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### Course: AP®︎/College Chemistry>Unit 8

Lesson 3: Weak acid and base equilibria

# Weak acid equilibria

Unlike strong acids, weak acids only partially ionize in aqueous solution. As a result, the concentration of H₃O⁺ in a weak acid solution is typically much less than the initial acid concentration. The relative strength of a weak acid is described by the acid ionization constant, Kₐ, which is the equilibrium constant for the reaction of the weak acid with water. Created by Jay.

## Want to join the conversation?

• Why is pure water left out of the equation?
• It is conventional to omit pure liquids (like water), solids, and solvents from equilibrium expressions. This is a consequence of exactly how each quantity is defined. They aren't defined simply as the concentrations of the chemical species but rather as the ratio of the concentration of a species to its concentration in its standard state. This standard state is defined differently depending on the chemical species. For solutes, the standard state is 1 M, For gases the standard state is 1 bar, and for solids and liquids the standard states are the pure solid or liquid. Mathematically this means that the concentrations of a solid or liquid are the same as their concentration in the standard state creating a ratio equal to 1. This results in solid or liquid quantities essentially being canceled out since they simplify to 1. If it's a solvent then the concentration is so close to that of a pure liquid that the quantity is essentially 1 as well. This is actually also why equilibrium constants are unitless since defining all the quantities as such results in them being unitless making the entire expression dimensionless. This convention is arbitrary and there's no deeper meaning as to why we do it other than to make the math simpler.

Hope that helps.
• I was doing tests to Lesson 3 and got a need to use a hint for one task. "To achieve a certain pH a weaker acid will always require a higher initial concentration than a stronger acid." This sentence seems to be contradictory considering that "At a given concentration, a weaker acid forms a solution with a higher pH than a stronger acid does.". I can't logically get it. Is it a general rule with some complicated reason underlying it?
(1 vote)
• "At a given concentration, a weaker acid forms a solution with a higher pH than a stronger acid does."

If we have a strong and a weak acid at the same concentration, the strong acid solution will have a lower pH than the weak acid solution. A lower pH corresponding to a more acidic solution. The strong acid will dissociate more so than the weak acid. This means there will be more acidic protons in solution in the strong acid solution than for the weak acid solution. More acidic protons mean the solution is more acidic with a lower pH.

“To achieve a certain pH a weaker acid will always require a higher initial concentration than a stronger acid."

If we have a strong and weak acid and we want them to be at the same pH, we’ll need a higher concentration of the weak acid to do so. Again if we keep the two acids at the same concentration, then the strong acid will be more acidic than the weak one. So we need a more concentrated weak acid solution to compensate for its weaker strength.

Hope that helps.
• Since pH + pOH=14.00 is derived from the euqation: [H+]*[OH-]=K(water), why we can still apply pH + pOH = 14.00 to all aqueous solution even when K(acid)>>1, which is much larger than K(water). Wouldn't it be pH + pOH = pK(acid)(probably much smaller than 14.00)?
(1 vote)
• The autoionization of water reaction, described by Kw, is a separate reaction from an acid dissociation reaction, described by Ka.

Ka>>1 means that the acid dissociation reaction lies far to the right, towards the products. It also means that the acid is strong and has a large majority of the original acid molecules dissociate. As opposed to a weak acid with a small Ka which only dissociates partially.

If we have a strong acid dissociate, then it contributes some hydronium to the aqueous solution to make it more acidic. But just as the acid reaction reaches an equilibrium, the water itself also reaches its own equilibrium. At 25°C Kw is 1.0x10^(-14) and that means the water autoionization reaction, Kw = [H3O^(+)][OH^(-)], always has to equal that number at that temperature. If we increase the hydronium concentration, then the hydroxide concentration decreases too to maintain that Kw value.

And since the Kw equation is equal to the pKw = 14 = pH + pOH equation, changes in the hydronium concentration cause subsequent changes in the pH and the pOH. But again since Kw is constant at 25°C, and pKw is based on Kw, the previous equation always equals 14. Strong acid or weak acid, doesn’t matter the Ka, pKw is 14.

Hope that helps.