Achieve W1 Q7

[Tala Ayoub 1H]

Consider the following system at equilibrium.
A(aq)↽−−⇀B(aq) ΔH=+850 kJ/mol
What can be said about Q and K immediately after an increase in temperature?
Q<K because K increased.
I know that the answer to this question is Q<K because K increased but I'm a bit confused, I know that this reaction is exothermic since energy is being released so therefore, the reactants would be favored, I assume that since reactants are favored either Q or K would decrease, could someone please explain.
Thank you!

[Natalie Keung 1D]

I believe that this reaction is actually endothermic because there is the positive delta H.
If you think of it as requiring energy to create the products, a higher temperature would increase the amount of products, increasing K. This should explain why Q<K.

[905756606]

Since the change in enthalpy is positive, this indicates that the forward reaction is endothermic. An increase in temperature causes the equilibrium to shift towards the RHS in favour of the products, since the system wants to counteract that change (Le Chatelier's principle) by using up energy in the forward reaction. Since the concentration of products increases, the value of K increases too. Hope this helps!

[Akane Shimoyoshi 3F]

Since delta H is positive, this reaction is endothermic and heat is a reactant. With an increase in temperature, the equilibrium shifts toward products since there is essentially an increase on the reactant side. Because K increases from the product increase, Q<K.

[Holly Do 2J]

This reaction in endothermic, meaning it absorbs heat to get the product, because of the positive delta H value. Because we know this, if we are increasing the temperature, adding more heat, the reaction will favor the products because heat is needed to create product.

[Zoe Dhalla 3I]

A positive change in enthalpy indicates an endothermic reaction. For endothermic reactions, an increase in temperature causes an increase in K. The reaction quotient is not affected by the temperature change, since there is no change in concentrations. Thus, Q <K, and the net reaction is toward products.

Also, consider that heat is consumed (absorbed) in an endothermic reaction.

Just as a reaction shifts right when more reactant is added, the reaction must shift right when heat is added by increasing temperature.

[Madison Kiggins 1E]

Hi, this reaction would be endothermic due to the positive delta h which indicates that heat is required to form the products. When heat is added, the products will be favored as there is essentially an increase in reactants with the additional heat. Therefore, the K value will go up, and Q<K immediately after the increase.

[indigoaustin 3H]

Since ∆h is positive, this reaction is actually endothermic, meaning it requires heat. An increase in temperature causes the equilibrium to shift towards the products since heat is a reactant. The system is originally at equilibrium where Q=K, so when it shifts towards the products K increases making Q<K.

[Savannah Licciardello 2C]

Because you are given a positive entropy, this is an endothermic reaction, which favors the product. As a result, the K value will increase, making it larger than the Q value so Q<K. As the reaction favors the product, the reaction shifts right.

[Saebean Yi 3E]

What can be said about Q and K after an increase (or decrease) in temperature has to do with whether the reaction is endothermic or exothermic. The reaction is actually endothermic, not exothermic, because the ΔH value is positive.

So, if the reaction is endothermic, an increase of temperature would increase K. A simple way to remember this, is to think of it like an equation. Since it's endothermic, the equation would be like: heat + reactants = products. If you increase temperature (add more heat), that would cause the equilibrium to shift towards the products. As a result, since products is on the "top" of the fraction when calculating K, K would increase. And I think this should also answer the other part of the problem.

[Hannah Choi 1K]

Since there is an increase in heat (ΔH is positive) we can conclude that this is an endothermic reaction, meaning heat is required to form the product. Because heat is needed to form the product, adding heat will add more product. The increase in product will cause K to increase (recall K = [P]/[R]). Therefore, Q < K, and the reaction shifts right in favor of the products.

[Clarissa Le 1I]

Since ΔH is positive, the reaction is endothermic! So when heat is applied, the reaction will favor product formation / shift to the right. K would then increase because of a change in temperature.

[taline krumian 1L]

A positive delta H value means the reaction is endothermic, as the system is gaining heat. Therefore, increasing temperature increases the K value, causing Q<K. The reaction will also shift right.

[Daniela G 2C]

Because delta H is positive, identifying a positive change in enthalpy, the reaction can be identified as endothermic. In endothermic reactions, a positive increase in temp is associated with K increasing.

[Harshitha_Pandian_3F]

Delta H is positive which means that the reaction is endothermic. This means heat enters the system so the increase in temperature would shift the equilibrium towards the products so as K increases the product increases.

[Anubhav_Chandla1G]

Similar to what everyone has said above, since the change in enthalpy was positive, the reaction can be denoted as an endothermic reaction since the reaction gained energy via heat. Due to this, the temperature of the reaction increased positively, leading to K increasing.