# Ka relationship to ph Relationship between Ka, pKa and acid strength: If the pH of a solution of a weak acid and the pKa are known, the ratio of the concentration of the conjugate . It can be used to calculate the concentration of hydrogen ions [H+] or hydronium ions [H3O+] in an aqueous solution. \(K_a\), the acid ionization constant, is the equilibrium constant for chemical reactions involving weak acids in aqueous solution. Calculate the \(K_a\) value of a. I'm guessing that most of us that work in the lab would find it more convenient to measure a pH rather than to calculate a theoretical value.

So this is the acid ionization constant or you might hear acid dissociation constant, so acid dissociation. So either one is fine. All right and we know when we're writing an equilibrium expression, we're gonna put the concentration of products over the concentration of reactants.

## pKa and Dissociation Equilibrium

Over here for our products we have H3O plus, so let's write the concentration of hydronium H3O plus times the concentration of A minus, so times the concentration of A minus.

All over the concentration of our reactant, so we have HA over here, so we have HA. So we could write that in and then for water, we leave water out of our equilibrium expression. It's a pure liquid.

Find the Ka of an acid (Given pH) (0.1 M Hypochlorous acid) EXAMPLE

Its concentration doesn't change and so we leave, we leave H2O out of our equilibrium expression. All right, so let's use this idea of writing an ionization constant and let's apply this to a strong acid.

HCL is gonna function as a Bronsted-Lowry acid and donate a proton to water which is going to be our Bronsted-Lowry base. And so we could think about a loan pair of electrons in the auction taking our proton, leaving those electrons behind.

And so the auction is now bonded to three hydrogens because it picked up a proton, giving this a plus one charge. Once again let's follow those electrons in red. This electron pair picks up this proton to form this bond, so we form H3O plus or hydronium.

And these electrons in green move off onto the chlorine, so let's show that. We form the chloride anion. Let me go ahead and draw in the electrons in green and let me go ahead and write a negative one charge here like that. So this is just a faster way of doing it and HCL is a strong acid. The equilibrium is so far to the right that I just drew this one arrow down over here.

So KA is equal to a concentration of H3O plus. So concentration of our products times concentration of CL minus, all over, right, we have HCL and we leave out water. So we have a very, very large number in the numerator and extremely small number in the denominator.

If you think about what that does for your KA, that's gonna give you an extremely high value for your KA. All right, so KA is much, much, much greater than one here. That's how we recognize a strong acid. An acid ionization constant that's much, much greater than one. Now let's think about the conjugate base. All right, so let's go back up here. So we had a HCL and CL minus as our conjugate acid base pair and the stronger the acid, the weaker the conjugate base. So let me write that here. The stronger the acid, so stronger the acid, weaker the conjugate, weaker the conjugate base. And one way to think about that is if I look at this reaction, we can think about competing base strength. All right, so here we have Bronsted-Lowry. Base water is acting as a Bronsted-Lowry base and accepting a proton. And over here if you think about the reverse reaction, the chloride anion would be trying to pick up a proton from hydronium for the reverse reaction here but since HCL is so good at donating protons, that means that the chloride anion is not very good at accepting them.

### Acids and bases | Chemistry | Science | Khan Academy

So the stronger the acid, the weaker the conjugate base. Water is a much stronger base than the chloride anion. Finally let's look at acetic acids. Acetic acid is going to be our Bronsted-Lowry acid and this is going to be the acidic proton. Water is gonna function as a Bronsted-Lowry base and a lone pair of electrons in the auction is going to take this acidic proton, leaving these electrons behind on the oxygen.

The square brackets indicate the concentration of respective components. Based on this equation, Ka expresses how easily the acid releases a proton in other words, its strength as an acid.

Carboxylic acids containing -COOHsuch as acetic and lactic acids, normally have a Ka constant of about to Consequently, expressing acidity in terms of the Ka constant alone can be inconvenient and not very intuitive. Therefore, pKa was introduced as an index to express the acidity of weak acids, where pKa is defined as follows. In addition, the smaller the pKa value, the stronger the acid. For example, the pKa value of lactic acid is about 3. This relationship is described by the following equation.

This equation can be rearranged as follows. If the pH changes by 1 near the pKa value, the dissociation status of the acid changes by an extremely large amount.