‘For a titration between a strong acid and a weak alkali, methyl orange is used as the indicator.
For a titration between a weak acid and a strong alkali, phenolphthalein is used as the indicator.
For a titration between a strong acid and a strong alkali, either methyl orange or phenolphthalein can be used as the indicator, although methyl orange is usually chosen.
For a titration between a weak acid and a weak alkali, no indicator is suitable, and a pH meter, conductivity meter or temperature probe has to be used.’
As this titration is between a strong acid and a weak alkali, I can see from this information that the correct indicator to use is methyl orange. I will use this indicator to show the end point of this titration.
Preliminary Preparations:
A sample of both acid and alkali solutions are put into two separate test tubes. A few drops of methyl orange indicator are added to each test tube and the colour changes are observed. This ensures that correct solutions are present because in acid the indicator will turn red and in alkali it will remain orange.
I plan to carry out a trial titration followed by further accurate titration. The first result will be ignored as it is highly likely to be anomalous.
Method:
- The sodium carbonate solution is poured into the burette using a funnel. To ensure there are no bubbles some of the solution is allowed to run through the burette.
- 25cm³ of the sulphuric acid solution is measured out using a pipette. This is then placed into a conical flask.
- 10 drops of methyl orange indicator are added to the acid.
- The flask is placed under the burette on a white tile so that changes in colour will be more noticeable.
- The starting point on the burette is noted and the sodium carbonate solution is run into the acid until the indicator turns from red to pale yellow.
- The volume of sodium carbonate used is measured at the bottom of the meniscus and noted.
- The titration is repeated until concordant results are achieved.
Justification:
I believe that my experiment will provide precise and reliable results due to the following factors:
- It is possible to read the burette to two decimal places, therefore the readings will be very accurate.
- Repeat titrations will enhance the accuracy of the results.
- Using the tare facility of the balance will ensure that an accurate amount of sodium carbonate is used.
- Use of the pipette ensures that an exact amount of acid is used.
- Conditions are the same for each titration (same amount of acid and indicator used every time; clean glassware used; same solutions used).
- All used glassware will be thoroughly washed with distilled water. It would be preferable if new glassware could be used for each titration but this would be impractical.
Risk Assessment:
- Acid is an irritant so it is important that eye protection is worn at all times.
- Any spillages must be dealt with safely and promptly using the correct solutions.
- Hands must be washed thoroughly after handling these substances.
- Glassware should be handled with caution.
Sources:
In devising this plan I used the following sources for information:
- Modified Exemplar Mark Schemes – 2001/2 (Analysing an Acidic Solution)
- ‘Letts AS Chemistry’ by Rob Ritchie
- Salters Advanced Chemistry – ‘Chemical Ideas’
- The Cambridge Encyclopaedia
ANALYSING EVIDENCE AND DRAWING CONCLUSIONS
Results:
I will use the average titre in my calculations to promote a higher level of accuracy.
(I will ignore the first reading in calculating the average as it appears to be anomalous)
Calculating the Concentration of Na2CO3 Used:
Although the concentration was calculated when making up the sodium carbonate solution, mistakes may have been made or the equipment used may not have been ideal. Therefore, it is necessary to find the exact concentration from the results obtained.
The information necessary for these calculations are:
Na2CO3(aq) + H2SO4(aq) → Na2SO4(aq) + H2O(l) + CO2(g)
-
The concentration C1 and the reacting volume V1 of Na2CO3(aq).
-
The concentration C2 and the reacting volume V2 of H2SO4(aq).
From the titration results, the amount of Na2CO3 (in mol) can be calculated:
Amount of Na2CO3 = C1 x = 0.100 x = 0.00247 mol
Calculating the Concentration of H2SO4:
From the equation, the amount of H2SO4 (in mol) can be determined:
Na2CO3(aq) + H2SO4(aq) → Na2SO4(aq) + H2O(l) + CO2(g)
1 mol 1 mol (balancing numbers)
Therefore, 0.00247 mol Na2CO3 reacts with 0.00247 mol H2SO4
Amount of H2SO4 that reacted = 0.00247 mol
The concentration (in mol dmˉ³) of H2SO4 can be calculated by scaling to 1000cm³:
25cm³ H2SO4(aq) contains 0.00247 mol H2SO4
1cm³ H2SO4(aq) contains mol H2SO4
1dm³ (1000cm³) H2SO4(aq) contains x 1000 = 0.099 mol H2SO4 (to 3 d.p.)
I have thus calculated that the concentration of H2SO4 is 0.099 mol dmˉ³. However, my scientific knowledge tells me that the actual concentration is more likely to be 0.1 and certain procedural errors have slightly altered the figures used in my calculations.
EVALUATING EVIDENCE AND PROCEDURES
I believe that my results were quite reliable although my first reading was not entirely concordant with the others. Although the difference was only 0.3-0.4 cm³ I have not included this result in my calculations to find the average as the repeat titrations produced much more concurrent results and the average would be more reliable with the anomaly excluded.
Calculating the Percentage Error:
Percentage error is one of many methods used to express your level of confidence in an average of many measurements. It is used when a known or accepted value for the measurement is available.
The following formula is used to determine the percentage error:
Percentage Error of Measurement 1:
24.10cm³ Percentage Error = 3.6%
Percentage Error of Measurement 2:
24.75cm³ Percentage Error = 1%
Percentage Error of Measurement 3:
24.65cm³ Percentage Error = 1.4%
Percentage Error of Measurement 4, 5, 6 and Average:
24.70cm³ Percentage Error = 1.2%
With a percentage error of 1.2% I feel that I am justified in saying that my results were very accurate under the circumstances and with the equipment available. I have ignored the first reading as it is not concordant with the other measurements taken.
I think the common source of error throughout the experiment would be my misjudgement of the endpoint of titration. I seem to have ‘stopped short’ of the endpoint repetitively and had I been totally confident of this, my results would have been far more accurate.
The equipment used in the experiment was the best available to me for use. I believe I utilised all the facilities well and took appropriate additional precautions to ensure the experiment would be a success.
Evaluating Equipment:
- Burette - This is ideal for titration experiments as the amount of solution allowed to run through can be carefully controlled with use of the tap. Towards the end of titration the solution can be added drop by drop. This way the exact endpoint can be reached. Also, the burette can be read to 2 decimal places. This promotes a higher level of accuracy than most other equipment. This method could be improved by using apparatus that you can read to more than 2 decimal places but the burette is very adequate.
- Pipette - this is a more accurate method of measuring out liquid than, for example, a measuring cylinder. This is because it is very thin and clearly marked at 25cm³. More acid than 25cm³ would have resulted in more sodium carbonate solution being used and less acid than 25cm³ would result in less sodium carbonate being used. The pipette was ideal for this and the only improvement would be if it were even thinner.
- Glassware - As it was impractical to keep using clean glassware, I ensured that the conical flask was washed thoroughly with distilled water and dried after each repeat. This was necessary because the experiment wouldn’t be a fair test if there was a change in the amount of acid used or if the acid was diluted.
- Volumetric Flask - Like the pipette, the thin neck and clear 250cm³ mark enabled the sodium carbonate solution to be made with a high level of accuracy. It was also easy to shake so that the sodium carbonate would dissolve completely. It was ideal for this procedure.
- White Tile - Placing the white tile underneath the conical flask made the colour change far easier to view and so the endpoint could be found with greater accuracy. This cannot be blamed for personal fault.
- Balance - Use of the tare facility of the balance enabled a higher level of accuracy when measuring out the solid sodium carbonate.
- Weighing Bottle - This was the best method available to us for weighing out the sodium carbonate. However, when transferring it from the weighing bottle to the volumetric flask some of the fine powder may have stuck to the weighing bottle or simply blown away. This would have altered the concentration of the alkali solution and therefore affected the overall results.
- Funnel - This was used to promote ease of pouring from the volumetric flask to the burette.
Evaluating Procedures:
-
I think I should have put the acid into the burette and a set amount of alkali into the flask instead of the other way round. The Na2CO3 solution could have evaporated, leaving solid Na2CO3 blocking the burette and an increase in the concentration of Na2CO3.
- Repeats - I believe that enough repeat titrations were performed to achieve concordant results. The accuracy of the results was elevated by the repeats and a suitable average was attained. Were there no or few repeats then it would be impossible to single out the anomalies and the conclusion would be incorrect.
- Taking Measurements - I was very careful when reading from the burette and in recording the measurements. It would have been quite difficult to make a mistake in this as the tap is turned off and the solution is steady at the mark so the exact position of the bottom of the meniscus can be seen clearly.
- Procedural Errors - I think that I was unclear as to the point of complete neutralisation. I obviously stopped the titration too early and I think this is why my results are a little under what was expected.
Greatest Sources of Error:
Making the solution of Na2CO3 homogeneous: even when the solution was shaken in the volumetric flask, the concentration of ions may have been different in the neck of the flask than in the bottom. This would have affected the results as the solution may have not had an accurate concentration.
The end point: It was fairly difficult to tell the end point as after the solution turned yellow, it gradually became pink again. This was probably because carbon dioxide in the air made the solution more acidic. This source of error could be taken away by performing the experiment under carbon dioxide free conditions (if the flask were sealed, the results would be more accurate).
In conclusion, I believe that effect of the faults found in the equipment and method is almost negligible when it comes to analysing the accuracy of the results. The greatest error was in my own inaccuracy in judging the endpoint when equivalence was reached. Were I more clear on this, I am confident that my results would have showed almost, if not exactly, 25cm³ of the sodium carbonate solution used to titrate 25cm³ of the sulphuric acid solution.
An Investigation to Find the Accurate Concentration of an Acid
Elizabeth Dodwell