2).Tip the crystals into large beaker
3).Add 100 cm^3 of distilled water
4).Gently heat over Bunsen flame (careful not to allow boiling)
5).Once solution turns clear/fully dissolved remove from heat
6).Transfer contents to volumetric flask
7).Wash beaker out with distilled water, transfer washings to volumetric flask (repeat 2x)
8).Carefully make up volume in volumetric flask with distilled water (slowly adding until 1cm from graduation mark then use pipette slowly add the water until meniscus just touches graduation mark)
9).Insert bung/stopper and invert several times.
1).Set up resort stand and burette clamp
2).Wash burette out with small amount of hydrochloric acid
3).Run small amount of hydrochloric acid out to ensure jet if full
4).Wash out conical flask in small amount of sodium carbonate
5).Place white tile under burette
6).Invert volumetric flask, using pipette remove 25 cm^3 transfer to conical flask
7).Place conical flask under burette on top of white tile
8).Record the volume reading on the burette before starting the titration (read to nearest 0.05cm^3)
9).Add 3 drops of indicator to sodium carbonate solution – colour orange is observed
10).Add small amounts of acid from burette, swirling conical flash after each addition
11).Stop when the indicator colour changes from orange to pink
12).Record final burette reading and calculate volume of solution run into flask (titre)
13).Record results into results table
Which indicator should I use?
I have the choice of phenolphthalein or methyl-orange. I found in preliminary experiments that half the volume of hydrochloric acid was require to that of when using methyl-orange as an indicator. This was because phenolphalein has an active useful pH range of 8.3 – 10.00 pH, this meant it only recorded a change between these parameters, therefore the solution had not neutralised so accurate calculations could not be made when using this indicator. Therefore I decided upon using methyl-orange as an indicator because it was more relevant to this acid/base titration. Methyl orange is orange in colour then forms a conjugant acid pink/red in colour when neutralised (one drop to much acid added).
How will I ensure my experiment will be fair/accurate?
1).I will make one batch of sodium carbonate solution, therefore the exact same molarity will remain throughout the investigation (limiting final % error)
2).I will use the same volume of sodium carbonate solution therefore I should obtain coherent results
3).I will use a white tile beneath the conical flask so to show colour change most accurately (end titration)
4).I will use the same number of drops of indicator to ensure same degree of colour change
5).I will use the same glass wear each time (so using same calibrations for measurement)
6). I will wash glass wear in corresponding solutions to ensure no cross contamination
7).I will record all measurements to 0.05 cm^3 to limit error
8).I will measure to the meniscus each time
9).I will use a pipette and volumetric flask rather than measuring cylinder because it increases accuracy
10).I will repeat experiments until I get concordant results to with in 0.1cm^3
What safety precautions must I observe?
1).I am working with acid/alkaline so must wear goggles
2).I will also wear a lab coat to protect my clothing
3).I must be careful to slowly pour liquids, so to avoid spillage on myself and work area (which could be potentially hazardous due to nature of chemicals use)
4).I will tie my hair back to stop it hindering my work
(molarity * volume) / 1000 = amount/relative formula mass = number moles
how many moles of Na2CO3 were in each (25cm^3) titration?
(0.1*25)/100 = num moles
From the balance equation (below) I know that one mole of sodium carbonate reacts with two moles of hydrogen chloride therefore from this I can calculate the molarity of the hydrochloric acid solution:
(molarity*volume titrated)/100 = (2*number moles of sodium carbonate)
(molarity*/23.6)/100 = 2*10-3
molarity = (5*10-3*1000)/23.6
molarity = 0.2118644068
By using a burette when triturating I found the experiment more controlled and measurements could be taken o a more accurate degree (compared to measuring cylinders). Also using a pipette it ensured a more accurate method of measuring out amount of sodium carbonate/base. I carried out and initial/rough titration to give me an approximate idea of how much acid was requires neutralising the sodium carbonate; from this it enabled me to be more accurate in the following titrations. When coming near to the point of neutralisation I slowed the burette to adding drop by drop, this allowed coherent results to be taken (with the exception of titration 1 & 3 these results were anomie’s and so were discounted when taking an average titre).
Main sources of error were in the recording/measurement of results as due to the calibration of glass wear (this was an inherent error) i.e. I could only measure to the nearest 0.05 cm^3. also in the measurement of sodium carbonate the balance I used was only accurate to 2 places, it would have been more accurate if it were to measure to more than 2 places. However I made my recordings of measurements for the burette by instead of rounding to the nearest 0.1 to the nearest 0.5 – thus increasing accuracy. When preparing the sodium carbonate solution I transferred washings also into the volumetric flask so also improving my accuracy by ensuring all the sodium carbonate was included – so not affecting the molarity of the calculated solution. Also before the experiment I washed all glass wear with corresponding solutions, therefore stopping cross contamination effecting results – this was also the reason for using distilled water when preparing solution.
Calculating the error within calculations
(Error/volume or amount)*100 = % error
>error in glass wears
Pipette = (0.2/25)*100
Volumetric flask = (0.5/250)*100
>error in measurement
balance = (0.01/7.05)*100
= +/- 0.14%
burette = (0.054/23.6)*100
= +/- 0.21%
>Total error =
pipette + volumetric flask +balance + burette (%error) = +/-%error
= +/- 1.35 % error
How much could this error have affected results?
This is the total maximum +/- error that could have occurred in the experiment, this leaves the calculated molarity possibly inaccurate by
+/- 1.35% of 0.213mol/dm
Maximum molarity =1.0135*0.213 = 0.219 mol/dm^3
Minimum molarity = 0.9865*0.213 = 0.21 mol/dm^3
The inherent errors identified above, as a result of glass wear/measurement could have contributed to the calculated molarity of the unknown concentration of hydrochloric acid being out by as much as 2.75 %. This could result in the true molarity lying between 0.21 mol/dm^3 and 0.219 mol/dm^3. Although there was a possible added error, which could not be accounted for in the figures – this being the point at which the experiment stopped (i.e. the exact point at which the indicator changed from one colour to the other). When researching indicators I came across another indicator which could possibly be better, bromothymol blue which had a useful range of 6.0 –7.0 pH, which distinctly changes in colour from blue to yellow, this may have been a better indicator to have used because it is nearer than methyl-orange to neutral, i.e. more accurate amounts of data collected. I am quite confident in my results, although I have identified errors within the experiment, possible error of +/- 1.35%, this sounds like a large error for a small quantity, however realistically this could be less; this being because of conducting repeats – this minimises the total error. Also I washed glass wear before and between procedures in the substance/solution which was to be used in it, this limited cross contamination which could have effected results from occurring, thus limiting error.