If the reaction is exothermic, decreasing the temperature will shift the equilibrium to the right, favoring NH3 production.

Increasing the temperature of an exothermic reaction shifts the equilibrium to the left, favoring the reactants.

A negative ΔH indicates that the reaction releases heat, thus it is exothermic.

At equilibrium, ΔG = 0; ΔG = ΔH - TΔS = 0; T = ΔH/ΔS = 150,000 J / 200 J/K = 750 K.

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.

Molar mass of NaCl = 58.44 g/mol. Moles needed = 0.2 M * 0.5 L = 0.1 moles. Mass = moles * molar mass = 0.1 * 58.44 g = 5.844 g.

The reaction between KOH and HCl is 1:1. Therefore, 0.1 moles of HCl will require 0.1 moles of KOH.

2,3-dimethylbutane has no chiral centers, so it has only 1 stereoisomer.

2-pentene has two geometric isomers: cis-2-pentene and trans-2-pentene.

1,2-dichloropropane has one chiral center, leading to 2 stereoisomers.

1 mole of NaCl dissociates into 2 particles (1 Na⁺ and 1 Cl⁻), so there are 2 particles in solution.

Using the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]). pKa of acetic acid is approximately 4.76. Since [A-] = [HA], pH = 4.76.

H2SO4 dissociates to give 2 H+ ions. Moles of H2SO4 = 0.1 M * 0.050 L = 0.005 moles. Therefore, moles of NaOH required = 0.005 moles * 2 = 0.01 moles.

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

Using Beer-Lambert Law (A = εlc), A = 200 L/(mol·cm) * 0.01 mol/L * 1 cm = 2.