Volumetric Analysis (1) Carrying out an acid-base titration.

You are probably wondering what volumetric analysis is on about. 

Volumetric analysis is the use of the titration technique to determine the concentration of a solution whose concentration is not known. 

In a titration, a measured volume of a solution of known concentration is added (from a burette) to a known volume of another solution (measured from a pipette) whose concentration we don’t know until the reaction is complete. 

The point at which the reaction is complete, the equivalence point or the stoichiometric point, can be shown visibly by a colour change of an indicator dye in the measured solution. 

This point is known as the end point of the titration. 

Because volumes of solution are being measured this technique is termed volumetric analysis.

This is from Chemwiki:

“Titration is the slow addition of one solution of a known concentration (called a titrant) to a known volume of another solution of unknown concentration until the reaction reaches neutralization, which is often indicated by a color change.”

You will probably have already carried out a titration.

The kit looks like this:

I’ve see several diagrams on the net where I would take exception to the way the set up has been arranged

First off, here is a set up where the indicator has been placed in the solution in the burette!!

No, because that alters the concentration of the solution in the burette since acid-base indicators are usually weak acids. 

and here:

In subsequent posts I’m going to discuss three types of titration

1)  An acid-base titration

2)  A redox titration

3)  A precipitation titration

Let’s begin by looking at an Acid Base titration

There three types of Acid Base titrations

a)  strong acids with strong bases

b)  strong acids with weak bases

c)   weak acids with strong bases

The titration technique is the same for each but the indicator may need to change to reveal the specific end point as each end point occurs over a different pH range.

Let’s look at the technique first then at the particular differences between these three acid base titrations

The Titration Technique

The example I’m going to discuss is that between a strong base, sodium hydroxide solution and the weak acid, sodium hydrogen phthalate C₈H₅O₄K.

This titration is often carried out because sodium hydroxide solution on standing in the lab changes concentration.

It reacts with carbon dioxide in the air to form sodium carbonate–you can sometimes see a white solid around the neck of a bottle of sodium hydroxide or even small slivers of the solid in the alkaline solution. 

2NaOH(aq)   +   CO₂(g)       =       Na₂CO₃(s)  +   H₂O(l)

So to check the sodium hydroxide solution we are going to use is the concentration it says it is on the bottle we need to carry out a titration.

The acid used (sodium hydrogen phthalate C₈H₅O₄K) is a stable solid that does not decompose in solution or on standing in the lab. 

It is known as a primary standard because its solution is used to standardise the solution of sodium hydroxide. 

Sodium hydrogen phthalate C₈H₅O₄K is an acid in aqueous solution because it has one hydrogen atom that dissociates. 

It is monoprotic or monobasic producing one mole of hydrogen ions per mole of compound. 

To simplify the structure I’m going to call sodium hydrogen phthalate C₈H₅O₄K simply HA. 

Dr Carr’s Rescue Box gives you more details of this interesting compound. 

So the equation for the reaction is the simple one of acid plus base giving salt plus water. 

HA(aq)    +   NaOH(aq)    =     NaA(aq)    +    H₂O(l)

Or as an ionic equation      H⁺      +     OH^–    =     H₂O

To reveal the end point of the titration we use an acid/base indicator in this example it is phenolphthalein. 

Phenolphthalein is colourless in acidic solution and magenta in alkaline solution. 

The point at which one drop of the alkaline solution (or even less) turns the indicator the faintest pink or magenta is the end point of the titration. 

The acid base reaction is now complete because an exact number of moles of H+ ions have just neutralised the exact same number of OH– ions. 

Here are some tips on carrying out the practical itself.

1.   Rinse a burette with the sodium hydroxide solution.  Then fill it with the same solution and don’t forget to fill the jet below the tap this is a very common mistake I have found with students.  It leads to over-estimates of the volume run out of the burette to neutralise the solution in the flask. 

2.   Use a pipette filler to first rinse out the pipette then fill it with the acidic solution in this example.  Transfer the specified volume (usually 25ml) to the 250ml conical (Erhlenmeyer) flask. 

3.   Add 2 or 3 drops of phenolphthalein indicator to the solution in the flask. 

4.   Run sodium hydroxide solution from the burette into the conical flask swirling the flask as you do so.  As this titration is a rangefinder (or trial run) you just need an approximate volume of sodium hydroxide needed to neutralise the acid in the flask. So stop running the sodium hydroxide solution into the flask when the indicator has turned pink.

5.   Repeat the titration but this time as you near the approximate end point add the sodium hydroxide very slowly, drop by drop until the solution in the flask remains just faintly pink after swirling. 

6.   Read the burette to the nearest 0.05ml.  You may find a white card placed behind the burette helps you to see the gradations and a line drawn down the middle of the card helps you to pin point the actual value to the nearest 0.05ml. Some burettes are produced with this blue line and white background on the glass of the burette tube.

7.   Here is what I mean:

8.   The YouTube video helps at this point.  You can see in the picture that the line behind the burette really helps you to read its value precisely. 

9.   Notice too that in this example we are reading what’s left in the burette but in the titration we need to read what’s run out of the burette.  So we must read down and therefore we can see that 42.3ml have run out of the burette.  Reading the burette is another place where many students in my experience have made mistakes.   

10.                To finish the experiment you need to repeat the titration at least three times or until you have concordant readings (that is readings to within 0.1ml of each other) for the volume of sodium hydroxide needed to neutralise the acid in the flask. 

11.                Achieving concordant readings is another challenge because it means you have to wash and rinse out your apparatus carefully and thoroughly between titrations. 

12.                The final tip is to wash out the burette very thoroughly with distilled water after the experiment is over since sodium hydroxide is able to dissolve glass and alter the bore and capacity of the burette if left in the instrument for a length of time. 

13.                At the start of your experiment you might have a results table that might look like this:

Pipette solution	Sodium hydrogen phthalate			0.1 mol/dm3		25ml
Burette solution	Sodium hydroxide			approx. 0.1mol/dm3
Indicator	Phenolphthalein
Burette readings	Rangefinder	1	2	3	(4)
Final reading (ml)
First reading (ml)
Volume used (ml)
Mean titre (ml)

In the next post I’ll discuss titration curves and show you how to carry a calculation based on a set of titration results from an experiment like this one.