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The aim of this investigative experiment is to discover the accurate concentration of sulphuric acid (H2SO4), which is found in a solution. The concentration is thought to be between 0.05 mol dm�3 and 0.15 mol dm�3. I have been given access to anhydrous potassium carbonate (K2CO3) and a range of indicators.

Plan

In order to obtain the concentration of the acid in the solution I have to titre the known solution of potassium carbonate with the unknown sulphuric acid. The indicator I will be using to indicate when the reaction is fully completed is methyl orange. This is because I am using a strong acid and a weak alkali and methyl orange is the most appropriate indicator for this type of acid-alkali titration.

To begin with I will have to prepare a standard solution of potassium carbonate that will be used in the titration with sulphuric acid. The potassium carbonate is of known strength and volume in contrast with the unknown concentration of sulphuric acid. This is the equation for the titration:

H2SO4 (aq) + K2CO3 (aq) –> K2SO4 (aq) + H20 (l) + CO2 (g)

Preparing a standard solution:

These are the requirements, showing apparatus and the quantities of the materials to be used.

Apparatus:

� Safety glasses

� Spatula:

� Weighing bottle

� Balance weighing to within 0.01g

� Wash bottle of distilled water

� Beaker, 2503cm

� Conical flask, 250cm3

� Pipette Filler

� Pipette, 25cm3

� Tripod stand

� 2 clamps

� White tile

� Burette, 50ml

� Filter funnel

Solutions:

� Sulphuric acid, 50cm3

� Methyl orange.

� Anhydrous potassium carbonate, 2.55g

Fair test and accuracy

� When reading, eye must be level with meniscus to read the burette.

� We will also wash all equipment when doing the other solutions. This is due to the different solutions mixing, so they give a different compound. We wash these in water because when you mix an acid or alkali with water, the pH level will not change since water is neutral. If we did not use water or anything neutral, then we will be changing the solution and this factor can give us inaccurate results.

Procedure – preparing my standard solution

1: Using a spatula put between 2.4g to 2.6g of potassium carbonate on the weighing bottle. Weigh the potassium carbonate on the balance, and make sure that you have an accurate weight between 2.4g to 2.6g of potassium carbonate, weighing to the nearest 0.01g.

2: Put about 250cm3 of distilled water into a 250cm3 beaker. Carefully move the bulk of the potassium carbonate from the weighing bottle into the beaker.

3: Shake the beaker to break up the solid.

4: Show your beaker to the teacher.

Procedure – titration

1: Set-up the tripod stand and the burette, held firmly in position by the clamps, you can refer to the diagram.

2: Fill the burette using a filter funnel with 50ml of sulphuric acid. Air bubbles should be avoided. Read off the zero mark at eye level to ensure that the bottom of the meniscus is on the mark.

3: Use a pipette filler and pipette to transfer 25cm3 of potassium carbonate in your beaker to a 250cm3 conical flask. Air bubbles must be avoided.

4: Add 3 drops of the methyl orange indicator to the conical flask (methyl orange is being used as the titration is between a strong acid and a weak alkali).

5: Now perform a rough titration by running the sulphuric acid, whilst swirling the flask of alkali, until the solution turns a pink colour. This is the end point where the acid and the alkali have neutralised each other. Note and record the rough value of the volume of acid in a table.

6: Now perform a second titration, this is as before except that this time run in the acid until about 2ml less than the rough value has been added, then proceed to run in the acid a drop at a time, keeping a close eye on the colour of the solution. You try to get an orange colour in the solution.

Record this end point in the table as before. Remember to use the pipette as before with the exact same volume of the standard solution, and the exact same volume of sulphuric acid.

7: Continue to do the titration until you get two readings within 0.1cm of each other. Remember to record the volume of sulphuric acid required each time. Now work out the average of the accurate readings and record this in the same table.

Once the titration has been completed and a sufficient average value has been calculated, you can go on to evaluate and analyse the experiment. Here is a risk assessment of the chemicals used in my experiment.

Risk Assessment

Hazards

� Do not shake the conical flask too vigorously or the solution will spill.

� Always be prepared for a spill.

� Make sure that the burette is closed when not in use.

� Do not fill the burette right to the top.

� Make sure the clamp stand is held tightly by your partner.

� Make sure the burette is held tightly by the clamp stand.

� Make sure the solution are put in a safe place and is not put on the edge of the table.

� Make sure every solution, liquid, etc, solutions are put in a safe place and is not put on the edge of the table.

� MAKE SURE YOU HAVE YOUR SAFETY GLASSES ON DURING THE WHOLE EXPERIMENT

Potassium carbonate

Although the solution I am using is very dilute I should still be aware that contact with eyes, skin and clothing must be avoided. As such I am to be wearing goggles and a lab coat. Also any slippages would cause the area to become slippery and dangerous, if they are not cleaned up quickly.

Methyl orange

The same precautions apply for this as did with the potassium carbonate.

Sulphuric Acid

It is very corrosive. If it is swallowed, wash out mouth and give one or two glasses of water. Don’t induce vomiting. Seek medical advice as soon as possible.

If it is splashed into the eye flood the eye with gently running tap water for 10 minutes. Seek medical advice.

If it is spilt on clothes or skin remove the clothes quickly and wipe as much liquid as possible away with a dry clothe before drenching the area with lots of water. If a large area of skin was affected or blistering seek medical advice.

If it is spilt in the laboratory, wear eye protection and gloves and cover with mineral absorbent and scoop into a bucket. Add anhydrous potassium carbonate and leave to react. Once the reaction has occurred add lots of cold water.

Results – titration

My plan for the titration should provide precise and reliable results if all chemicals are weighed and measured to the correct mass or volume. The apparatus should be set-up properly with no judgmental errors such as the reading of the volume on the dropping pipette or burette. This is the table by which I will be recording the volumes of sulphuric acid used in the titration, hence my results; I will calculate the average volume from these results. Using this form of table should allow me to provide a clear and concise way of representing my results and should aid me in my analysing and evaluating section of this investigation.

Rough/cm3

1/cm3

2/cm3

3/cm3

4/cm3

Final

14.8

29.85

44.88

15.04

30.01

Initial

0

14.8

29.85

0

15.04

Titres

14.8

15.05

15.03

15.04

15.06

Average = 15.05cm3 (2.dp)

Calculation

Unknown = Sulphuric Acid

Known = Potassium Carbonate

To calculate the concentration of the potassium carbonate solution used:

2.55g solid anhydrous potassium carbonate

Mass

RMM

RMM of K2CO3 = (2 x 39.1) + 12.0 + (3 x 16.0)

= 138.20g

Moles K2CO3 = 2.55/138.2 = 0.0185m

= 1.85 x 10-2m

Mole x 1000

Volume volume = 250cm3

1.85 x 10-2 x 1000

250

Therefore, initial molarity of potassium Carbonate = 7.4 x 10-2 Mol/dm3 (0.074 Mol/dm3)

Concentration of acid solution:

Equation of neutralisation of sulphuric acid with potassium carbonate:

H2SO4 (aq) + K2CO3 (aq) –> K2SO4 (aq) + H20 (l) + CO2 (g)

Which means that the mole ratio of sulphuric acid: potassium carbonate is 1:1

Amount of potassium carbonate needed to neutralise 25cm3 sulphuric acid (average titre) = 15.00cm3

Moles K2CO3 in 15.00cm� of 0.074 mol/dm� solution

= 0.01505 x 0.074 = 0.00113 moles K2CO3

Ratio K2CO3:H2SO4 = 1:1 therefore moles H2SO4 = 0.00111

0.00113 moles H2SO4 in 25cm�

Concentration of H2SO4

= 0.00113 x 1000

25

Therefore initial molarity of sulphuric acid

= 0.0452 mol/dm� (4.52 x 10-2 Mol/dm3)

Evaluation

The results and calculations show what I believe to be a fairly inaccurate concentration of sulphuric acid that is 0.0452 mol/dm�. I carried out the experiment as stated in my method until I had concordant titres of within 0.1cm� of each other. I did not include my rough titre in the average as it was performed to a much lower degree of accuracy and was merely to familiarise myself with the equipment and experiment.

I am almost certain that several errors have occurred in the experiment, mainly human error but also some caused by the procedure, technique and equipment. My reading of the meniscus in the burette could be inaccurate and therefore create an error. Even though I rinsed all my apparatus with distilled water before use (and the burette with sulphuric acid) it is still likely that some contamination occurred which would have affected my results. The amount of sulphuric acid I used in each titration was determined by when I saw a colour change in the indicator in the conical flask. It is very possible that I could have missed the exact moment when the colour change and therefore neutralisation occurred. This would mean too much or too little acid may have been added to the flask giving an inaccurate representation. The percentage errors associated with the experiment can be calculated by:

(Precision error x 100) / Actual reading

Errors caused by glassware and equipment

� Volumetric flask when filled correctly has a precision error of 0.2cm� and therefore has a percentage error of 0.08%

� All burette readings include 2 decimal places in which the second figure was a 0 or a 5 as I could only determine between these measurements. This gives an error of 0.2% for each reading.

� The digital balance gave readings of 2 decimal places, which means the actual reading could be +/- 0.005g of the recorded reading; (0.005 x 100 /1.06 = 0.47) so the balance delivers a percentage error of 0.47%.

Errors caused by technique

� Mixing of the solution in the conical flask may not be totally through.

� The burette and pipette may not have been thoroughly washed out with the solutions used.

� The conical flask may not have been thoroughly washed out with distilled water between titres.

� The end point may not be accurate if the solution from the burette is not added drop by drop with continuous swirling.

� Too much or too little indicator may have been added each time

It is not possible to place a value on the effect of human error on the reliability and accuracy of results. However, further repetition of the experiment would limit the effect human error has on results.

Improvements to the investigation would be mainly aiming to reduce the human error. This could be done by using equipment that displays values and measurements digitally, or detect the colour change more accurately.

Overall I do not believe my results could have been that inaccurate seeing as my titres were the same. I feel that the procedure allowed me to discover the accurate concentration of the acid to a fairly accurate and reliable degree.

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