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Chemical equilibrium is a dynamic state in which the rate of the forward reaction equals the rate of the reverse reaction. At this point, there is no net change in the concentrations of reactants and products. When we deal with a chemical equation like \[ \mathrm{H}_{2}(\mathrm{g}) + \mathrm{I}_{2}(\mathrm{g}) \rightleftarrows 2 \mathrm{HI}(\mathrm{g}) \]it reaches equilibrium when the concentrations of \( \mathrm{H_2} \), \( \mathrm{I_2} \), and \( \mathrm{HI} \) remain constant over time. This does not mean the reactions have stopped but that the forward and backward reactions occur at the same rate.
An essential feature of equilibrium is that it can be disturbed by changing the conditions in a process known as Le Chatelier's Principle. If you change the concentration of reactants or products, or alter the temperature or pressure, the system will adjust to counteract that change and reestablish equilibrium. Understanding these principles helps in predicting the direction in which a reaction will proceed when conditions are changed.

Reaction stoichiometry involves the quantitative relationship between reactants and products in a chemical reaction. In the given reaction equation, \( \mathrm{H}_{2}(\mathrm{g}) + \mathrm{I}_{2}(\mathrm{g}) \rightleftarrows 2 \mathrm{HI}(\mathrm{g}) \), the stoichiometry shows that one mole of \( \mathrm{H}_2 \) reacts with one mole of \( \mathrm{I}_2 \) to produce two moles of \( \mathrm{HI} \). This 1:1:2 stoichiometric ratio is crucial for calculating changes in concentrations as the reaction proceeds.
In our specific case, because 78.6% of \( \mathrm{I}_2 \) is consumed, an equal amount of \( \mathrm{H}_2 \) must also be consumed. Since two moles of \( \mathrm{HI} \) are produced for every mole of \( \mathrm{I}_2 \) consumed, the amount of \( \mathrm{HI} \) formed can be calculated by doubling the change in \( \mathrm{I}_2 \). This exact ratio is vital for balancing chemical equations and performing calculations related to the yield of products.

Concentration changes play a central role in understanding chemical reactions at equilibrium. The initial concentrations of \( \mathrm{H}_{2} \) and \( \mathrm{I}_{2} \) were both \( 0.0088 \ \text{mol/L} \). With 78.6% of each being used, the concentration change, denoted \( \Delta c \), was \( 0.00693 \ \text{mol/L} \).
Using those changes, we determine equilibrium concentrations by subtracting \( \Delta c \) from the initial concentrations of \( \mathrm{H}_2 \) and \( \mathrm{I}_2 \), resulting in \( 0.00187 \ \text{mol/L} \) for each. The concentration of \( \mathrm{HI} \) increased by \( 2 \times 0.00693 = 0.01386 \ \text{mol/L} \).
Understanding how to calculate these changes is important for applying the equilibrium constant expression \[K = \frac{[\mathrm{HI}]^2}{[\mathrm{H_2}][\mathrm{I_2}]} \]. This equation links concentration changes to the equilibrium constant \( K \), providing insight into the extent of reaction and stability of the substances at equilibrium.