Core Practical 9: Determine Ka for a weak acid

Question 1

sodium hydroxide solution is

Answer 1

an irritant

Question 2

the method for this experiment is

Answer 2

• calibrate the pH meter and set up the datalogger
• pipette 25.0 cm3 of 0.1 mol dm−3 ethanoic acid solution into a 250 cm3 conical flask
• fill a burette with sodium hydroxide solution
• add two drops of phenolphthalein to the acid in the conical flask (the less indicator you, the more accurate the results as indicators are weak acids)
• titrate the acid with the NaOH solution until the mixture just turns pink
• pipette a further 25.0 cm3 of 0.1 mol dm−3 ethanoic acid solution into the 250 cm3 conical flask
• record the pH of this solution

Question 3

the value of Ka is an indication of strength, the larger the value of Ka, the ……………….. acid

Answer 3

the stronger the acid

Question 4

if the pH of the solution is 4.77, then what is [H+]?

Answer 4

pH = -lg[H+]
[H+] = 10 -pH
[H+] = 1.7 x 10-5

Question 5

Calculate a value of Ka for ethanoic acid is the pH of the final solution is 4.77

Answer 5

pH = -lg[H+]
[H+] = 10 -pH
[H+] = 1.7 x 10-5

CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)

Ka = [H+] [A-] / [HA]
Ka = [H+] [CH3COO-] / [CH3COOH]

Ka = [1.7 x10-5] [0.1] / [0.1] = 1.7x10-5

Question 6

Ka is

Answer 6

(the acid ionisation constant)
the equilibrium constant for chemical reactions involving weak acids in aqueous solution. The numerical value of Ka is used to predict the extent of acid dissociation. A large Ka value indicates a stronger acid (more of the acid dissociates) and small Ka value indicates a weaker acid (less of the acid dissociates).

Question 7

sources of uncertainty and how to combat them are:

Answer 7

Sources of uncertainty include inaccuracy of burette readings, and difficulty identifying the exact end-point.
Read glassware from the bottom of the meniscus; use a white tile so you can see the colour change clearly