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 pH (TITRATION) CURVES

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تاريخ التسجيل : 12/11/2009
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مُساهمةموضوع: pH (TITRATION) CURVES   الخميس 10 ديسمبر 2009, 1:57 pm


pH (TITRATION) CURVES




This page describes how pH changes during various acid-base
titrations.



The equivalence point of a titration



Sorting out some confusing terms

When you carry out a simple acid-base titration, you use an indicator
to tell you when you have the acid and alkali mixed in exactly the
right proportions to "neutralise" each other. When the indicator
changes colour, this is often described as the end point
of the titration.

In an ideal world, the colour change would happen when you mix the
two solutions together in exactly equation proportions. That particular
mixture is known as the equivalence point.

For example, if you were titrating sodium hydroxide solution with
hydrochloric acid, both with a concentration of 1 mol dm-3,
25 cm3 of sodium hydroxide solution would need exactly the
same volume of the acid - because they react 1 : 1 according to the
equation.



In this particular instance, this would also be the neutral
point
of the titration, because sodium chloride solution has a
pH of 7.

But that isn't necessarily true of all the salts you might get
formed.

For example, if you titrate ammonia solution with hydrochloric acid,
you would get ammonium chloride formed. The ammonium ion is slightly
acidic, and so pure ammonium chloride has a slightly acidic pH.

That means that at the equivalence point (where you had mixed the
solutions in the correct proportions according to the equation), the
solution wouldn't actually be neutral. To use the term "neutral point"
in this context would be misleading.

Similarly, if you titrate sodium hydroxide solution with ethanoic
acid, at the equivalence point the pure sodium ethanoate formed has a
slightly alkaline pH because the ethanoate ion is slightly basic.

To summarise:



  • The term "neutral point" is best avoided.
  • The term "equivalence point" means that the solutions have been
    mixed in exactly the right proportions according to the equation.
  • The term "end point" is where the indicator changes colour. As
    you will see on the page about indicators, that isn't necessarily
    exactly the same as the equivalence point.





Simple pH curves

All the
following titration curves are based on both acid and alkali having a
concentration of 1 mol dm-3. In each case, you start with 25
cm3 of one of the solutions in the flask, and the other one
in a burette.


Although you
normally run the acid from a burette into the alkali in a flask, you may
need to know about the titration curve for adding it the other way
around as well. Alternative versions of the curves have been described
in most cases


Titration
curves for strong acid v strong base

We'll take
hydrochloric acid and sodium hydroxide as typical of a strong acid and a
strong base.





Running
acid into the alkali





You can see
that the pH only falls a very small amount until quite near the
equivalence point. Then there is a really steep plunge. If you
calculate the values, the pH falls all the way from 11.3 when you have
added 24.9 cm3 to 2.7 when you have added 25.1 cm3.







Running alkali into the acid



This is very similar to the previous curve except, of course, that
the pH starts off low and increases as you add more sodium hydroxide
solution.




Again, the pH doesn't change very much until you get close to the
equivalence point. Then it surges upwards very steeply.



Titration curves for strong acid v weak base
]

This time we are going to use hydrochloric acid as the strong acid
and ammonia solution as the weak base.





Running acid into the alkali





Because you have got a weak base, the beginning of the curve is
obviously going to be different. However, once you have got an excess
of acid, the curve is essentially the same as before.

At the very beginning of the curve, the pH starts by falling quite
quickly as the acid is added, but the curve very soon gets less steep.
This is because a buffer solution is being set up - composed of the
excess ammonia and the ammonium chloride being formed.



Notice that the equivalence point is now somewhat acidic ( a bit less
than pH 5), because pure ammonium chloride isn't neutral. However, the
equivalence point still falls on the steepest bit of the curve. That
will turn out to be important in choosing a suitable indicator for the
titration.




Running alkali into the acid


At the beginning of this titration, you have an excess of
hydrochloric acid. The shape of the curve will be the same as when you
had an excess of acid at the start of a titration running sodium
hydroxide solution into the acid.

It is only after the equivalence point that things become different.

A buffer solution is formed containing excess ammonia and ammonium
chloride. This resists any large increase in pH - not that you would
expect a very large increase anyway, because ammonia is only a weak
base.






Titration curves for weak acid v strong base


We'll take ethanoic acid and sodium hydroxide as typical of a weak
acid and a strong base.




Running acid into the alkali


For the first part of the graph, you have an excess of sodium
hydroxide. The curve will be exactly the same as when you add
hydrochloric acid to sodium hydroxide. Once the acid is in excess,
there will be a difference.




Past the equivalence point you have a buffer solution containing
sodium ethanoate and ethanoic acid. This resists any large fall in pH.



Running alkali into the acid





The start of the graph shows a relatively rapid rise in pH but this
slows down as a buffer solution containing ethanoic acid and sodium
ethanoate is produced. Beyond the equivalence point (when the sodium
hydroxide is in excess) the curve is just the same as that end of the
HCl - NaOH graph.




Titration curves for weak acid v weak base


The common example of this would be ethanoic acid and ammonia.



It so happens that these two are both about equally weak - in that
case, the equivalence point is approximately pH 7.



Running acid into the alkali


This is really just a combination of graphs you have already seen.
Up to the equivalence point it is similar to the ammonia - HCl case.
After the equivalence point it is like the end of the ethanoic acid -
NaOH curve.




Notice that there isn't any steep bit on this graph. Instead, there
is just what is known as a "point of inflexion". That lack of a steep
bit means that it is difficult to do a titration of a weak acid against a
weak base.




A summary of the important curves



The way you normally carry out a titration involves adding the acid
to the alkali. Here are reduced versions of the graphs described above
so that you can see them all together.








More complicated titration curves




Adding hydrochloric acid to sodium carbonate solution



The overall equation for the reaction between sodium carbonate
solution and dilute hydrochloric acid is:



If you had the two solutions of the same concentration, you would
have to use twice the volume of hydrochloric acid to reach the
equivalence point - because of the 1 : 2 ratio in the equation.

Suppose you start with 25 cm3 of sodium carbonate
solution, and that both solutions have the same concentration of 1 mol
dm-3. That means that you would expect the steep drop
in the titration curve to come after you had added 50 cm3 of
acid.

The actual graph looks like this:






The graph is more complicated than you might think - and curious
things happen during the titration.

You expect carbonates to produce carbon dioxide when you add acids to
them, but in the early stages of this titration, no carbon dioxide is
given off at all.

Then - as soon as you get past the half-way point in the titration -
lots of carbon dioxide is suddenly released.

The graph is showing two end points - one at a pH of 8.3 (little more
than a point of inflexion), and a second at about pH 3.7. The reaction
is obviously happening in two distinct parts.

In the first part, complete at A in the diagram, the sodium
carbonate is reacting with the acid to produce sodium hydrogencarbonate:



You can see that the reaction doesn't produce any carbon dioxide.

In the second part, the sodium hydrogencarbonate produced goes on to
react with more acid - giving off lots of CO2.



That reaction is finished at B on the graph.

It is possible to pick up both of these end points by careful choice
of indicator. That is explained on the separate page on indicators.



Adding sodium
hydroxide solution to dilute ethanedioic acid


Ethanedioic acid was previously known as oxalic acid. It is a diprotic
acid
, which means that it can give away 2 protons (hydrogen
ions) to a base. Something which can only give away one (like HCl) is
known as a monoprotic acid.




The reaction with sodium hydroxide takes place in two stages because
one of the hydrogens is easier to remove than the other. The two
successive reactions are:





If you run sodium hydroxide solution into ethanedioic acid solution,
the pH curve shows the end points for both of these reactions.




The curve is for the reaction between sodium hydroxide and
ethanedioic acid solutions of equal concentrations.







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pH (TITRATION) CURVES
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