For equilibrium, the net force acting on the system must be zero.

The resultant force can be found using the Pythagorean theorem since the forces are perpendicular.

The equilibrium constant Kc is dimensionless because it is a ratio of concentrations raised to their stoichiometric coefficients.

Adding SO2 will shift the equilibrium to the right to produce more SO3, according to Le Chatelier's principle.

Removing SO3 will shift the equilibrium to the right to produce more SO3.

Removing SO2 will shift the equilibrium to the left to produce more SO2, according to Le Chatelier's principle.

Decreasing the volume increases the pressure, and the equilibrium will shift towards the side with fewer moles of gas, which is the right side (2 moles of SO3).

Increasing the concentration of a reactant shifts the equilibrium position to the right, favoring the formation of products (SO3).

Increasing the pressure will shift the equilibrium towards the side with fewer moles of gas, which is the right side in this case (2 moles of SO3).

Removing SO3 will shift the equilibrium to the right to produce more SO3, according to Le Chatelier's principle.

Increasing pressure shifts the equilibrium towards the side with fewer moles of gas, which is the right side in this case.

Decreasing the temperature for an exothermic reaction will shift the equilibrium to the right, favoring the formation of CH3OH.

Removing H2 will shift the equilibrium to the left to produce more H2, thus favoring the reactants.

Removing H2 will shift the equilibrium to the left to produce more H2, according to Le Chatelier's principle.

Removing H2 will shift the equilibrium to the left to produce more H2, according to Le Chatelier's principle.

Decreasing the volume increases the pressure, which shifts the equilibrium to the side with fewer moles of gas, favoring the production of CH3OH.

Decreasing the volume increases the pressure, and the equilibrium will shift to the side with fewer moles of gas, which is the left side in this case.

Decreasing the volume increases the pressure, and the equilibrium will shift towards the side with fewer moles of gas, which is the left side in this case.

If the reaction is exothermic, increasing the temperature will shift the equilibrium to the left, favoring the reactants.

When a reaction is reversed, the equilibrium constant for the reverse reaction is the reciprocal of the equilibrium constant for the forward reaction.

When a reaction is reversed, the equilibrium constant is inverted (1/K).

[H+] = 10^(-pH) = 10^(-3) = 0.001 M

A catalyst does not affect the equilibrium constant; it only speeds up the rate at which equilibrium is reached.

A catalyst speeds up the rate of both the forward and reverse reactions equally, thus having no effect on the position of equilibrium.

A catalyst increases the rate of both the forward and reverse reactions equally, thus it does not shift the position of equilibrium.

Adding a common ion decreases the solubility of a salt in a saturated solution due to the common ion effect.

A buffer solution is designed to resist changes in pH, so adding a strong acid will cause only a small change in pH.

Adding a strong base to a buffer solution will increase the pH, but the buffer will resist significant changes in pH due to the presence of the weak acid.

Adding an inert gas at constant volume does not change the partial pressures of the reactants and products, so the equilibrium position remains unchanged.

Adding an inert gas at constant volume does not change the partial pressures of the reactants and products, thus it has no effect on the equilibrium position.