Doubling the concentration of A will shift the equilibrium to the right, but Kc remains constant at 4.

Since Kc < 1, it indicates that at equilibrium, the concentration of A is greater than that of B.

Kp is independent of volume changes; it remains 25.

The rate constant k for a first-order reaction is given by k = 0.693/t1/2. Thus, k = 0.693/10 min = 0.0693 min^-1.

For a first-order reaction, the half-life (t1/2) is given by t1/2 = 0.693/k. Rearranging gives k = 0.693/t1/2 = 0.693/10 min = 0.0692 min^-1.

Using the Arrhenius equation, a doubling of the rate constant corresponds to an activation energy of approximately 40 kJ/mol for a 10°C increase.

Using the Arrhenius equation and the rule of thumb that a 10°C increase roughly doubles the rate constant, we can estimate the activation energy to be around 40 kJ/mol.

If doubling the rate occurs when tripling the concentration, the reaction is first order with respect to A, as rate ∝ [A]^n implies 2 = 3^n.

Using the Arrhenius equation and the two-point form, Ea can be estimated as Ea ≈ 2.303RT²(Δlnk/Δ(1/T)). With R = 8.314 J/(mol·K), Ea ≈ 40.8 kJ/mol.

For a first-order reaction, rate = k[A]. Thus, k = rate / [A] = 0.05 M/s / 0.1 M = 0.5 s^-1.

If the rate is proportional to the square of the concentration of A, the rate law is rate = k[A]^2.

If tripling the rate occurs when doubling the concentration, the reaction is first order with respect to A, as rate ∝ [A]^n implies 3 = 2^n, leading to n = 1.

E = E° - (0.059/3)log(1000) = 0.77 - 0.059/3 * 3 = 0.70 V.

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

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.

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.

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.

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.

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.

According to Le Chatelier's Principle, increasing the temperature of an exothermic reaction shifts the equilibrium to the left, favoring the reactants.

Using q = mcΔT, ΔT = q/(mc) = 4200 J / (100 g * 4.18 J/g°C) = 10 °C.

q = mcΔT = 50 g × 4.18 J/g°C × (75-25)°C = 1045 J.

A negative q value indicates that heat is released by the system, signifying an exothermic process.

A negative ΔH indicates that the reaction releases heat to the surroundings, classifying it as exothermic.

A positive q value indicates that the system absorbs heat from the surroundings.

A large equilibrium constant (K) indicates that at equilibrium, the concentration of products is much greater than that of reactants, favoring products.

A negative ΔH indicates that the reaction releases heat to the surroundings, classifying it as exothermic.

A negative ΔH indicates that the reaction releases heat, classifying it as exothermic.

According to Gay-Lussac's Law, pressure is directly proportional to temperature at constant volume (P/T = k).

According to Gay-Lussac's Law, at constant volume, the pressure of a gas is directly proportional to its temperature in Kelvin.