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.

Transmittance (T) can be calculated using the formula T = 10^(-A). For A = 0.5, T = 10^(-0.5) = 0.316, which is approximately 50%.

Using Beer-Lambert Law, A = εcl = 200 L/(mol·cm) * 0.1 mol/L * 1 cm = 20.

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.

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.

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.

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.

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.

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.

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.