When to Use Kw & Ka for Acid Dissociation Calculations

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Discussion Overview

The discussion revolves around the appropriate use of the acid dissociation constant (Ka) and the ion product of water (Kw) in calculating the pH of acid solutions. Participants explore the distinctions between strong and weak acids, the implications of concentration on these calculations, and the relationships between Ka, Kb, and Kw.

Discussion Character

  • Technical explanation
  • Conceptual clarification
  • Debate/contested
  • Mathematical reasoning

Main Points Raised

  • Some participants note that for strong acids, the concentration of H+ ions equals the concentration of the acid, leading to straightforward pH calculations.
  • Others argue that weak acids do not fully dissociate, necessitating the use of Ka to determine H+ concentration and pH.
  • There is a discussion about when Kw can be used, particularly in cases of very dilute acids where H+ from water autodissociation cannot be ignored.
  • One participant questions whether the use of Kw applies only to water or can also be relevant for some acids.
  • Participants discuss the relationship between Ka and Kb, with one asking if Kb can always be derived from Ka using Kw, and whether this holds true for all acids.
  • There is a clarification that Ka and Kb must correspond to conjugate acid-base pairs, and that the concentration of OH- relates to H+ concentration in specific ways.

Areas of Agreement / Disagreement

Participants express varying views on the conditions under which Kw and Ka should be used, with no clear consensus reached. The discussion remains unresolved regarding the precise thresholds for using each constant and the implications for different acid concentrations.

Contextual Notes

Limitations include the lack of clarity on specific concentration thresholds for applying Kw versus Ka, and the dependence on definitions of strong and weak acids. The discussion also reflects uncertainty about the generalizability of relationships between Ka, Kb, and concentrations of H+ and OH-.

Nishantkumar19
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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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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?
 
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?
 
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).
 
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.
 
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).
 
Okay thanks
 

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