This is the reaction between Sulphuric Acid and Sodium Hydroxide:
Sulphuric acid and Hydrochloric acid each release different amounts of Hydrogen ions because they are different types of acids. Acids, which form one H+ ion from each acid molecule, are called Monoprotic. Acids, which form two, are called Diprotic. Acids, which form three, are called Triprotic (e.g. Orthophosphoric Acid [H3PO4])
I predict that for Monoprotic acids (e.g. Hydrochloric) :
Concentration of Acid * Volume of Acid = Concentration of Alkali * Volume of Alkali
The amount of acid needed to neutralise an alkali =
(MAl * VAl)/MA = (Molarity of Alkali * Volume of Alkali)/Molarity of Acid
I predict that for Diprotic acids (e.g. Sulphuric) :
Concentration of Acid * Volume of Acid = Concentration of Alkali * Double the Volume of Alkali
Molarity = how many molecules of the acid or alkali per 1000 cm3 (1 litre) of water.
Equipment
Goggles
(to protect the eyes)
Mat
(to protect the surface of the bench)
Burette
(to hold the acid)
Conical Flasks
(to hold the alkali)
Indicator
(to show the pH of the solution)
Funnel
(to pour the acid into the burette)
Measuring Cylinders
(to measure out the water, acid & alkali and to dilute the acids and the alkali)
Chemicals
H2O
(Water)
NaOH
(Sodium Hydroxide)
HCl
(Hydrochloric Acid)
H2SO4
(Sulphuric Acid)
Indicator
Method
This experiment is done using titration. This comes from the glossary of GCSE Chemistry Classbook, " Titration is a method of investigating the volumes of solution that react together."
In my experiment I chose to use two acids and one alkali. The two acids were Hydrochloric and Sulphuric, and were always kept at a concentration of 0.5 M. The alkali was Sodium Hydroxide and the concentration of this was varied. The alkali always had a volume of 25 ml, however dilute it was. I decided to do five experiments with each of the acids. The concentrations of the alkali were 1 M, 0.5 M, 0.4 M, 0.2 M & 0.1 M.
This was how I did each of the experiments:
1. Close the tap on the burette and fill the burette with the acid. Pour an amount of alkali in the conical flask and then dilute to the right concentration and volume. Add some indicator so the pH of the solution is known.
2. Set up the Burette in a clamp above a ceramic tile.
3. Place the conical flask with the alkali in under the end of the burette. With your left hand turn the tap gently and with your right hand gently swirl the liquid inside as the acid is dripping in.
4. Carry on letting the acid drip in and carry on swirling the flask until there is an indication that the solution in the conical flask is neutral. An example of which is that the indicator has changed to a different colour.
When the solution in the conical flask has become neutral, turn the tap off and stop the acid dripping in. You can then measure how much acid is left in the burette and subtract that amount from the starting volume to give the volume of acid that was needed to neutralise the alkali.
This process is repeated for each concentration of alkali with each acid.
This experiment was a fair test because :
- The strength of the acids was kept constant throughout the whole experiment.
- The same equipment was used.
- The same type of alkali was used for each experiment.
- Sensible volumes of acids and alkalis were used.
- The same volume of alkali was used each time.
- Only one thing was changed each time (the concentration of the alkali)
Key factors that could influence the results were :
- Accuracy of the dilution of the Alkali.
- Accuracy of the dilution of the Acids.
- How accurate the concentrations of the supplied Acids and Alkalis were.
- Which Indicator was used (Universal changes colour too gradually, so the pH of the solution wasn't as accurate as it could have been with other indicators.)
Great care had to be taken when using acids and alkalis. This was for safety reasons because both acids and alkalis are corrosive and will cause damage if they come into contact with either skin or eyes.
Results
It was necessary to try and repeat each result just to confirm that each result was correct and not anomalous.
Hydrochloric Acid
(Click on the graph to see an enlarged version of it.)
Sulphuric Acid
(Click on the graph to see an enlarged version of it.)
Conclusion
The results came out almost exactly as I had predicted them. None of them appeared to be anomalous. They were slightly different to the predictions because of some inaccuracies. The volume of water used to dilute the alkali could have been slightly too much or too little and the volume of alkali could have also been slightly inaccurate. Also, according to the results I got, either the molarity of the acid or the molarity of the alkali was slightly inaccurate. In addition, tap water was used, which is usually slightly acidic. This is because Carbon Dioxide dissolves in the water forming very weak carbonic acid which means that there will be H+ ions in the water which can start to neutralise the alkali. This means that some of the alkali will get neutralised by the water. It was also very hard to decide when the solution was neutral because the Universal Indicator changed very gradually from purple to blue and then went very quickly to red. Although, when we were very careful, we could get yellow and green.
This experiment has actually proved that the following is correct for all acids :
Even though the equations in my predictions were correct, the above equation is better because it is general to all acids and all alkalis.
On my graphs, since 0 concentration of Alkali will mean 0 volume of Acid, I decided that the best fit line should carry on down to the origin. Because there were no data there, I had to extrapolate the line downwards and because of this it is shown as a red dotted line. The data must be slightly inaccurate because the best fit line passes very close to the origin rather than actually passing through it.
Improvements
If this experiment was used again, I would try to improve it in the following ways :
- Instead of using Universal Indicator, I would use an indicator which only changed colour when the solution was Neutral rather than gradually changing as the pH became lower (e.g. Methyl Orange or Phenolphthalein).
- To get the pH more accurate next time, a pH meter could be used.
- To get it very accurate, the acids and alkalis supplied should be an exact amount (I am not sure whether the ones I used were exact).
- To use very recently distilled water, because the water will become acidic if left for a long time.
References
Books
Dunstan, S. 1968. Principles of Chemistry. Van Nostrand Reinhold
· Pages 187-88
· Pages 215- 253
McDuell, B. 1997. GCSE Chemistry Classbook. Letts
· Pages 171-72 · Pages 28-32 · Page 262
Computer programs
Microsoft Excel Version 97
Microsoft Word 97
Paint Shop Pro 4
