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Position the burette in the clamp-stand and ensure that it is below eye level (this can be done by placing it on a stool instead of the desk). Using the funnel and glass beaker containing the 0.1 mol dm3 sodium carbonate solution fill the burette until it is just over the 50 cm3 mark. Now place the beaker with the acid solution in it beneath the burette and whilst keeping your eye-level opposite the meniscus open the tap slowly until the base of the meniscus is on the 25cm3 mark. Fill a burette with 0.1 mol dm3 sodium carbonate solution. Make sure that the burette jet is also full of solution.
- Record the volume reading in the burette before starting the titration. Then add sodium carbonate solution, in small volumes, to the acid solution in the conical flask. Swirl the flask after each addition.
- The four drops of Methyl Orange turns the liquid orange in a alkaline solution. The beaker is placed on a white tile so that when total neutralisation occurs the peach colour can be easily seen. Run in small volumes of sodium carbonate solution until the first occurrence of a permanent peach colour in the titration mixture. This is the end point.
- Record the final reading by reading off the burette at eye-level from the base of the meniscus and calculate the volume of sodium carbonate solution that has been used.
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The first attempt will be a rough titration; you will have gone beyond the end-point and added more sodium hydroxide than is needed to react with all the acid in the flask. You should have a general idea of what the end-point is. Do several more titration’s until you record four volumes that are within 0.1cm3 of each other. When you near the end points of the titration’s, add the sodium hydroxide more carefully, adding only one drop at a time until the peach colour remains permanent, recording the results
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From the titres which are within 0.1cm3 of each other, calculate an average volume of sodium hydroxide solution used in the titration flask.
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Combine this average volume with the concentration of the sodium carbonate solution to calculate the amount (in moles) of sodium carbonate, Na2CO3, which just reacts with the acid in the titration flask.
The acid in the sample is sulphuric acid. The equation for the reaction in this titration is therefore:
H2SO4 (aq) + Na2Co3 (aq) → Na2SO4 (aq) + H2O (l) + CO2 (g)
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Write down the concentration of the sulphuric acid in units of mol dm3.
Apparatus set-up diagram:-
Safety precautions
In this experiment we are working with both acids and alkalis, which are corrosives that could cause serious harm. The unknown concentration of the acid is also potentially dangerous…
Safety precautions to be taken
- Goggles must be worn at all times.
- All acids must be poured below eye level so to allow any spillage to reach eyes, as acid is an irritant.
- Lab coats must be worn and all relevant safety equipment word. I.e. gloves
- Long hair must be tied back.
- Do not mouth pipette, always use the filler bulb, so as not to risk sucking acid into your mouth, and risking irritant/corrosive damage.
- Ensure to wash any skin or part of the body if they come in contact with either acid or alkali as both are irritant and/or corrosive.
Implementing:
I followed closely the procedure as described in my plan and have paid attention to the safety procedures is have aforementioned, therefore I have encountered no problems. Below are my accurate results, I recorded 6 results in total the first being my rough titre, in which I surpassed the neutralisation point. I had 1 anomaly as I did not meet the 0.1 differential maximum gap and the other 5 results all fall with the 0.10 cm3 accurate parameters, which were set earlier.
Results table:-
(24.20+24.20+24.20+24.20+24.30) = 24.22 cm3 (mean average)
5
Average titre= 24.22 cm3
Analysing evidence and drawing conclusion:
Equation for the reaction
H2SO4 (aq) + Na2CO3 (aq) → Na2SO4 (aq) + H2O (l) + CO2 (g)
i.e.- 1 mol. Neutralises 1 mol.
To determine the concentration of the sodium carbonate solution I will first need to calculate the no. of moles in the solution. To do this I will use the following formula:
Number of moles = mass
Relative atomic mass
No. Moles = 2.65 = 0.025 mol.
106
Therefore the number of moles that are present in the 25cm3 solution is 0.0025 moles.
The equation clearly shows the 1:1 moles ratio between H2SO4 and Na2CO3
25cm3 Na2CO3 neutralises .0025mol. H2SO4
24.22cm3 H2SO4 solution contains 0.0025 mols
Concentration of H2SO4 =No. moles x 1000
Volume
1dm3/1000cm3 H2SO4 solution contains (0.0025 x 1000)
24.22 = 0.1032 mol
Concentration of H2SO4 = 0.10mol dm-3
From this we can conclude that 24.22 cm3 of 0.1 mol H2SO4 neutralises 25cm3 of 0.1 mol. Na2CO3
Evaluating Evidence
In the main the results I obtained were very accurate, as I followed a well planed method and used very precise equipment. The ‘test run’ provided a useful estimate in which we can use to calculate the neutralisation point of the acid/alkali solution. My accurate average is calculated from the final 5 results, any anomalous results.
There was on slightly anomalous result, which could have been the result of poor use of the available equipment as all of the instruments had a high degree of accuracy. The most likely cause is that of not judging the point of neutralisation correctly as the solution had not turned completely colourless or had missed the same point. The Burette scale could also be misinterpreted as you had to almost read backwards and down the scale, which could be quite awkward. Also since everybody has different levels of visibility judging colour changes can be subtly effected and trying to judge the bottom of meniscus can also result in different readings. Other issues were of contamination in the form of incorrect use of the pipette filer bulb which had often been filled with the alkali solution; this could seriously affect any results obtained.
The following calculations show how errors in the accuracy of the equipment used may have affected my results:
Percentage error = Error x 100
Reading
Pipette error = 0.06 x 100 =0.24%
25
Burette error = 0.05 x 100 = 0.34%
14.575
Volumetric flask error = 0.02 x 100 = 0.08%
250
Although the above results only attribute to a small percentage of the total error they still contribute to the total result. An error in the volumetric flask can effect the concentration of the sodium carbonate, pipette errors can attribute towards the amount of Na2CO3 being neutralised. Lastly the burette errors affect the volume of acid which is needed to neutralise the Na2SO4.
If I was to repeat this investigation then I feel there are some minor things than I would change, so as to improve the accuracy of my retainable results. The amount of obtained results would increase from 5 to 10 or possibly more, this would give a more accurate average and show up any anomalies. Given the accuracy of the pipette and burette the dropper method used to apply the methyl orange indicator could be improved by such means as using a measuring pipette. Also the change in colour technique relies heavily upon any persons eyesight, an electronic ph metre would absolve this issue any would give us accuracy to the 0.01pH. Contamination of the equipment is a big issue, a possible solution is to have a new set of apparatus for each of the experiments, I realise this is an impossibility as equipment is expensive. Another alternate is to have a form or dishwasher. If I implicated these changes I feel the experiment will run much smoother.
