Consider the reaction: A solution is made containing an initial [Fe3+] of 1.0 * 10-3 M and an initial [SCN-] of 8.0 * 10-4 M. At equilibrium, [FeSCN2+] = 1.7 * 10-4 M. Calculate the value of the equilibrium constant (Kc).

In chemistry, the concept of chemical equilibrium is paramount to understanding how reactions reach a state of balance. Unlike in physical equilibrium where things might seem static, in chemical equilibrium, both the forward and reverse reactions are ongoing processes. However, they occur at the same rate, leading to a constant concentration of reactants and products over time.

Applying this to our exercise, once the reaction between Fe3+ and SCN- to form FeSCN2+ reaches chemical equilibrium, the rate of formation of FeSCN2+ equals the rate at which it breaks back down into Fe3+ and SCN-. It's here where the reaction seems to 'pause', not because it stopped, but due to the opposing reactions perfectly compensating each other.

The reaction quotient, or Q, plays an instrumental role as a mathematical predictor of the direction of a chemical reaction. It has the same form as the equilibrium constant expression but uses initial concentrations instead of equilibrium concentrations.

In essence, Q analyzes the current status of a reaction: If Q < Kc, the reaction will proceed forward to produce more products. Conversely, if Q > Kc, the reaction will shift to form more reactants. When Q = Kc, the system is at equilibrium. For the purpose of our problem, we deal with the equilibrium constant directly because we're provided with equilibrium concentrations.

Stoichiometry is the heart of chemical reactions. It is the study of the quantitative relationships, or ratios, within a chemical reaction, as defined by the balanced equation. Stoichiometry allows us to predict how much product we can produce from given amounts of reactants or vice versa.

In the given problem, stoichiometry dictates that for every mole of FeSCN2+ formed, one mole each of Fe3+ and SCN- is used. This 1:1:1 stoichiometric ratio is essential in calculating how the concentrations of the reactants change as they form the product, reaffirming the changes we see as the system moves toward equilibrium.

Equilibrium concentration is the amount of reactants and products present in a reaction mixture when the reaction has reached equilibrium. Knowing the starting concentrations and the change caused by the reaction, we can use stoichiometry to find the final, or equilibrium, concentrations of all species involved in the reaction.

In the exercise, we determine these values by subtracting the concentration of FeSCN2+ formed from the initial concentrations of Fe3+ and SCN-. The equilibrium concentrations are crucial for calculating the equilibrium constant. They reflect the precise relationship and ratio of reactants to products that stabilizes at chemical equilibrium.