Chemguide: Support for CIE A level Chemistry


Learning outcomes 7.2(e) and 7.2(f)

These statements are about titration curves and the use of indicators. I am covering them together because the two things are so closely related.

Before you go on, you should find and read the statements in your copy of the syllabus.


Titration curves

It makes sense to look at titration curves before you look at indicators. Start by reading the page about pH (titration) curves.

This is quite a complicated page, and it is easy to get confused between the shapes of the various curves. In the first instance, stop when you get to the heading "More complicated titration curves".

What I would suggest you do is ignore these more complicated graphs until after you have read about indicators later on. I will remind you further down this page.


Note:  In the exams from June 2007 up to November 2013 (based on the same syllabus statement), a question involving complex graphs like the ones in this last section has only been asked once, for a total of 2 marks. You could get one of those marks by using some simple information given in the question. The examiners reported that only a very small minority of candidates got the second mark.


Don't worry about the references to buffer solutions throughout the page. Once you know about buffer solutions later, you could come back and look at this again, but it is more important to be able to draw the curves.

Concentrate on how the curves relate to each other. Make sure that you can draw the four basic curves shown in the summary section (just before you get into the complicated shapes).

And make sure that you can understand why they are different - just in terms of the likely pHs involved.

So:

  • If you start with a strong base, the pH will probably be around 13 - 14.

  • If you start with a weak base, the pH will be lower - say 11 or 12.

  • If you use a strong acid, the pH will probably end up around 1 - 2.

  • If you use a weak acid, the pH will probably end up around 4.

If you understand this, you should be able to draw a reasonable approximation to these curves.

Now do the same thing for the curves where you add an alkali to an acid. Look carefully at the graphs on the pH curves page and work out what happens if you start with a strong or weak acid, and add a strong or weak base. You need to know what your likely start and finish points are going to be.


Indicators

Read the page about acid-base indicators.

Ignore the final bit about the titration of sodium carbonate solution for now.

You don't need to know about the structure of methyl orange (or any other indicator), but you do need to know:

  • the colour changes for methyl orange and phenolphthalein;

  • that the indicator has to change colour on the steep bit of the titration curve;

  • that either methyl orange or phenolphthalein can be used to titrate a strong acid against a strong base;

  • that to titrate a strong base with a weak acid you can use phenolphthalein;

  • that to titrate a weak base with a strong acid you can use methyl orange;

  • that no indicator will give you an accurate end point for a weak base / weak acid titration.

And a final comment on this last point:

Both the teacher support material and the CIE Chemistry Coursebook mention the use of the indicator bromothymol blue in this context, to demonstrate the fact that it won't give a good result.

Bromothymol blue has a pH range of 6.0 to 7.6, and so bridges the end point of a typical weak acid / weak base titration.

However, the colour change isn't sharp. It will change gradually from blue through green to yellow while you add perhaps 1 cm3 of weak acid to a weak base. You can't get an accurate titration out of this.


The more complicated titration curves

Once you are reasonably happy about this, go back to the pH (titration) curves page, and read the end section on "More complicated titration curves".

Then look at how a careful choice of indicators lets you pick up both end points in the titration of sodium carbonate solution with dilute hydrochloric acid on the page about acid-base indicators.

Don't spend too much time on this, but you should be aware that it is possible to get these more complex titration curves under some circumstances.


Finding pKa for a weak acid from a titration curve

I have no idea whether you are likely to be asked about this or not. This isn't mentioned by the syllabus, and nothing has been asked so far. It shouldn't be asked, but that doesn't always seem to stop CIE asking things anyway!

It is mentioned in the teacher support material, which says:

Measuring the pH at regular intervals during the titration of a weak acid with a strong base can be used as a method for calculating the pKa of the acid, using the equation
pKa = pH + log10 ([HA]/[A- ]).

There is nothing to suggest that you should learn this equation, but I will just quickly show you how to use it to find a value for pKa using a titration curve just in case you are given it, and asked to use it.

Suppose you have a weak acid in a flask, and run in a strong base like sodium hydroxide solution from a burette, plotting the pH as you add it.

You can tell when the acid has been neutralised when the titration curve suddenly goes steep.

You read off the volume of sodium hydroxide needed for neutralisation, and work out what half that value would be. So, for example, if the volume needed to neutralise the acid turned out to be 20.4 cm3, then the volume you would calculate would be 10.2 cm3.

Then you use the graph to find the pH when that half-volume of sodium hydroxide solution had been added. That pH is also the pKa of the acid.

Why does this work?

The ionic equation for the reaction between hydroxide ions and a weak acid HA is:

Think about what happens when the acid has been half neutralised. Half of it has been turned into A- ions; the other half remains as HA.

That means that at half-neutralisation:

[HA] = [A- ]

Now relate that to the formula that CIE quote:

Because the top and bottom of the log term on the right are equal, that term works out to be log 1. And log 1 = 0. So that term disappears and

pKa = pH

The pKa is the pH of the solution at half-neutralisation.


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© Jim Clark 2011 (last modified April 2014)