consider the following equilibrum reaction:
2NO yield N2 + O2.
Take:delta H=-18.0KJ/M
To determine the value of the equilibrium constant (K) for the given reaction, we can use the relationship between K and the change in enthalpy (∆H).
The equilibrium constant (K) is defined as the ratio of the concentrations of the products to the concentrations of the reactants, each raised to their respective stoichiometric coefficients. In this case, the stoichiometric coefficients from the balanced equation are:
Reactants: 2NO
Products: N2 + O2
For the given reaction, we have the relationship:
K = [N2] * [O2] / [NO]^2
Since we do not have any specific concentrations provided, we cannot determine the numerical value of the equilibrium constant (K), but it can still be described qualitatively based on the value of ∆H.
The value of ∆H gives us information about the energy change that occurs during the reaction. In this case, the given ∆H is -18.0 kJ/mol. A negative value of ∆H implies that the reaction is exothermic, meaning that it releases heat energy.
By applying Le Chatelier's principle, we can understand the effect of ∆H on the equilibrium constant (K). According to Le Chatelier's principle, an increase in temperature (which corresponds to a positive ∆H) favors the endothermic direction, while a decrease in temperature (which corresponds to a negative ∆H) favors the exothermic direction.
Since the given ∆H is negative (-18.0 kJ/mol), the reaction is exothermic. As a result, a decrease in temperature will favor the formation of products (N2 and O2), shifting the reaction towards the right. Consequently, the equilibrium constant (K) for this exothermic reaction will be larger than 1.
In summary, although we cannot determine the specific value of the equilibrium constant (K) without concentration values, we can conclude that it will be greater than 1 for the given exothermic reaction.