Consider the reaction: MISSED THIS? Read Section 16.8 : Watch KCV 16.8 , ME 16.9 $$ \begin{array}{r} \mathrm{NiO}(\mathrm{s})+\mathrm{CO}(g) \rightleftharpoons \mathrm{Ni}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g}) \\ K_{c}=4.0 \times 10^{3} \text { at } 1500 \mathrm{~K} \end{array} $$ If a mixture of solid nickel(II) axide and \(0.20 \mathrm{M}\) carbon monoxIde comes to equllibrium at \(1500 \mathrm{~K},\) what is the equilibrium concentration of \(\mathrm{CO}_{2} ?\)

The equilibrium constant expression relates the concentrations of reactants and products at equilibrium for a reversible chemical reaction. It's a vital tool for chemists to predict the extent of a chemical reaction and balance between the reactants and products.

For a given reversible reaction, like \[\mathrm{aA} + \mathrm{bB} \rightleftharpoons \mathrm{cC} + \mathrm{dD},\] the expression is written as \[K_c = \frac{[\mathrm{C}]^c [\mathrm{D}]^d}{[\mathrm{A}]^a [\mathrm{B}]^b}\] where \[K_c\] is the equilibrium constant in terms of concentration, and \[a, b, c,\] and \[d\] are the stoichiometric coefficients for reactants \[A\] and \[B\], and products \[C\] and \[D\], respectively. It's important to note that the concentrations of solids and pure liquids do not appear in this expression, as their concentration does not change during the reaction.

Le Chatelier's principle provides a predictive qualitative tool for understanding how a system at equilibrium responds to external changes. It states that if an external stress is applied to a system at equilibrium, the system will adjust in such a way as to partially counteract the stress and regain a new equilibrium state.

Stresses that can affect chemical equilibrium include changes in concentration, temperature, and pressure. For example, if the concentration of a reactant is increased, the system will respond by consuming more of the added reactant to produce more products. Likewise, if the temperature of an exothermic reaction is increased, the equilibrium will shift to favor the reactants, as the system seeks to decrease the temperature by absorbing heat.

To calculate the equilibrium concentrations of reactants and products, one typically starts with the initial concentrations, defines the change that occurs as the system reaches equilibrium, sets up the equilibrium constant expression, and then solves for the unknown.

In the given example, we start with an initial concentration of CO of 0.20 M and CO2 of 0 M. When the system reaches equilibrium, as defined by the expression \[ K_c = \frac{[\mathrm{CO_{2}}]}{[\mathrm{CO}]}\], we denote the change in concentration of CO2 as \[x\]. By substituting the equilibrium concentrations into the expression and solving for \[x\], the balance between the reactants' consumption and the production of the products is quantified. This systematic approach to solving for equilibrium concentrations underscores the dynamic nature of chemical equilibria and provides a quantitative measure of the composition of the system at equilibrium.

Chemical reaction equilibrium is a state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentration of reactants and products over time. However, this does not imply that the reaction has stopped; rather, reactions continue to occur in both directions at the same rate.

Reaction equilibrium is dynamic and can be influenced by changes in reaction conditions. The equilibrium constant, K, indicates the ratio of product to reactant concentrations at equilibrium and is constant for a given reaction at a specific temperature. Therefore, understanding chemical equilibrium is crucial for predicting how much product can be produced in a chemical process and for controlling industrial chemical reactions.