When to Use Kw & Ka for Acid Dissociation Calculations

In summary, the concentration of H+ in strong acids that dissociate completely is the same as the concentration of the initial solution, meaning 1M of a strong acid will create 1M of H+ with a pH of 0. In weak acids, the acid dissociation constant must be used to calculate the concentration of H+. However, in some cases, the Kw can be used to calculate the pH of acids with low concentrations. The use of Kw depends on whether the H+ from water autodissociation can be ignored or not. For weak acids, Ka can also be used to find Kb, with the H+ concentration being 10^-4 resulting in a base concentration of 10^-10.
  • #1
Nishantkumar19
16
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So I've learned that in strong acids that dissociate completely, the concentration of H+ is the same as the concentration of the initial solution. So 1M of a strong acid will create 1M of H+, meaning the pH is 0.

I've also learned that in weak acids, the whole thing doesn't dissociate. I have to use the acid dissociation constant to calculate the concentration of H+.

But in some parts of my textbook, it also says I can use Kw to calculate the pH of acids, eg. if a solution has a concentration of 10^-4 H+ ions, then the pH is 4.

But wouldn't I need to first find the acid dissociation constant of that solution?

In other words, when can I use Kw, and when do I have to calculate Ka? Can Kw be used only for water, or does it apply to some acids as well?

Thanks a lot for your help.
 
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  • #2
Nishantkumar19 said:
But in some parts of my textbook, it also says I can use Kw to calculate the pH of acids, eg. if a solution has a concentration of 10^-4 H+ ions, then the pH is 4.

How does it differ from the

Nishantkumar19 said:
in strong acids that dissociate completely, the concentration of H+ is the same as the concentration of the initial solution. So 1M of a strong acid will create 1M of H+, meaning the pH is 0.

Kw starts to be important when the concentration of acid (be it strong or weak) is so low, you can't ignore H+ from the water autodissociation.

What is pH of a 10-8 M HCl?
 
  • #3
Okay so concentration is different from whether an acid is strong or weak?

Meaning if you have a diluted acid, then you use Kw, but if you have a concentrated acid, whether it's strong or weak, you have to use Ka to calculate the pH?
 
  • #4
Nishantkumar19 said:
Meaning if you have a diluted acid, then you use Kw, but if you have a concentrated acid, whether it's strong or weak, you have to use Ka to calculate the pH?

More or less.

For a very weak acid it may happen that even if it is not highly diluted you need to take Kw into account.

In general, as I already wrote, it depends on whether the H+ from the water autodissociation can be ignored or not. It can be ignored when it is orders of magnitude lower than the concentration of H+ from the dissociation of acid (note: doesn't matter whether the acid is strong or weak, what matters is how much H+ is produced during dissociation).
 
  • #5
Okay thanks.

Also, in a weak acid, with the value of Ka, is it possible to find Kb? Can the Kw be used to do this for all acids?

Eg. if my value of Ka is 10^-4, is Kb always going to be 10^-10? For every single acid?

And if the H+ concentration is 10^-4, then is the base concentration always going to be 10^-10 as well?

Thanks for your help.
 
  • #6
Nishantkumar19 said:
Eg. if my value of Ka is 10^-4, is Kb always going to be 10^-10? For every single acid?

And if the H+ concentration is 10^-4, then is the base concentration always going to be 10^-10 as well?

Easy to prove with a simple math. Just beware - Ka and Kb must be those of the conjugated pair of acid and base, and in the latter case if the H+ is 10-4 M it is concentration of OH- that is 10-10, not necessarily a concentration of a "base" (whatever it is intended to be).
 
  • #7
Okay thanks
 

1. What is the difference between Kw and Ka in acid dissociation calculations?

Kw and Ka are both constants used in acid dissociation calculations, but they have different meanings. Kw, or the autoionization constant, is the equilibrium constant for the dissociation of water into hydrogen ions (H+) and hydroxide ions (OH-). Ka, or the acid dissociation constant, is the equilibrium constant for the dissociation of an acid into its conjugate base and a hydrogen ion.

2. When should I use Kw in acid dissociation calculations?

Kw is typically used when dealing with solutions that contain both an acid and a base. Since both H+ and OH- ions are present, Kw is used to calculate the concentrations of these ions at equilibrium. It is also useful when determining the pH of neutral solutions, as the concentration of H+ and OH- are equal in neutral solutions.

3. When should I use Ka in acid dissociation calculations?

Ka is used when dealing with a specific acid and its conjugate base. It is used to calculate the concentration of the acid and its conjugate base at equilibrium. It is also used to determine the strength of the acid, as stronger acids will have a higher Ka value.

4. Is it possible to use both Kw and Ka in the same acid dissociation calculation?

Yes, it is possible to use both Kw and Ka in the same calculation. This would be necessary when dealing with a solution that contains both an acid and a base, as well as a weak acid and its conjugate base. By using both constants, you can determine the concentrations of all the ions present in the solution.

5. How do I determine which constant to use in a specific acid dissociation calculation?

The choice between using Kw or Ka in an acid dissociation calculation depends on the specific situation. If you are dealing with a solution that contains both an acid and a base, then Kw would be the appropriate constant to use. If you are dealing with a specific acid and its conjugate base, then Ka would be the appropriate constant. It is important to consider the components of the solution and what information you are trying to determine in order to choose the correct constant.

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