(Chemistry) Ka of unknown acid by titrations.

  • Thread starter Thread starter cshum00
  • Start date Start date
  • Tags Tags
    Acid Chemistry
Click For Summary
SUMMARY

The discussion focuses on determining the acid dissociation constant (Ka) of an unknown weak acid through titration with NaOH. A total of 7.37 mL of NaOH was required for complete titration, with a half-titration volume of 3.69 mL. The measured pH of the solution was 4.7, leading to the conclusion that pKa equals 4.7, and thus Ka is calculated as 2.0 x 10-5 using the formula Ka = 10-pH. The discussion also explores alternative methods for calculating Ka, emphasizing the need for initial concentrations of the acid and its conjugate base.

PREREQUISITES
  • Understanding of weak acid titration principles
  • Familiarity with the Henderson-Hasselbalch equation
  • Knowledge of pH measurement techniques
  • Basic concepts of acid-base equilibrium
NEXT STEPS
  • Study the Henderson-Hasselbalch equation in detail
  • Learn how to calculate concentrations from titration data
  • Explore the concept of monoprotic acids and their dissociation
  • Investigate methods for determining initial concentrations in titrations
USEFUL FOR

Chemistry students, educators, and laboratory technicians involved in acid-base titration experiments and those seeking to understand the calculation of acid dissociation constants.

cshum00
Messages
215
Reaction score
0

Homework Statement


Let's say that i have a unknown weak acid.
  • I prepare 10mL of the weak acid diluted in 25mL of water.
  • I titrate the solution with NaOH. 7.37mL of NaOH was needed for a full titration.
  • Therefore, half-titration of this acid was 3.69mL.
  • I measure the pH of the solution with a pH-meter and the pH was 4.7.

I want to find the Ka of the unknown acid.

Homework Equations


Henderson-Hasselbalch pH formula for a buffer solution is:
pH = pKa + log([base] / [acid])

The Attempt at a Solution


Since the weak acid is monoprotic, the half titration of the solution gives an equal concentration of both acid and conjugate base.

pH = pKa + log(1) = pKa

Does it mean that Ka is the concentration of [H+]?
Ka = 10^-pH = 10^-4.7 = 2.0 e-5?

Edit: Is it possible to calculate in a different way using the given data above?
Ka = [H+][A-]/[HA]
I can calculate the concentration of [H+] from the pH but i don't know the initial concentration of [HA]. And what about the concentration of [A-] ion?
 
Last edited:
Physics news on Phys.org
Yes.
 

Similar threads

Chemistry Are there typos in this snippet causing confusion?
  • · Replies 2 ·
Replies
2
Views
2K
Chemistry How to take into account autoprotolysis in weak base solution?
  • · Replies 8 ·
Replies
8
Views
3K
Chemistry Comparing methods of computing pH at equivalence point in titration
  • · Replies 4 ·
Replies
4
Views
2K
Chemistry Is pH of strong acid strong base titration always 7 at stoichiometric point?
  • · Replies 2 ·
Replies
2
Views
2K
Chemistry Moles of NaOH to Create Buffer Solution pH=6 in 0.42M Ethanoic Acid
  • · Replies 3 ·
Replies
3
Views
3K
Chemistry Book seems to say that we can have weak acid and weak conjugate base
  • · Replies 2 ·
Replies
2
Views
2K
Titration of acetic acid w/ NaOH problem
  • · Replies 2 ·
Replies
2
Views
3K
Chemistry How to calculate pH at equivalence point using this alternative?
  • · Replies 7 ·
Replies
7
Views
2K
Calculating pKa from pH: Understanding Titration
  • · Replies 3 ·
Replies
3
Views
3K
Question about acid and base indicators
  • · Replies 7 ·
Replies
7
Views
7K