Amphoteric substances are chemical substances able to act as both acids and bases. This is IUPAC's definitions:
Amphiprotes, which is molecules/ions able to act as both Brønsted acids and bases, e.g. water and HCO3−.
Ampholytes are molecules that can act as both acid and base, as it contains both groups, e.g. amino acids.
Whan reading some educational materials, the two types amphoteric substances are treated as if they are the same, and all the substances are just referred to as ampholytes. This can lead to quite a few misunderstandings.
There is a distinct difference between amphoteric substances and buffers, which is easily overlooked, when starting to work with acid/base chemistry, amphoteric substances are one molecule or one ion having both functions, whereas buffers are mixtures of acid and base, i.e. two different molecules/ions. Likewise you have to be careful not to get ampholytes mistaken for zwitter ions, which is ions having both positive and negative charges at the same time. There is an overlap, as ampholytes are also zwitter ions, but not all zwitter ions are ampholytes (a lot of people like this analogy: all elephants are grey, but grey things are not all elephants).
The most important of the amphiprotic substances is water, as all acid/base chemistry depends on water, and water's autoprotolysis becomes a dominating effect at highly diluted solutions. For educational chemistry problems, autoprotolysis is rarely relevant or interesting, as the focus is on the stoichiometric calculations, and you work with concentrations where the contribution from water is negligible, but for industrial processes, e.g. for discharge of wastewater in regards to safety and environmental legislation, the autoprotolysis of water is essential in regards to pH requirements.
The autoprotolysis of water, and thus why water is amphiprotic, looks like this:
2 H2O(l) H3O+(aq) + OH−(aq)
An amphiprote that isn't water could be KH2PO4. It is equivalent to phosphoric acid titrated to the equivalence point using potassium hydroxide. The ion H2PO4− is now able to work as an acid, by splitting off an additional H+, e.g. by reacting with OH−:
H2PO4−(aq) + OH−(aq) HPO42−(aq) + H2O(l)
The same ion can also work as base and accept an H+, like this:
H2PO4−(aq) + H+(aq) H3PO4(aq)
Ampholytes having both acidic and basic functional groups are common, e.g. amino acids. Looking at the amino acid glycine, which is 3-amino-propanoic acid, it has an acid group in one end, and an amine, which is an basic group, in the other end. Therefore, when adding hydroxide ions, there will be a reaction with the carboxylic acid:
Beware, ampholytes can have multiple acidic or basic functional groups, e.g. aspartic acid, which is an amino acid having two carboxylic acid groups and one amino group, or asparagine, which is an amino acid having one carboxylic acid, one amine, and one amide.
Formal writing for amphoteric compounds
The way to write the reactions for amphoteric substances in general, is these three:
In the Danish educational material that I know of, this is called the reaction of ampholytes, but the equations are valid for amphoteric compounds in general, and the examples being used are often amphiprotes instead.
The notation may appear a bit misguided, as the ampholyte here must have the charge −1 or lower, and the base must have the charge −2 or lower. Also the writing contradict what we normally consider correct balancing of reaction equation, as the charge is out of balance. This is done to make the reaction equations generic, so they can be used for all ampholytes, regardless of their initial charges. The notation may also leave the impression that there is only one of each of the acidic and basic groups. This is not the case either. You just look at it as a collective acid or base. Likewise it may also appear backwards, that at the reaction where the ampholyte reacts as a base (2nd reaction equation), it is written as the dissociation of an acid. This is done due to the math, when calculating equilibria and thus pH, for the ampholytes.
pH calculations for amphoteric substances
pH calculation on amphoteric compounds quickly becomes complicated. Basically you have to treat the calculations multiple equilibria affecting each other.
By multiplying the two acid dissociation constants, you get:
Ka(acid) · Ka(Ampholyte) =
[H3O+] · [Base]
[H3O+] · [Amp]
[H3O+] 2 · [Base]
We now have an equation with three unknowns. To move along from here, you use a simplification called the ampholyte equation.
The ampholyte equation, which is also used for amphiprotes, and very popular for educational purposes, works like this:
At very dilute solutions, where both acid and base are weak and water's autoprotolysis is still negligible, [Acid] ≈ [Base], and can you make the approximation that
Ka(acid) · Ka(Ampholyte) = [H3O+] 2
pKa(acid) + pKa(Ampholyte)
Note that according to the equation, pH is independent of concentration. This is obviously not correct, an thus the reason why the equation has very limited use, e.g. for high school chemistry.
The way the equation is constructed, it means that if you have an amfiprote like HCO3−, then the two pKa values are pKa1 and pKa2 for H2CO3, which makes the equation very easy to work with. If you have an ampholyte, you use pKa for the acid and pKa for the base. If you have multiple acid or base groups, e.g. like aspartic acid, you use the pKa for the dominating, i.e. the strongest, acid. At differences in pKa of more than 1, for the two acids or bases, the contribution from the weakest acids or bases will be negligible compared to the general inaccuracy inherent in the equation.