Some AgNO3, NaCl, and AgCl are added to water and the following reaction consumes some AgCl until equilibrium is established. The temperature and pressure remain constant. Note: AgNO3 and NaCl are highly soluble in water. AgCl has limited solubility in water.
Ag+(aq) + Cl−(aq) ⇌ AgCl(s)
Which of the following conditions must be satisfied: (i) initially; and (ii) at equilibrium?
(i) ΔrG > 0 and Q > K (ii) ΔrG = 0
(i) ΔrG° > 0 and Q > K (ii) ΔrG° = 0
(i) ΔrG > 0 and Q < K (ii) ΔrG = 0
(i) ΔrG° < 0 and Q < K (ii) ΔrG° = 0
(i) ΔrG < 0 and Q < K (ii) ΔrG = 0
i think its (i) ΔrG > 0 and Q > K (ii) ΔrG = 0
pls confirm
I'm sorry but I don't understand the problem. In my mind there is no way that AgCl can be consumed. The solution MUST precipitate more AgCl so AgCl is formed and not consumed. Also, adding AgNO3 and NaCl to a solution that was already at equilibrium with AgCl(s) the MORE AgCl would ppt.
To determine the conditions that must be satisfied initially and at equilibrium, we need to understand the concepts of ΔrG, Q, and K.
ΔrG is the change in Gibbs free energy, which measures the spontaneity of a reaction. It is related to the equilibrium constant, K, through the equation:
ΔrG = ΔrG° + RTln(Q/K)
where ΔrG° is the standard Gibbs free energy change, R is the gas constant, T is the temperature in Kelvin, Q is the reaction quotient, and ln stands for the natural logarithm.
Q is the reaction quotient, which is the ratio of the product of concentrations of the products raised to their stoichiometric coefficients divided by the product of concentrations of the reactants raised to their stoichiometric coefficients. It is calculated using the concentrations at any given point in the reaction.
K, on the other hand, is the equilibrium constant, which is the ratio of the product of the concentrations of the products raised to their stoichiometric coefficients divided by the product of concentrations of the reactants raised to their stoichiometric coefficients. It is a constant value at equilibrium.
Now, let's analyze the given options:
(i) ΔrG > 0 and Q > K: This means that the reaction is non-spontaneous, as ΔrG is positive. Also, Q is greater than the equilibrium constant K. However, at equilibrium, ΔrG should be zero (ii) ΔrG = 0), so this condition does not hold.
(i) ΔrG° > 0 and Q > K: This option suggests that the standard Gibbs free energy change is positive, indicating a non-spontaneous reaction. And just like the previous option, Q is still greater than K. Again, at equilibrium, ΔrG should be zero (ii) ΔrG = 0), so this condition is not satisfied either.
(i) ΔrG > 0 and Q < K: Here, we have a non-spontaneous reaction (ΔrG > 0) and Q is smaller than K. Once again, at equilibrium, ΔrG should be zero (ii) ΔrG = 0), so this condition does not apply.
(i) ΔrG° < 0 and Q < K: This option implies that the standard Gibbs free energy change is negative, indicating a spontaneous reaction. Additionally, Q is less than K, suggesting that the reaction favors the products. At equilibrium, ΔrG should be zero (ii) ΔrG = 0), so this condition satisfies our requirements.
(i) ΔrG < 0 and Q < K: This condition is similar to the previous one. ΔrG < 0 indicates a spontaneous reaction, and Q is still smaller than K. Moreover, at equilibrium, ΔrG should be zero (ii) ΔrG = 0). Hence, this condition is also satisfied.
To summarize, the correct conditions that must be satisfied initially and at equilibrium are:
(i) ΔrG < 0 and Q < K
(ii) ΔrG = 0
Option: (i) ΔrG < 0 and Q < K (ii) ΔrG = 0