Decreasing the pressure will shift the equilibrium to the side with more moles of gas, which in this case is the right side (C).

Increasing pressure shifts the equilibrium towards the side with fewer moles of gas, which is the right side (1 mole of CH3OH).

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

A decrease in Kc with an increase in temperature indicates that the reaction is exothermic.

Removing H2 decreases its concentration, causing the equilibrium to shift to the left to produce more reactants and restore balance.

Decreasing the concentration of NH3 will shift the equilibrium to the right to produce more NH3, according to Le Chatelier's Principle.

Decreasing the volume increases the pressure, and according to Le Chatelier's Principle, the equilibrium will shift to the side with fewer moles of gas, which is to the right in this case.

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

Removing NH3 will shift the equilibrium to the right to produce more NH3, in accordance with Le Chatelier's Principle.

Decreasing the concentration of NH3 will shift the equilibrium to the right to produce more NH3, according to Le Chatelier's Principle.

Dilution decreases the concentration of reactants and products, and according to Le Chatelier's Principle, the equilibrium will shift to the left to produce more reactants.

For an exothermic reaction, increasing the temperature shifts the equilibrium to the left, favoring the reactants, as the system tries to absorb the added heat.

Increasing temperature generally increases the enthalpy of a substance due to increased kinetic energy.

Increasing temperature typically increases the enthalpy of a system due to the increased kinetic energy of the molecules.

For an endothermic reaction, increasing the temperature increases the equilibrium constant (K) as the reaction favors the formation of products.

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

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

Removing a product will shift the equilibrium to the right to produce more of that product, according to Le Chatelier's Principle.

Decreasing the volume increases the pressure, which shifts the equilibrium towards the side with fewer moles of gas.

Adding a product increases its concentration, causing the equilibrium to shift to the left to favor the formation of reactants and restore balance.

According to the Arrhenius equation, an increase in temperature generally increases the rate constant, as it provides more energy to overcome the activation energy barrier.

Atomic radius increases down a group. Potassium (K) has a radius of approximately 227 pm.

For a single bond (bond order = 1), the average bond energy is approximately 348 kJ/mol, but for C-C it is typically around 346 kJ/mol.

In a first-order reaction, the rate is directly proportional to the concentration of the reactant. Halving the concentration will halve the rate.

E = E° - (RT/nF)ln(Q); E = 0.34 - (0.0257/2)ln(100) = 0.30 V.

Doubling the concentration of reactants increases Q, which decreases E according to the Nernst equation.

Heat released = ΔH * n = -150 kJ * 3 mol = -450 kJ.

ΔH per mole = 75 kJ / 2 mol = 37.5 kJ/mol.

ΔH for 1 mole = -500 kJ / 5 mol = -100 kJ/mol.

The combustion of carbon to form CO2 releases -393.5 kJ, which is the enthalpy change for the reaction.