pH curves (titration curves)

And sodium ethanoate is produced beyond the

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Unformatted text preview: 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. Note: Bec aus e y ou almos t nev er do titrations with this c ombination, there is no real point in giv ing the graph where they are added the other way round. It is n't diffic ult to work out what it might look lik e if y ou are interes ted - tak e the beginning of the s odium hy drox ide added to ethanoic ac id c urv e, and the end of the ammonia added to hy droc hloric ac id one. The reas on that it is diffic ult to do thes e titrations is dis c us s ed on the page about indic ators . 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. Where would you like to go now? To the acid­base equilibria menu . . . To the Physical Chemistry menu . . . To Main Menu . . . © J im Clark 2002...
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