Safety
Justification of Quantities of Materials being used
Due to the fact that the concentration of the sulphuric acid is between 0.05 and 0.15 moldm-3 I will make the sodium carbonate 0.1 moldm-3 for convenience. Above I have stated that I will use 2.65g of sodium carbonate. I have derived this from the following calculations:
Molecular Mass Na2CO3 = (23 x 2) + 12 + (16 x 3)
= 106
Therefore 1 mole of Na2CO3 is 106g
I require 0.1 mole so = 106 / 10
= 10.6g
Therefore 10.6g will be dissolved in 0.25dm3 so there will be ¼ the amount of sodium carbonate = ¼ x 10.6g
= 2.65g
For this reason I will be using 2.65g of sodium carbonate. I will also only be adding a few drops of methyl orange indicator so that I can clearly observe any colour changes. The volume of sulphuric acid needed to neutralise the sodium carbonate will obviously depend on its concentration, so I can not at this point determine what volume of sulphuric acid I need.
Justification of Method
I have chosen to use a burette as it has a large resolution- i.e. it can be read very accurately (0.05cm3). This will enable me to obtain more accurate results and therefore to calculate a more reliable reading for the concentration of the sulphuric acid.
For the same purpose I have used the volumetric flask as I can accurately make up a concentration of sodium carbonate/ water solution. The instruments I am using are precision instruments which will allow me to measure things more accurately.
The fact that I am rinsing things I have placed the sodium carbonate in and then adding it to the solution will also make my results more accurate as I will not be losing any of the sodium carbonate and therefore will not be making a slightly lower concentration. Again when I calculate the concentration of the sulphuric acid I will get a more reliable answer.
In these ways I believe that when I carry out my experiment I will be ensuring the test is fair and therefore I will obtain more accurate and reliable results.
Details of Sources Used
The sources I have used in writing this plan up were worksheets Elements of Life 2.1 which assisted me in writing the procedure to use during this experiment. The handouts given out by the teacher were also used to discover information of the correct type of indicator to use and the hazards of some chemicals.
Results
Mass of sodium carbonate used = 2.65g
Indicator colour change = Turns from yellow to pink
Analysis
Molecular Mass Na2CO3= 106
Therefore 1 mole = 106, so 0.1 moles = 10.6g
As we are only using ¼ dm3 we require ¼ of 10.6g, so:
Mass of Na2CO3 used = 10.6 / 4
= 2.65g
Concentration of
Na2CO3 solution = Volume (in that titration) / (Total) Volume
= 0.025 / 0.25
= 0.1 moldm-3
Moles Na2CO3 used = Concentration x Volume (in that titration)
= 0.1 x (25/1000)
= 0.0025 moles
Na2CO3 + H2SO4 Na2SO4 + H2O + CO2
Moles H2SO4 used = Moles Na2CO3 used (as seen in the equation)
= 0.0025 moles
So if 26.875cm3 of H2SO4 contains 0.0025 moles of H2SO4 then
Molarity of H2SO4 = (Moles / Volume) x 1000
= (0.0025 / 26.875) x 1000
= 0.093 M (3 d.p.)
So the concentration of the Sulphuric Acid Solution was 0.093 M
Evaluation
From my results I obtained 2 anomalous results, including the rough titration I carried out. However titrations 2 and 3 were 0.05cm 3 within each other, which shows the experiment, was very accurate.
The results I obtained, including the anomalous results, were very precise- i.e. they were read to 2 decimal places. The results are also reliable in that they all show that more sulphuric acid was needed to neutralise the sodium carbonate, and these figures were all in the same region. The proximity of each value to one another also enforces the fact that the experiment was carried out very precisely and accurately, and thus produced very reliable results.
The apparatus I used was also very reliable- the burette being able to read up to 2 decimal places of a cm3. The other apparatus I used also gave precise readings with small percentage errors. Below I have calculated the percentage errors of my experiment, using the formula for percentage error shown below
Percentage error = (error x 100) / reading
Percentage error for a
250cm3 volumetric flask = (0.2 x 100) / 250
= 0.08 %
Percentage error for a
25cm3 pipette = (0.06 x 100) / 25
= 0.24 %
Percentage error for a
26.875cm3 burette reading = (0.05 x 100) / 26.875
= 0.19 % (2 d.p.)
Percentage error for scales
of 2.65g reading = (0.005 x 100) / 2.65
= 0.19 % (2 d.p.)
Combined percentage error for the complete titration is therefore 0.69 % (2 d.p.). This is very small for an overall percentage error for the whole titration and so I can say that the apparatus used was very accurate. Ideally the combined percentage error would equal zero, but the very small percentage error I obtained has no significant impact on the experiment. I therefore again say that I think the experiment was very accurate and reliable.
The sources of error in my procedure were few. The apparatus was extremely accurate, and so the only real source of error I can see is human error. The main human errors would be observing the colour change. If during one titration I stopped adding the acid when the colour of the indicator was darker, I would clearly have added more acid. This would cause me to overestimate the concentration of the sulphuric acid- that is I would think the acid was less concentrated. However once again I must say that I obtained two titrations within 0.05cm3 of each other.
Improving the procedure would be difficult as the apparatus I used was all very accurate. However if I did alter something it would be to use a digital indicator meter which measures the ph level of the acid/alkali solution. By doing this I would not overshoot when adding the acid to the alkali (as I would add acid until the meter read 7) and therefore obtain more accurate readings and less anomalous ones. Apart from this I can think of no other real improvements to the procedure, apart from reading the burette at eye level to prevent reading a higher or lower number.
