If a weak acid is say only 5% ionized at equilibrium, then the ionization rxn would be reactant favored, correct? I am a little on the fence about understanding this if anyone could better explain why this would be? I am assuming because it is not completely ionized, the amount of products would be less than the amount of reactants and so it would be reactant favored.
Here is the way I do it.
1. Write the reaction. Let's use acetic acid, CH3COOH (which I will call HAc), as an example.
HAc ==> H^+ + Ac^-
2. Write the Ka expression. For a weak base write the Kb expression. For an equilibrium reaction, write Keq. They all work the same.
Ka = 1.8 x 10^-5 = (H^+)(Ac^-)/(HAc).
3. Analyze.
It's as simple as this. Ka is small; therefore, the numerator of the fraction must be smaller than the denominator. Which means, of course, that the ions are less than the un-ionized material and that means the products are not favored (the reactants are favored). I look at it that the forward reaction is not favored; the reverse reaction is favored. (Another way I find useful is to say, "not many ions, a lot of un-ionized acid")
4. Students always want to know what I consider small; i.e., what Ka, Kb, Keq, is small and what is large. For a 1:1 set of coefficients, a Ka, Kb, Keq of 1 means numerator of fraction = denominator which means the reaction is 50%. So I use 1 as a benchmark although reactions that are not 1:1 vary slightly but as a quick assessment, K of 1 is a good number to choose.
Yes, you're on the right track. When a weak acid is only 5% ionized at equilibrium, it means that only a small portion of the acid molecules have dissociated into ions. In this case, the reaction would indeed be reactant favored.
To better explain why this is the case, let's consider the dissociation reaction of a weak acid (HA) in water:
HA ⇌ H+ + A-
At equilibrium, the reaction can be described by the equilibrium constant, Ka. The equilibrium constant is the ratio of the concentrations of the products (H+ and A-) to the concentration of the reactant (HA). In this case, since only 5% of the acid has dissociated, the concentration of the products (H+ and A-) would be much lower than the concentration of the reactant (HA).
Therefore, the concentration of the reactants would be higher compared to the concentration of the products at equilibrium, indicating that the reaction is reactant favored.
It's important to note that the degree of ionization (or the proportion of the acid that has dissociated) depends on the strength of the acid. Strong acids (like HCl) fully dissociate into ions, while weak acids (like acetic acid) have a lower degree of ionization.
I hope this explanation helps clarify why the reaction would be reactant favored when a weak acid is only 5% ionized at equilibrium! If you have any further questions, feel free to ask.