Q<K?
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Re: Q<K?
If Q<K, the reaction will proceed in the forward direction and produce more products. The concentration of reactants is relatively greater than the concentration of products compared to their equilibrium concentrations.
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Re: Q<K?
Likewise, If Q>K then the concentrations of products is greater than reactants with regards to the reaction's equilibrium constant. This means the reverse reaction would occur in order to reach equilibrium
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Re: Q<K?
The following helps me envision a Q/K comparison:
Since Q=[P]/[R] and Q<K
If <K and K is a constant (doesn't change), the only way such that could be less than K is if denominator (R) was large. It helped to plug in some easy numbers (ex. 1) so that I could get a really firm grasp of the concept.
Since Q=[P]/[R] and Q<K
If <K and K is a constant (doesn't change), the only way such that could be less than K is if denominator (R) was large. It helped to plug in some easy numbers (ex. 1) so that I could get a really firm grasp of the concept.
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Re: Q<K?
Recall that Q is the ratio (in M or partial P) of reactants and products at any time during a given reaction. Meanwhile, K is the ratio when the reactants and products are at equilibrium.
When comparing Q to K, I tend to think of the comparison when comparing K to itself. If K > 1E3, then the reaction strongly lies to the right and product formation is favored. If K < 1E-3, then reaction strongly lies to the left and reactant formation is favored.
The same holds true for comparing Q to K:
If Q < K at any time during the reaction then [R] > [P] and the forward reaction (shift to the right) is favored and more product is formed until equilibrium is achieved again.
If Q > K at any time during the reaction then [R] < [P] and the reverse reaction (shift to the left) is favored and more reactant is formed until equilibrium is achieved again.
On a side note, I liked how William set up the comparison.
When comparing Q to K, I tend to think of the comparison when comparing K to itself. If K > 1E3, then the reaction strongly lies to the right and product formation is favored. If K < 1E-3, then reaction strongly lies to the left and reactant formation is favored.
The same holds true for comparing Q to K:
If Q < K at any time during the reaction then [R] > [P] and the forward reaction (shift to the right) is favored and more product is formed until equilibrium is achieved again.
If Q > K at any time during the reaction then [R] < [P] and the reverse reaction (shift to the left) is favored and more reactant is formed until equilibrium is achieved again.
On a side note, I liked how William set up the comparison.
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Re: Q<K?
If Q<K at some time during the reaction, then the concentration of reactant is greater than the concentration of products therefore the forward reaction is favored. Hope this helps!
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Re: Q<K?
If O<K then this means that the reaction quotient, Q, is less than the value of the equilibrium constant, K, meaning that [R] is greater than [P] thus they want to favor the forward reaction in order to bring the R/P ratio back to equilibrium and reach the equilibrium constant once again. Hope this helps.
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Re: Q<K?
When Q < K, the system favors the forward reaction. This means that there are more reactants than products so the system moves in a forward fashion to try and reach equilibrium.
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Re: Q<K?
If Q<K, then this means the concentration of the reactants are greater than the concentrations of the products ( [R] > [P] ) and the forward reaction is favored.
Conversely, if Q>K, then [P] > [R] and the reverse reaction is favored.
Conversely, if Q>K, then [P] > [R] and the reverse reaction is favored.
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Re: Q<K?
When Q<K, there is an excess of reactants, so the value of Q would be less than the equilibrium constant. Therefore, in order for the reaction to go to equilibrium, we would need a forward reaction that would increase the amount of product and allow for the reaction to reach the equilibrium constant.
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Re: Q<K?
If Q<K, [R]>[P], meaning the forward reaction is favored. The concentration of reactants is greater than the concentration of the products, so the reactants move towards the products in order to reach equilibrium.
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Re: Q<K?
If Q<K, the reaction will proceed in the forward direction because there are more reactants than products in the equation than there are when the reaction is at equilibrium. If Q>K, the reverse reaction occurs because there are more products than reactants in the reaction than are when it is at equilibrium. And lastly, if Q=K, the reaction is at equilibrium!
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Re: Q<K?
When Q<K, the reaction will make more products in order to bring the system to equilibrium. In other words, it will favor the forward reaction.
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Re: Q<K?
If Q is less than K then the reaction proceeds in the forward reaction and the products are favored.
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Re: Q<K?
If Q<K the reaction will shift right/proceed forward. If Q>K the reaction shifts left/the reverse reaction is favored.
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Re: Q<K?
If Q<K, the reaction favors the product side and shifts right (acting in the forward direction) to achieve equilibrium.
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Re: Q<K?
If Q<K , then the reaction favors the products. The ratio of products to reactants is less than that for the system at equilibrium—the concentration or the pressure of the reactants is greater than the concentration or pressure of the products. Because the reaction tends toward reach equilibrium, the system shifts to the RIGHT to make more products.
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Re: Q<K?
Q=[P]/[R], so if Q<K the concentration of reactants is greater than products. This means that the reaction will proceed in the forward direction until it reaches the equilibrium constant again.
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Re: Q<K?
When Q<K, that means the reaction quotient is a smaller value than the equilibrium constant, so when comparing Q to the ratio of K, there is too much reactant and not enough product given that Q is less than K, meaning that the reaction will shift to the right (the forward reaction) and produce product until the system reaches equilibrium (when Q is equal to K). (Q=[P]/[R], K=[P]/[R]).
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Re: Q<K?
It means that there were more reactants than usual, so the reaction will favor the products (forward rxn).
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Re: Q<K?
Hey there!
When K>Q, this means that the reaction will shift right and favor the formation of products (the forward reaction) since at this particular point in time there are more reactants than there would be at equilibrium.
I hope this helps!
When K>Q, this means that the reaction will shift right and favor the formation of products (the forward reaction) since at this particular point in time there are more reactants than there would be at equilibrium.
I hope this helps!
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Re: Q<K?
Q < K means that in that instant, the forward reaction is favored since the ratio of products to reactants is less than normal. Because of Le Chatlier's principle, the system works to re-establish equilibrium which, in the case of Q < K, means making more products.
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Re: Q<K?
If Q<K, then at the moment Q is being calculated, there are too many reactants and not enough products to reach the equilibrium value, since Q and K both = products/reactants. Therefore, the reaction will need to shift right/proceed forward to produce more products, therefore raising the Q value to reach equilibrium.
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Re: Q<K?
When Q is less than K it means that there are significantly more reactants than products, favoring the forward reaction. This means more product will be formed.
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Re: Q<K?
When Q<K, the forward reaction is favored, meaning that there is a greater concentration of reactants, which are creating more products.
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Re: Q<K?
When Q<K, the products are favored. Since, Q/K = [P]/[R], the situation where Q<K would mean that the reaction quotient has less products or too many reactants needed to establish equilibrium. Meaning, in order to establish equilibrium, the forward reaction would be favored to produce more products, so that Q = K.
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Re: Q<K?
Q<K ⮞
The fraction representing Q must increase its numerator/decrease its denominator (product⭧/reactant⭨) to reach K, so the reaction will favor products (shift right).
The fraction representing Q must increase its numerator/decrease its denominator (product⭧/reactant⭨) to reach K, so the reaction will favor products (shift right).
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Re: Q<K?
If Q > K the forward reaction is favored because there was a higher concentration of reactants than products. Conversely, if Q < K the reverse reaction is favored for the opposite reasoning.
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Re: Q<K?
if Q is less than K, then rate of the forward reaction is faster and more products than reactants are being made (reaction shifts left)
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Re: Q<K?
When Q is less than K, the concentration of the reactants in greater than the concentration of the products and a forward reaction is favored.
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Re: Q<K?
If Q<K, the reaction currently has more reactants/less products than the reaction would at equilibrium. Thus, the reaction will proceed in the forward direction.
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Re: Q<K?
If Q<K at some point during the reaction, this means that the concentration of reactants is greater than the concentration of products. This makes the forward reaction be favored meaning more products will be made.
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Re: Q<K?
When Q is less than K, the reaction has a greater concentration of reactants compared to that of the products. Additionally, this means a forward reaction is favored.
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Re: Q<K?
When Q<K, that means that the concentration of reactants is higher than the products, therefore the reaction will shift to the right in the forward direction.
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Re: Q<K?
Hello! A visual that really helped me understand the relationship between reaction quotients and equilibrium constants is the image below.
If Q is less than K, the reaction is going in the forward direction because there are fewer products formed than the system would like to be at. Because of this, the system will shift towards the production of products to compensate for the lack of products.
If Q is more than K, the reaction is going in the reverse direction because there are more products formed than the system would like to be at. Because of this, the system will shift towards the formation of reactants to compensate.
If Q is equal to K, there is no shifting taking place.
The image shows that K is the ideal state of the system, and if Q is above or below that, the system will naturally proceed in a way to achieve the state of equilibrium.
If Q is less than K, the reaction is going in the forward direction because there are fewer products formed than the system would like to be at. Because of this, the system will shift towards the production of products to compensate for the lack of products.
If Q is more than K, the reaction is going in the reverse direction because there are more products formed than the system would like to be at. Because of this, the system will shift towards the formation of reactants to compensate.
If Q is equal to K, there is no shifting taking place.
The image shows that K is the ideal state of the system, and if Q is above or below that, the system will naturally proceed in a way to achieve the state of equilibrium.
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Re: Q<K?
Hi, Q<K means that the reaction will shift to the right in the forward direction since the concentration of reactants is higher than the products.
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Re: Q<K?
Q and K are both ratios of reactant to product, and the only difference is that Q is not under equilibrium. Therefore, because the product is the numerator and the reactant is the denominator, if Q<K, it means that the current product is less than the product concentration under equilibrium. Thus, the reaction will shift right, the side that favors more product formation.
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Re: Q<K?
When Q<K, the reactants are still going to products during the RxN. The forward reaction is favored.
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Re: Q<K?
When Q<K that means the reaction favors the products and the forward reaction will be more favored, causing production of products.
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Re: Q<K?
When K is greater than Q , our forward reaction is favored, meaning there are more products than reactants
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Re: Q<K?
When Q is less than K it means that the Reactant concentration increased and therefore the forward reaction will be favored
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Re: Q<K?
If the reactant quotient(Q) is less than the equilibrium constant (K), then the current ratio of products to reactants is less than what is expected when the reaction reached equilibrium. Thus, the chemical equation will favor the forward direction to increase the concentration of products in order to reach equilibrium.
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Re: Q<K?
If Q is greater than K than that means there is a higher concentration of reactants vs. products and so the reaction will proceed in the forward direction (toward the products).
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Re: Q<K?
Hi! Q<K means during RxN, R>P (reactants > products), and the forward reaction is favored & proceeds toward products. When Q>K during RxN, then P>R and the reverse reaction is favored & proceeds toward reactants.
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Re: Q<K?
The relationship between Q and K determines the direction in which the reaction will proceed. When Q>K, then the products are favored, and when Q<K, the reactants are favored.
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Re: Q<K?
Q determines the current state of the reaction, while K displays the ideal product-reactant concentration ratio for equilibrium. If Q<K, it indicates there are more reactants and fewer products than there should be at equilibrium, and the reaction would adjust. Similarly, the opposite is occurring if Q>K.
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Re: Q<K?
If Q<K, it means that the reaction has not yet reached equilibrium, and more products have yet to be created. Remember that Q=K at equilibrium and Q = [P]/[R] at any given time. Hope this helps!
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Re: Q<K?
If Q<K, the concentration of reactants is greater than the concentration of products. This means that the reaction will proceed in the forward direction and favor the formation of products.
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Re: Q<K?
Hi,
When Q<K, the reaction is not in equilibrium and we will have more products formed as the forward reaction is favored.
When Q<K, the reaction is not in equilibrium and we will have more products formed as the forward reaction is favored.
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Re: Q<K?
If Q is less than K, that means the equilibrium state has a larger product to reactant ratio than the current state Q. This means the reaction will proceed in the forward direction.
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Re: Q<K?
K represents the constant when the chemical equation reaches a dynamic equilibrium, and Q indicates the instant status of a chemical reaction. Therefore, when Q is smaller than K, we can infer based on this information in which direction the chemical reaction would proceed in. The K constant is calculated using products over reactants, so if Q is smaller than K, it means that Q has more reactants than products in its product to reactant ratio compared to K. Therefore, in order to proceed to chemical equilibrium which is represented by K. it will orient itself in favor of aligning its value with K. Therefore, it will make its reactants go down and products go up, which means the chemical reaction will proceed in the forward direction, decreasing reactants and increasing products.
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Re: Q<K?
When Q<K, that means that the reaction favors the forward reaction because there is still more reactant to form products with.
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Re: Q<K?
Q<K means that the forward reaction is favored. There is still more reactants to form products.
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Re: Q<K?
When Q is less than K, there are more reactants than products than there should be at equilibrium. This means the equilibrium should shift to the products.
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Re: Q<K?
if you remember, Kc or Qc= [products]/[reactants]. If Q<K, this means there are more reactants and the formation of products would be favored in order to reach equilibrium. Hope this helps!
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Re: Q<K?
Since Q and K are both calculated by the equation products/reactants, Q being less than K means that there are more reactants (greater denominator) in the current reaction than there is at equilibrium, shown by K. This means that the reaction still has to form more products, such that the forward reaction is favored.
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Re: Q<K?
If Q is less than K, this means that the reaction is not at equilibrium and that the concentration of the products at that point in the reaction is less than what the concentration of products is at equilibrium, and that the forward reaction is favored as that reaction continues towards equilibrium.
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Re: Q<K?
If Q<K then this means that there are more reactants in relation to the products. Thus the reaction shifts to the right.
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Re: Q<K?
If Q>K then this means that there are more products in relation to reactants. Thus the reaction will shift to the left and favor the reactants.
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Re: Q<K?
when q<k the amount of products at the time is greater than the amount of products at equilibrium so the reaction will shift left towards reactants to reach equillibrium
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Re: Q<K?
Hi!
If Q<K, this means that in the current state of the system, there are more reactants than there should be at equilibrium, so the system will favor products and shift to the right to produce more products to reach equilibrium.
Hope this helps!
If Q<K, this means that in the current state of the system, there are more reactants than there should be at equilibrium, so the system will favor products and shift to the right to produce more products to reach equilibrium.
Hope this helps!
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Re: Q<K?
It means that the reaction will favor the production of products to move towards equilibrium. Because the reaction quotient measures [P]/[R] of the reaction not in equilibrium it is telling of the direction that the reaction will move in to achieve equilibrium. Having a value less than K means that its concentration of products is less than what it should be in equilibrium. Thus the reaction will make more products so that it moves towards equilibrium.
Re: Q<K?
When Q is less than K, that means the reaction will be proceeding in the forward direction, where products are being formed. When Q>K, that means the process is going in the reverse direction, where products will go towards reactants.
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Re: Q<K?
Q<K it means then it will proceed towards the forward direction while as if Q>K then the reaction will proceed in the reverse direction
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Re: Q<K?
When Q is less than K, the reaction will go in the forward direction which means that reaction will favor the production of products.
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Re: Q<K?
When Q<K, you have more reactants so the reactions will shift in the forward direction to produce more products
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