This is an application of Le Chatlier's Principle. When you take away heat from the reaction, the reaction shifts toward the left in order to compensate from the heat loss. The reaction may then be rewritten to include energy as a reactant.

Since the energy is on the reactant side, the reaction is endothermic.

According to Le Chatelier's principle, when you compress a system, its volume decreases, so partial pressure of the all the gases in the system increases. The system will act to try to decrease the pressure by decreasing the moles of gas.

Recall that ion-product constant of water, , is at and .

An endothermic reaction signifies that heat is at the reactant side. By the LeChatelier's principle, increased heat to shifts the equilibrium to the right side, favoring the increase of and . This means that and both increase, decreasing pH and pOH to less than 7, each. As a result, pH + pOH is less than 14.

Recall that ion-product constant of water, , is at and .

An endothermic reaction signifies that heat is at the reactant side. By the LeChatelier's principle, increased heat to shifts the equilibrium to the right side, favoring the increase of and . This means that and both increase, decreasing pH and pOH to less than 7, each. As a result, pH + pOH is less than 14.

By increasing the concentration of one of the reactants, the reaction will compensate by shifting to the right to increase production of products.

Increasing the concentration of one of the products (such as increasing \[C\]), however, would have the opposite effect. Increasing the temperature of an exothermic reaction would shift the reaction to the left, while increasing the temperature of an endothermic reaction would lead to a rightward shift. Finally, decreasing the volume leads to an increase in partial pressure of each gas, which the system compensates for by shifting to the side with fewer moles of gas. In this case, the right side has three moles of gas, while the left side has two; thus decreasing volume would shift equilibrium to the left.

According to Le Chatelier's principle, when you compress a system, its volume decreases, so partial pressure of the all the gases in the system increases. The system will act to try to decrease the pressure by decreasing the moles of gas.

This is an application of Le Chatlier's Principle. When you take away heat from the reaction, the reaction shifts toward the left in order to compensate from the heat loss. The reaction may then be rewritten to include energy as a reactant.

Since the energy is on the reactant side, the reaction is endothermic.

By increasing the concentration of one of the reactants, the reaction will compensate by shifting to the right to increase production of products.

Increasing the concentration of one of the products (such as increasing \[C\]), however, would have the opposite effect. Increasing the temperature of an exothermic reaction would shift the reaction to the left, while increasing the temperature of an endothermic reaction would lead to a rightward shift. Finally, decreasing the volume leads to an increase in partial pressure of each gas, which the system compensates for by shifting to the side with fewer moles of gas. In this case, the right side has three moles of gas, while the left side has two; thus decreasing volume would shift equilibrium to the left.

When writing an equilibrium constant expression, remember that products are on the top, and reactants are on the bottom of the expression. The coefficients for the compounds in the balanced reaction become the exponents for the compounds seen in the expression.

When Q = K, the reaction has reached equilibrium and is in a state of dynamic equilibrium, where the forward and reverse reactions are occurring at the same rate. Since there is no net change in concentration, the free energy is zero.