Exam 2 lecture 7

When a chemical reaction takes place in a closed system, the quantities of components change until an end point is reached where the composition remains undisturbed, calle equilibrium. Things are still happening, but the net change is zero. We define the reaction quotient, Q, or equilibrium constant, K_eq, using the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients. When dealing with solutions, we often assume concentration equals activity. For gases, we use partial pressures, and for solids or liquids, their activity is considered to be 1. It's important to distinguish between Delta G standard, which is the standard reaction free energy for pure reactants under standard conditions, and Delta G, which is the reaction free energy or "driving force" as the composition of the reaction mixture evolves during a reaction. When Delta G equals zero, there is no driving force, and the reaction has reached equilibrium. As the reaction proceeds, this driving force changes because the reactants or products are not present in their standard states. Thermodynamics is directly related to equilibrium through the equation Delta G = Delta G standard + RT ln Q, where R is the gas constant and T is the temperature in Kelvin. At equilibrium, Delta G is zero, which means Delta G standard = -RT ln K. If Delta G standard is greater than zero, the reaction favors reactants. If Delta G standard is less than zero, the reaction favors products. Equilibrium constants can also be very temperature dependent, meaning they change with variations in temperature. The direction of a reaction can be predicted by comparing the reaction quotient Q to the equilibrium constant K. If Q is less than K, then Delta G is less than zero, and the reaction will proceed towards products. If Q equals K, then Delta G is zero, and the reaction is at equilibrium. Finally, if Q is greater than K, then Delta G is greater than zero, and the reaction will proceed towards reactants. We can also manipulate equilibrium constants. For a reverse reaction, K' = 1/K. If we add reactions together, the overall equilibrium constant is the product of the individual equilibrium constants. For example, if A <=> B has K_1 and B <=> C has K_2, then A <=> C has K_3 = K_1 * K_2. Setting up a reaction table, also known as an ICE (Initial, Change, Equilibrium) table, is a useful method for solving equilibrium problems. For a reaction like A + 2B <=> C, if we start with initial concentrations of 0.01 M for A and 0.02 M for B, and 0 for C, we can determine the changes in concentration to reach equilibrium. If the equilibrium constant K_eq is very small, we can often make an approximation where the change 'x' is negligible compared to the initial concentrations. Le Châtelier's Principle states that if a system at equilibrium is disturbed, the direction in which the system proceeds back to equilibrium is such that the disturbance is partially offset.