Acid-base titrations
Acid-base titrations are based on the neutralization reaction. They are sometimes called alkalimetric titrations and general name of the method is alkalimetry, although these are not used as often as just "acid-base titration".
Acid-base titrations can be used to determine most acids and bases, strong and not too weak, monoprotic and polyprotic. For example we can use acid-base titration to determine concentration of hydrochloric acid, sulfuric acid, acetic acid, as well as bases - like sodium hydroxide, ammonia and so on. In some particular cases, when solution contains mixture of acids or bases of different strengths, it is even possible to determine in one titration composition of a mixture - for example sodium hydroxide and sodium hydrogen carbonate. Using acid-base back titration it is also possible to determine amount of substances that can be easily dissolved in acids, like calcium carbonate. To do so we would add known amount of hydrochloric acid to calcium carbonate and after the solid is dissolved we would titrate excess acid with a strong base.
Most commonly used reagents are hydrochloric acid and sodium hydroxide. Solutions of hydrochloric acid are stable, solutions of sodium hydroxide can dissolve glass and absorb carbon dioxide from the air, so they should be not stored for long periods of time.
There are many standard substances that can be used in acid base titrations. Those most popular are sodium carbonate Na2CO3, borax (disodium tetraborate decahydrate) Na2B4O7·10H2O and potassium hydrogen phthalate KHC8H4O4, often called simply KHP.
Type of indicator depends on several factors. One of them is the equivalence point pH. Depending on the titrated substance and titrant used this can vary, usually between 4 and 10. However, even if it is often possible (see list of pH indicators) we are rarely selecting indicator that changes color exactly at the equivalence point, as usually increase of accuracy doesn't justify additional costs. Thus in practice you will probably use phenolphtalein when NaOH is used as the titrant and methyl orange when titrating with the strong acid.
Acid-base titrations can be used to determine most acids and bases, strong and not too weak, monoprotic and polyprotic. For example we can use acid-base titration to determine concentration of hydrochloric acid, sulfuric acid, acetic acid, as well as bases - like sodium hydroxide, ammonia and so on. In some particular cases, when solution contains mixture of acids or bases of different strengths, it is even possible to determine in one titration composition of a mixture - for example sodium hydroxide and sodium hydrogen carbonate. Using acid-base back titration it is also possible to determine amount of substances that can be easily dissolved in acids, like calcium carbonate. To do so we would add known amount of hydrochloric acid to calcium carbonate and after the solid is dissolved we would titrate excess acid with a strong base.
Most commonly used reagents are hydrochloric acid and sodium hydroxide. Solutions of hydrochloric acid are stable, solutions of sodium hydroxide can dissolve glass and absorb carbon dioxide from the air, so they should be not stored for long periods of time.
There are many standard substances that can be used in acid base titrations. Those most popular are sodium carbonate Na2CO3, borax (disodium tetraborate decahydrate) Na2B4O7·10H2O and potassium hydrogen phthalate KHC8H4O4, often called simply KHP.
Type of indicator depends on several factors. One of them is the equivalence point pH. Depending on the titrated substance and titrant used this can vary, usually between 4 and 10. However, even if it is often possible (see list of pH indicators) we are rarely selecting indicator that changes color exactly at the equivalence point, as usually increase of accuracy doesn't justify additional costs. Thus in practice you will probably use phenolphtalein when NaOH is used as the titrant and methyl orange when titrating with the strong acid.
Remember, that what we calculate is not end point - but equivalence point.
In the equivalence point we have solution containing pure salt that is a product of the neutralization reaction occurring during titration. Thus calculation of equivalence point pH is identical with calculation of pH of salt solution.
Depending on the type of titration there are at least three different cases to discuss.
In the case of titration of strong acid with strong base (or strong base with strong acid) there is no hydrolysis and solution pH is neutral - 7.00 (at 25°C).
In the case of titration of weak acid with strong base, pH at the equivalence point is determined by the weak acid salt hydrolysis. That means we have to find pKb of conjugated base and calculate concentration of OH- starting from there, then use pH=14-pOH formula. See pH of weak acids and bases lecture and pH cheat sheet for details of calculation.
In the case of titration of weak base with strong acid, situation is very similar - pH at the equivalence point is determined by the weak base salt hydrolysis. Thus we need pKa of conjugated acid to calculate H+ and pH. Check lecture and cheat sheet mentioned above for details.
In the case of polyprotic acids and bases calculations get much harder. You may try to follow methods described in the lecture on polyprotic acids and bases pH calculation, or you may use BATE - pH calculator.
Calculate pH at the equivalence point of formic acid titration with NaOH, assuming both titrant and titrated acid concentrations are 0.1 M. pKa = 3.75.
At the equivalence point we have a solution of sodium formate. As both concentrations of titrated acid and titrant are identical, and monoprotic formic acid reacts 1:1 with sodium hydroxide, we have to add identical volume of base to the given volume of acid. That in turn means that final volume is twice that of initial volume of acid sample, so after dilution concentration of formate must be half that of acid - that is 0.05 M.
We have titrated weak acid, so to calculate pH we have to calculate concentration of OH- from formate hydrolysis first. Formate is a weak base with
1
Let's try to use the most simplified formula first:
2
Using 10-10.25 and 0.05 we get
3
To be sure we can use the simplified formula we have to check, whether hydrolysis was below 5%. To do so, we should divide concentration of OH- by initial concentration of formate. That means
4
Obviously assumption about low hydrolysis degree is correct, and we can proceed with calculation of pOH:
5
and pH:
6
What is pH at the equivalence point of 0.0211 M H2SO4 titrated with 0.01120 M NaOH?
7.0
OK, that was very short answer, now a little bit longer one. This is case of strong acid titrated with strong base, so we expect pH at equivalence point to be that of neutral solution - that is, 7.00. In reality the answer will be slightly different. Three reasons for that. First, sulfuric acid has pKa1 = -3 (very strong acid) but second dissociation step has pKa2 = 2.0, so it is much weaker. Still strong, but weak enough so that its hydrolysis can't be ignored, especially in more concentrated solutions. Second, NaOH - while strong base - is much weaker than it is commonly assumed, with pKb = 0.2 (see pKb of NaOH in ChemBuddy FAQ for details), so in precise calculations its hydrolysis can't be neglected as well. Finally, there is a reason that we are ignoring in all our examples, but that can't be neglected in the real lab - that is, activity coefficients of all ions involved are not 1 (more on that in ChemBuddy lecture on ionic strength and activity coefficients). If we take all these things into account we can calculate pH of the solution to be 7.05, close enough to 7.0.
In the equivalence point we have solution containing pure salt that is a product of the neutralization reaction occurring during titration. Thus calculation of equivalence point pH is identical with calculation of pH of salt solution.
Depending on the type of titration there are at least three different cases to discuss.
In the case of titration of strong acid with strong base (or strong base with strong acid) there is no hydrolysis and solution pH is neutral - 7.00 (at 25°C).
In the case of titration of weak acid with strong base, pH at the equivalence point is determined by the weak acid salt hydrolysis. That means we have to find pKb of conjugated base and calculate concentration of OH- starting from there, then use pH=14-pOH formula. See pH of weak acids and bases lecture and pH cheat sheet for details of calculation.
In the case of titration of weak base with strong acid, situation is very similar - pH at the equivalence point is determined by the weak base salt hydrolysis. Thus we need pKa of conjugated acid to calculate H+ and pH. Check lecture and cheat sheet mentioned above for details.
In the case of polyprotic acids and bases calculations get much harder. You may try to follow methods described in the lecture on polyprotic acids and bases pH calculation, or you may use BATE - pH calculator.
Calculate pH at the equivalence point of formic acid titration with NaOH, assuming both titrant and titrated acid concentrations are 0.1 M. pKa = 3.75.
At the equivalence point we have a solution of sodium formate. As both concentrations of titrated acid and titrant are identical, and monoprotic formic acid reacts 1:1 with sodium hydroxide, we have to add identical volume of base to the given volume of acid. That in turn means that final volume is twice that of initial volume of acid sample, so after dilution concentration of formate must be half that of acid - that is 0.05 M.
We have titrated weak acid, so to calculate pH we have to calculate concentration of OH- from formate hydrolysis first. Formate is a weak base with
1
Let's try to use the most simplified formula first:
2
Using 10-10.25 and 0.05 we get
3
To be sure we can use the simplified formula we have to check, whether hydrolysis was below 5%. To do so, we should divide concentration of OH- by initial concentration of formate. That means
4
Obviously assumption about low hydrolysis degree is correct, and we can proceed with calculation of pOH:
5
and pH:
6
What is pH at the equivalence point of 0.0211 M H2SO4 titrated with 0.01120 M NaOH?
7.0
OK, that was very short answer, now a little bit longer one. This is case of strong acid titrated with strong base, so we expect pH at equivalence point to be that of neutral solution - that is, 7.00. In reality the answer will be slightly different. Three reasons for that. First, sulfuric acid has pKa1 = -3 (very strong acid) but second dissociation step has pKa2 = 2.0, so it is much weaker. Still strong, but weak enough so that its hydrolysis can't be ignored, especially in more concentrated solutions. Second, NaOH - while strong base - is much weaker than it is commonly assumed, with pKb = 0.2 (see pKb of NaOH in ChemBuddy FAQ for details), so in precise calculations its hydrolysis can't be neglected as well. Finally, there is a reason that we are ignoring in all our examples, but that can't be neglected in the real lab - that is, activity coefficients of all ions involved are not 1 (more on that in ChemBuddy lecture on ionic strength and activity coefficients). If we take all these things into account we can calculate pH of the solution to be 7.05, close enough to 7.0.
No comments:
Post a Comment