaction is a chemical reaction that releases heat into its surroundings. In such reactions, the energy of the products is lower than the energy of the reactants. This release of energy leads to a negative change in enthalpy (\( \Delta H \)) for the reaction.
For example, consider the reaction:
\(\text{C}_{2}\text{H}_{2}(g) + \text{H}_{2}\text{O}(g) \rightleftharpoons \text{CH}_{3}\text{CHO}(g) \) \( \Delta H = -151 kJ \)
The negative \( \Delta H \) indicates that this reaction is exothermic. Decreasing the temperature of such a reaction will shift the equilibrium to the right, increasing the production of products and reducing reactants. This shift happens because the system tries to counteract the loss in temperature by producing more heat.

In contrast to exothermic reactions, endothermic reactions absorb heat from their surroundings. The energy needed to form the products is higher than the energy of the reactants, resulting in a positive change in enthalpy (\( \Delta H \)).
For instance, consider the reaction:
\(\text{N}_{2}\text{O}_{4}(g) \rightleftharpoons 2 \text{NO}_{2}(g) \) (endothermic)
Since this reaction requires heat, decreasing the temperature will shift the equilibrium to the left, increasing the concentration of \( \text{N}_{2}\text{O}_{4} \) and decreasing the concentration of \( \text{NO}_{2} \).
This shift happens because the system compensates for the reduced temperature by favoring the reaction that absorbs heat. Thus, understanding endothermic reactions is essential for predicting how temperature changes impact equilibrium states in chemical reactions.