Thanks for joining me again for the newest installment of my crash course series on Chemistry!
In case you missed it, last time I wrote on Spontaneity, Entropy, & Free Energy, and today’s lesson will be on chemical equilibrium!
From last time, we know that
If we take the antilogarithm, we get
From this equation we can deduce that
- Reactions with the largest negative value of ΔG will tend to proceed to completion to the greatest extent.
- Conversely, if ΔG>0, K will be less than one, and although some products will be present at equilibrium, most of the material will be in the form of reactants.
- When ΔG=0, this corresponds to an equilibrium constant of one, where neither reactant or product is favored.
As previously discussed, the equation
relates the free energy change to the standard free energy and the pressure/concentrations of the reagents.
Q is the reaction quotient, defined as
for the reaction aA + bB -> cC + dD.
If ΔGr° < 0, the reaction is product favored.
If ΔGr° > 0, the reaction is reactant favored.
According to the following equation, the larger the value of ΔSr°, the larger K is. The tendency toward maximum entropy (“chaos”) for the system directly influences the magnitude of the equilibrium constant.
A negative ΔHr° corresponds to a larger K as well. The enthalpy part of the equation is really another way of describing the entropy of the surroundings. A negative ΔHr° corresponds to a positive ΔSsurr°.
Chemical reactions usually don’t go to completion, but to an equilibrium state. Equilibrium is a state in which there are no observable changes as time goes by. The concentration of reactants and products remain constant over time. However, this does not mean that equilibrium is a static state – the backwards rate is equal to the forward rate, which is known as dynamic equilibrium.
Systems move towards equilibrium state spontaneously – it proceeds at some finite rate without the action of outside influences. If the system is disturbed, it will eventually return to equilibrium state.
According to the law of mass action, a certain of reactant and product concentrations for a reversible reaction at equilibrium has a constant value K, the equilibrium constant.
For the equation
aA + bB <=> cC + dD
Each set of equilibrium concentrations/partial pressures is called an equilibrium position.
There is only one equilibrium position for a particular system at a particular temperature, but there are infinite numbers of possible equilibrium positions (infinite number of initial conditions)!
If K>>1, the reaction is product favored, and if K<<1, the reaction is reactant favored.
In thermodynamics, K is defined such that it has no units, because every concentration or pressure is actually a ratio to a standard value (1 M and 1 bar, respectively).
This eliminates the unit without altering the numerical part of the concentration or pressure.
In a homogeneous equilibrium all reactants and products are in the same phase.
For a gas phase reaction, assuming ideal gas behavior, the equilibrium constant is expressed in terms of pressure.
Also, this can be expressed in terms of concentration, because (c) of an ideal gas is directly related to pressure (P) at constant temperature.
In an aqueous reaction, solvents do not appear in equilibrium constant expressions, because the concentration of the solvent is a large quantity in comparison to other species in the solution, and can be treated as a constant.
Reactants and products involved in a reversible reaction are in different phases in a heterogeneous equilibrium.
If a pure liquid or a solid is a reactant or product, we can treat its activity (concentration) as a constant and omit it from the equilibrium constant expression.
Predicting the Direction of a Reaction
For reactions that have not reached equilibrium, we can compare the value of the reaction quotient (Q) and the equilibrium constant (K) to determine the direction in which the net reaction will proceed to achieve equilibrium.
- If Q<K ~ system proceeds from left to right
- If Q=K ~ system is at equilibrium
- If Q>K ~ system proceeds from right to left
Example Calculations for Equilibrium Concentrations/Partial Pressures
- Express the equilibrium concentration/partial pressures of all species in terms of the initial concentrations/partial pressures and a single unknown x, representing the change in concentration/partial pressure.
- Write the equilibrium constant expression in terms of the equilibrium concentrations/partial pressures. Knowing the value of the equilibrium constant, solve for x.
- Having solved for x, calculated the equilibrium concentrations/partial pressures of all species.
Example (from Chemistry LibreTexts):
Factors that Affect Chemical Equilibrium
This analysis will be based on the Le Chatelier’s Principle: if an external stress is applied to a system at equilibrium, the system adjusts in a way that the stress is partially offset.
Concentration: Increasing the concentrations of the reactants or equivalently increasing the partial pressure of the reactants shifts the equilibrium position to the right.
Volume and Pressure: Increase in total pressure (decrease in total volume) favors the net reaction that decreases the total number of moles of gases, and a decrease of pressure (decrease in volume) favors the net reaction that increases the total number of moles of gases. If the reaction does not affect the number of moles of gases, a pressure/volume change has no effect on the equilibrium position.
Temperature: Only a change in temperature can alter the equilibrium constant. A temperature increase favors an endothermic reaction (heat can be considered to be a reactant), and a temperature decrease favors an exothermic reaction (heat can be considered to be a product).
Catalyst: The presence of a catalyst neither alters the equilibrium constant or shifts the position of the equilibrium system. It will increase both the forwards and backwards rate equally so that the mixture will reach equilibrium sooner by reducing the activation energy and increasing the rate constant.
Isochoric isothermal addition of an inert gas: If an inert gas is added at constant volume and temperature, the partial pressure of each gas taking part in the reaction is unaffected by the addition of the inert gas. There will be no shift in the ideal gas equilibrium and there will be no change in the equilibrium constant.
Thanks for reading! I hope that this was helpful to you. If you have any further questions, feel free to leave a comment below, and I will be sure to address them. If you want to keep learning more about chemistry, subscribe to get all my newest posts!