Apparatus and Equipment used for the two titrations carried out
Chemicals:
Ethanoic Acid (CH3COOH(aq))
Sulphuric Acid (H2SO4(aq)) of 0.1 mol dm-3 concentration
Sodium Hydroxide (NaOH(aq)) of unknown concentration
Bromothymol Blue Indicator
Laboratory Equipment:
Safety goggles
Laboratory coat
2 x 2cm3 micro pipette
1 x 1cm3 micro pipette
1 x 12 Well plate
3 x 25cm3 Beakers
Marker to label beakers
1 x White card
1 x Indicator pipette
3 x stirrers
1 x Clamp to hold pipette
1 x Syringe with tube
Method
For safety precautions, lab coat and goggles must go on before any experiments take place. Once labelling the beakers with the marker pen, the sulphuric acid, vinegar, and sodium hydroxide were poured into each. As shown in the above diagram, the titration equipment was set up, upon a white card so that the colour change at the neutralisation stage is observationally clear.
For the first titration, where the sulphuric acid was titrated against the sodium hydroxide the following was carried out. Using the 2cm3 micropipette, 1cm3 of sodium hydroxide was transferred into a well on the well plate, ensuring the meniscus was on the 0cm3 mark. This was carried out four times, into 3 more wells. This was done so that an average of the titrations can be taken for the calculations. To each of these, 3 drops of bromothymol blue was added using the indicator pipette and stirred to mix the solution. This is used as an indicator for the colour change as and when the neutralisation stage occurs. The sulphuric acid was then carefully drawn up into the second 2cm3 micropipette and attached to the clamp on the well plate. In the first well of sodium hydroxide, the sulphuric acid was titrated drop by drop. This first titration would be the flush, so that there was a guide for the next three. This solution was continuously stirred especially when there was sign of slight colour change. Once the blue colour of the indicator within the solution significantly changed to green, the titration was ceased and the reading was taken and recorded. This was then carried out three more times in the other 3 wells with sodium hydroxide making sure readings are 0.02cm3 of each other.
The second titration involved titrating the sodium hydroxide against the ethanoic acid. The 1cm3 micropipette was used to draw up the ethanoic acid and fill each of four new wells. The Bromothymol blue indicator was then added and stirred in. The sodium hydroxide was drawn up using the first 2cm3 micropipette and once attaching it to the clamp above a single well, it was carefully titrated drop by drop. This first was taken as the flush, which was an estimate for the other three to find the equivalence point. This in turn, saved time in the next three titrations where the sodium hydroxide was flushed through until it reached close to the flush point and then titrated drop by drop to be more accurate. Therefore this titration was repeated three times, to find an average in the three runs which were to be within 0.02cm3 accuracy. All readings were recorded within a table.
Results Achieved
Titration One: Sulphuric Acid against Sodium Hydroxide
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
Indicator: Bromothymol Blue
Colour change: blue to yellow (green at equivalence point)
Table 1: Amount of sulphuric acid that was needed to neutralise 1cm3 of sodium hydroxide
Calculating concentration of sodium hydroxide from the above results:
Average Volume of H2SO4 = 0.39* + 0.40* + 0.41*
3
= 0.40cm3
Concentration of H2SO4 = 0.1 mol dm3
Volume of NaOH = 1cm3
Concentration of NaOH = ?
∴Moles of H2SO4 = concentration of H2SO4 X average volume of H2SO4 in dm3
= 0.1 X 0.40
1000
= 0.00004 moles
∴Moles of NaOH can be calculated from the balanced equation:
2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
Ratio: 2 : 1
Moles: 0.00004 x 2 : 0.00004
∴Moles of NaOH = 0.00004 x 2
= 0.00008 moles
∴Concentration of NaOH = Moles of NaOH .
Volume of NaOH in dm3
= 0.00008
1/1000
= 0.08 mol dm-3
= 0.1 mol dm-3 (to 1 decimal place)
Titration Two: Sodium Hydroxide against Ethanoic Acid
CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l)
Indicator: Bromothymol Blue
Colour change: yellow to blue (green at equivalence point)
Table 2: Amount of sodium hydroxide that was needed to neutralise 1cm3 of ethanoic acid
Calculating concentration of Ethanoic Acid from the above results:
Average Volume of NaOH = 1.26* + 1.24* + 1.25*
3
= 1.25cm3
Concentration of NaOH = 0.08 mol dm-3 (for accuracy)
Volume of CH3COOH = 1cm3
Concentration of CH3COOH = ?
∴Moles of NaOH = concentration of NaOH X average volume of NaOH in dm3
= 0.08 X 1.25
1000
= 0.0001 moles
∴Moles of CH3COOH can be calculated from the balanced equation:
CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l)
Ratio: 1 : 1
Moles: 0.0001 : 0.0001
∴Moles of CH3COOH = 0.0001 moles
∴Concentration of CH3COOH= Moles of CH3COOH .
Volume of CH3COOH in dm3
= 0.0001
1/1000
= 0.1 mol dm-3
Discussion
Being provided ethanoic acid together with sulphuric acid of a known concentration and sodium hydroxide of an unknown concentration, there was no direct possible way to find the concentration of the ethanoic acid. This lead to carrying out two separate titrations, which is called a back-titration, as we must find the concentration of the sodium hydroxide to then find the concentration of the ethanoic acid.
Using the sulphuric acid to titrate the sodium hydroxide, gave a clear set of results recorded in table 1. These results helped in calculating the concentration of the sodium hydroxide. The amount of sulphuric acid required to neutralise the sodium hydroxide was at the point 0.40cm3, this calculated the concentration to be 0.1mol dm-3.
The sodium hydroxide was then used to titrate the ethanoic acid. The results of this were recorded in table 2. The volume of sodium hydroxide required in this neutralisation was 1.25cm3, this led to the calculation of the concentration of ethanoic acid, which also was 0.1mol dm-3.
The choice to use the indicator bromothymol blue was decided as ethanoic acid is a strong acid and the sodium hydroxide is a strong base. This meant that at the neutralisation stage the pH of the solution will be close to pH7, which is where the approximate pH range for colour change is for this indicator.
The concentrations were not exact to the decimal place of 0.1mol dm-3, this could be due to a number of reasons. Human error plays quite a factor in these titrations carried out. The stirrers were used more then once during each titration, and working in such small amounts, some of the solution could have been transferred between wells. The micropipette should also be either attached to the clamp or held vertically to the well plate, else the meniscus will not be read correctly or at eye level to be accurate. The titrations should be watched very carefully too, as the colour change could be instant. In the above case there were subtle changes, therefore the judgement of the colour could differ from person to person.
Safety Precautions taken
All staff and students must be aware of the fire exits are and be able to locate the first aid kits. Safety goggles and lab coats must be worn in the laboratory at all times as part of a safety precaution. There must be no food or drink consumed, no smoking, running or throwing of equipment in the laboratory as it is prohibited. In case of any spillages or accidents with chemicals, wash with excess water.
