Unit 13: Applications of Inorganic Chemistry
Acid or Base?
Blue is M3 - Justify the selection of indicators for the titrations.
Task 2: Values of Ka – Practice work to determine Ka for a weak acid.
Determining the acid dissociation constant Ka for a weak acid
Acid Dissociation Constant For Ethanoic Acid:
Picture :
Chemicals and Equipment:
- 0.1M of ethanoic acid
- 0.1M of sodium hydroxide solution
- Goggles
- 1 burette (50cm3)
- Stand
- pipette (25.0 cm3)
- pipette filler
- 1 beaker (250cm3)
- 1 conical flask (250cm3)
- Distilled water
- pH meter
Method
1. Set up the stand and burette. Once completed, fill the burette with sodium
hydroxide solution using the small glass funnel.
2. Fill the pipette with 25cm3 of ethanoic acid and insert it in the conical flask.
3. Add 2-3 drops of phenolphthalein indicator and mix thoroughly with the solution.
4. Add the sodium hydroxide from the burette to the solution until it turns pink.
5. Repeat the titration until you get consistent results. Calculate the mean of the
results and enter it into the table.
,6. Repeat the practical, but add ethanoic acid to the beaker without the indicator.
7. Add V/2cm3 of sodium hydroxide to the previously measured mean findings.
8. After mixing, calculate the pH of the solution and record it in a table once stable.
Results:
Trial run 1st run 2nd run
(rough) (accurate) (accurate)
Initial reading / cm3 00.00 00.00 00.00
Final readings / cm3 26.6 26.5 26.5
Titre (volume used) / cm3 26.6 26.5 26.5
Mean titre (V) cm3 26.5
V/2 cm3 12.25
Ph of half-neutralised ethanoic acid solution (1st run) 4.3
Ph of half-neutralised ethanoic acid solution (2nd run) 4.6
Ph of half-neutralised ethanoic acid solution ( 3rd run, 4.6
if required)
Average pH value 4.6M
Calculation:
The acid dissociation constant for a weak acid is :
Ka equation is: Ka = [H+] or pKa = pH , but the expression can be rearranged
to give Ka = 10-pH
pKa = 4.6
CH3COOH(aq) <- -> CH3COO- (aq) + H+ (aq)
[CH3COO-](aq) [H+] (aq)
Ka =
[ CH3COOH ](aq
At half neutralising point:
[CH3 COOH] (aq) = [CH3COO- ](aq)
Therefore: Ka = [H+] (aq) = -4.6
Ka = 10-pH
Ka = 10-4.6
pH for ethanoic acid: Ka = 2.5 x 10-5 moldm-3
Therefore: -pH = log [H+] = - 4.6
[H+] = 10-4.6
2.5 x 10-5
=x
[CH3COO-](aq) [H+] (aq)
Ka =
, [ CH3COOH ](aq)
X2
0.1 –x
(2.5 x 10-5)2
0.1 – 2.5 x 10-5)
Ka = 6.25 x 10-9
Conclusion Compare ka with ethonic acid
An acid’s strength can be ascertained by calculating its acids dissociation constant, or
Ka. The degree to which the acid dissociates in water to produce hydronium ions (H+)
and the conjugate base is indicated by the value of Ka. It is a measurement of the
acid’s capacity to provide water molecules with protons. An acid is said to be weaker
when its Ka value is lower. For my experiment, by knowing the Ka value of ethanoic
acid, one may compare the relative strength of several acids and determine the pH of
a weak acid solution.
In this experiment, pH was calibrated, along with the equipment, factors that affected,
and titration readings. A white tile was used, the colours were lighter, less colour was
added than needed air bubbles were present, the pH metre was not calibrated
correctly, and the liquid may have spilt into the funnel at the top of the burette more
than was necessary, which could have changed the Ka value reading leading to mine
being higher than necessary. I needed to swirl the mixture repeatedly after every
volume of sodium hydroxide was dropped to ensure that it was evenly blended and
colour was apparent. However, that was not the case, even when I had reached the
readings when the colour had changed, I continued to see if there would be excess
colour change if more was added, which there was not. Additionally, in my doing this,
the pH metre bulb was not as inside measuring the hydrochloric acid as it should have
been, and extra sodium hydroxide was poured into the beaker after crossing the
meniscus line, which might have caused inaccuracy in the measurements. The pipette
filler exceeded the 25cm3 using the pipettes. The experiment was inaccurate
because the amount of phenolphthalein seemed to be less than the required drops by
one time or not by all. Because the colours might vary quickly based on whether there
was enough to generate a colour change, the endpoint mistake may experience these
kinds of interpretation problems. Human error and occasionally equipment
malfunctions can cause the reading from the burette to vary since little drops from an
, open tap might cause the measurements to change. The published Ka value for
ethanoic acid is 1.8 x 10-5; for ethanoic acid, Ka = 6.25 x 10-9 moldm-3.
Task 3: Buffer Solutions – Practical work to demonstrate and discuss the
action of buffers
Chemical and equipment:
- Sodium hydroxide
- Hydrochloric acid
- Sodium ethanoate solutions
- Ethanoic acid
- Beaker (150ml)
- pH meter
- distilled water
Method:
1. Combine 10ml of 2.0 mol dm-3 ethanoic acid and 10ml of 1.0 mol dm-3 sodium
ethanoate in a 150ml beaker.
2. After cleaning the pH metre with distilled water, insert the tip into the solution.
3. Once the pH metre reading is stable, record the values in a table.
4. Follow the table layout and add hydrochloric acid and sodium hydroxide to continue
the procedure.
Results:
The Ka of ethanoic acid is 1.8 x 10-5
pH of the acidic buffer is 3.80
Volume of Hydrochloric pH Reading for HCL pH Readings for NaOH
Acid or Base?
Blue is M3 - Justify the selection of indicators for the titrations.
Task 2: Values of Ka – Practice work to determine Ka for a weak acid.
Determining the acid dissociation constant Ka for a weak acid
Acid Dissociation Constant For Ethanoic Acid:
Picture :
Chemicals and Equipment:
- 0.1M of ethanoic acid
- 0.1M of sodium hydroxide solution
- Goggles
- 1 burette (50cm3)
- Stand
- pipette (25.0 cm3)
- pipette filler
- 1 beaker (250cm3)
- 1 conical flask (250cm3)
- Distilled water
- pH meter
Method
1. Set up the stand and burette. Once completed, fill the burette with sodium
hydroxide solution using the small glass funnel.
2. Fill the pipette with 25cm3 of ethanoic acid and insert it in the conical flask.
3. Add 2-3 drops of phenolphthalein indicator and mix thoroughly with the solution.
4. Add the sodium hydroxide from the burette to the solution until it turns pink.
5. Repeat the titration until you get consistent results. Calculate the mean of the
results and enter it into the table.
,6. Repeat the practical, but add ethanoic acid to the beaker without the indicator.
7. Add V/2cm3 of sodium hydroxide to the previously measured mean findings.
8. After mixing, calculate the pH of the solution and record it in a table once stable.
Results:
Trial run 1st run 2nd run
(rough) (accurate) (accurate)
Initial reading / cm3 00.00 00.00 00.00
Final readings / cm3 26.6 26.5 26.5
Titre (volume used) / cm3 26.6 26.5 26.5
Mean titre (V) cm3 26.5
V/2 cm3 12.25
Ph of half-neutralised ethanoic acid solution (1st run) 4.3
Ph of half-neutralised ethanoic acid solution (2nd run) 4.6
Ph of half-neutralised ethanoic acid solution ( 3rd run, 4.6
if required)
Average pH value 4.6M
Calculation:
The acid dissociation constant for a weak acid is :
Ka equation is: Ka = [H+] or pKa = pH , but the expression can be rearranged
to give Ka = 10-pH
pKa = 4.6
CH3COOH(aq) <- -> CH3COO- (aq) + H+ (aq)
[CH3COO-](aq) [H+] (aq)
Ka =
[ CH3COOH ](aq
At half neutralising point:
[CH3 COOH] (aq) = [CH3COO- ](aq)
Therefore: Ka = [H+] (aq) = -4.6
Ka = 10-pH
Ka = 10-4.6
pH for ethanoic acid: Ka = 2.5 x 10-5 moldm-3
Therefore: -pH = log [H+] = - 4.6
[H+] = 10-4.6
2.5 x 10-5
=x
[CH3COO-](aq) [H+] (aq)
Ka =
, [ CH3COOH ](aq)
X2
0.1 –x
(2.5 x 10-5)2
0.1 – 2.5 x 10-5)
Ka = 6.25 x 10-9
Conclusion Compare ka with ethonic acid
An acid’s strength can be ascertained by calculating its acids dissociation constant, or
Ka. The degree to which the acid dissociates in water to produce hydronium ions (H+)
and the conjugate base is indicated by the value of Ka. It is a measurement of the
acid’s capacity to provide water molecules with protons. An acid is said to be weaker
when its Ka value is lower. For my experiment, by knowing the Ka value of ethanoic
acid, one may compare the relative strength of several acids and determine the pH of
a weak acid solution.
In this experiment, pH was calibrated, along with the equipment, factors that affected,
and titration readings. A white tile was used, the colours were lighter, less colour was
added than needed air bubbles were present, the pH metre was not calibrated
correctly, and the liquid may have spilt into the funnel at the top of the burette more
than was necessary, which could have changed the Ka value reading leading to mine
being higher than necessary. I needed to swirl the mixture repeatedly after every
volume of sodium hydroxide was dropped to ensure that it was evenly blended and
colour was apparent. However, that was not the case, even when I had reached the
readings when the colour had changed, I continued to see if there would be excess
colour change if more was added, which there was not. Additionally, in my doing this,
the pH metre bulb was not as inside measuring the hydrochloric acid as it should have
been, and extra sodium hydroxide was poured into the beaker after crossing the
meniscus line, which might have caused inaccuracy in the measurements. The pipette
filler exceeded the 25cm3 using the pipettes. The experiment was inaccurate
because the amount of phenolphthalein seemed to be less than the required drops by
one time or not by all. Because the colours might vary quickly based on whether there
was enough to generate a colour change, the endpoint mistake may experience these
kinds of interpretation problems. Human error and occasionally equipment
malfunctions can cause the reading from the burette to vary since little drops from an
, open tap might cause the measurements to change. The published Ka value for
ethanoic acid is 1.8 x 10-5; for ethanoic acid, Ka = 6.25 x 10-9 moldm-3.
Task 3: Buffer Solutions – Practical work to demonstrate and discuss the
action of buffers
Chemical and equipment:
- Sodium hydroxide
- Hydrochloric acid
- Sodium ethanoate solutions
- Ethanoic acid
- Beaker (150ml)
- pH meter
- distilled water
Method:
1. Combine 10ml of 2.0 mol dm-3 ethanoic acid and 10ml of 1.0 mol dm-3 sodium
ethanoate in a 150ml beaker.
2. After cleaning the pH metre with distilled water, insert the tip into the solution.
3. Once the pH metre reading is stable, record the values in a table.
4. Follow the table layout and add hydrochloric acid and sodium hydroxide to continue
the procedure.
Results:
The Ka of ethanoic acid is 1.8 x 10-5
pH of the acidic buffer is 3.80
Volume of Hydrochloric pH Reading for HCL pH Readings for NaOH
