Consider the following reactions at equilibrium and determine which of the indicated changes will cause the reaction to proceed to the right. (1) \(\mathrm{CO}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons \mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g)\left(\right.\) add \(\left.\mathrm{CH}_{4}\right)\) (2) \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)\) (remove \(\mathrm{NH}_{3}\) ) (3) \(\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \rightleftharpoons 2 \mathrm{HF}(g)\) (add \(\mathrm{F}_{2}\) ) (4) \(\mathrm{BaO}(s)+\mathrm{SO}_{3}(g) \rightleftharpoons \mathrm{BaSO}_{4}(s)\) (add BaO) (a) \(2.3\) (b) \(1.4\) (c) \(2.4\) (d) \(2,3,4\)

Chemical equilibrium is a state in a chemical reaction where the rates of the forward and reverse reactions are equal, resulting in no overall change in the concentrations of reactants and products over time. This dynamic balance doesn't mean the reactants and products stop reacting; they do, but at a rate that doesn't change the overall amounts.

In the context of the exercise, looking at the provided chemical reactions, when they reach equilibrium, the ratio of the concentrations of products to reactants remains constant provided the system is closed and conditions remain constant. This ratio is described by the equilibrium constant (K). The process of reaching equilibrium can be affected by varying different parameters such as concentration, pressure, volume, or temperature. Understanding this balance is crucial for chemists to predict how a system will respond to changes in conditions.

The concept of reaction shifts is integral to predicting how a chemical system at equilibrium will respond to changes in conditions, a manifestation of Le Chatelier's principle. This principle states that if an external change is applied to a system at equilibrium, the system will adjust or shift in such a way as to counteract the change and establish a new equilibrium.

As seen in our exercise, adding or removing a substance can induce such a shift. If more product like CH4 is added, as in the first reaction, the equilibrium shifts to the left to decrease the product concentration, counteracting the change. In contrast, removing a product like NH3 in the second reaction shifts the equilibrium to the right, increasing the production of NH3 to balance out its removal. Recognizing these shifts helps understand and predict the direction in which reactions will proceed when disturbed.

Changes in the concentration of reactants or products can significantly influence the position of equilibrium in a chemical reaction. This is the basis for understanding how alterations in concentration can lead to reaction shifts, directly applying Le Chatelier's principle.

In the context of our exercise, when F2 is added to the third reaction, the increase in reactant concentration drives the system to produce more product (HF) to reduce the F2 concentration, thus shifting the equilibrium to the right. Similarly, removing a product causes the equilibrium to shift in the direction that tends to increase the concentration of the removed substance. However, it's important to note as exemplified in the fourth reaction, that changing the concentration of a solid, like BaO, does not alter the equilibrium position in a heterogeneous reaction as it does not appear in the expression for the equilibrium constant.