Q<K

[ursulavictorino1K]

Why is it that if Q<K, the forward reaction is favored if [R]>[P]?

[Kyle Rex 1K]

since there is a higher concentration of reactants, the excess of reactants will react to form product, which is why it favors the products since it goes in that direction

[Kishan Shah 2G]

As Kyle said, if Q<K then there is more reactants at that instance, so in order to reach equilibrium (Q=K), the reaction needs to proceed toward the products.

[CynthiaLy4F]

The forward reaction is favored when Q<K because at that given stage of the reaction, there is a lot more reactants than products. Therefore, for Q to increase and reach equilibrium, more products must be produced.

[Ryan Yoon 1L]

When Q<K, this signifies that there are more reactants than the product currently at the system. Since there are more reactants than products, the reaction will go forward.

[Tauhid Islam- 1H]

When Q is less than K, we know that in the system, there are more reactants compared to products, so for the system to reach the ratios of products to reactants that is detonated by K, the forward reaction is favored where you would end up making more products to get the ratio to be K.

[Miriam Villarreal 1J]

Q<K means there are more reactants than products therefore in order to reach equilibrium more product is needed and therefore rxn will shift to the right (forward)
Q>K means that there are more products than reactants therefore in order to reach equilibrium more reactant is needed and therefore rxn will shift left (reverse)

[Catherine Daye 1L]

This is because the numerator of Q needs to increase to reach equilibrium (K value), and the concentration of products is the numerator.

[Leyna Dang 2H]

When Q<K, the forward reaction is favored because the concentrations/partial pressures of the products are too low compared to reactants for equilibrium. Therefore, the excess reactants would form more products in order to reach equilibrium.

[Daria MacAuslan 1H]

The forward reaction is favored because of the fact that changing the amount of reactants by increasing it will allow for more reactions to take place, ultimately increasing the amount of products and favoring this forward reaction.

[WYacob_2C]

Q is also equal [products]/[reactants]. If Q is less than K, that means the concentration of reactants is too much, causing the ratio between products and reactants to be small. In order to fix this, the chemical reaction will favor the formation of products to increase this ratio, in hopes of meeting the K value.

[Nicholas_Gladkov_2J]

Why is it that if Q<K, the forward reaction is favored if [R]>[P]?

As Q and K are both the concentration of products over reactants, if Q is less than K, then in order for Q to be equal to K, Q has to get bigger. In order for Q to get bigger, the concentration of products have to be larger. Therefore, the forward reaction is favored.