Table of Contents
Understanding Amphoterism
Amphoterism is the ability of a substance to act both as an acid and as a base, depending on what it reacts with. In the Brønsted sense, an amphoteric substance can both donate and accept protons $($H$^+)$; in the Lewis sense, it can both accept and donate electron pairs in different reactions.
Amphoterism is closely related to, but not identical with, the notion of an amphiprotic substance (one that can donate and accept protons). All amphiprotic substances are amphoteric in the Brønsted sense, but not all amphoteric substances are necessarily amphiprotic (for example, some metal oxides show amphoterism without proton transfer being the main focus).
Amphiprotic Substances and Conjugate Pairs
Many amphoteric species are conjugate acid–base pairs of polyprotic acids. For such species, amphoterism arises because they sit “in the middle” between a stronger acid form and a stronger base form.
A typical pattern is:
- A diprotic acid: $H_2A$
- Its first deprotonation:
$$H_2A \rightleftharpoons H^+ + HA^-$$ - Its second deprotonation:
$$HA^- \rightleftharpoons H^+ + A^{2-}$$
Here the intermediate species $HA^-$ is amphiprotic:
- It can donate a proton (act as an acid): $HA^- \to A^{2-} + H^+$
- It can accept a proton (act as a base): $HA^- + H^+ \to H_2A$
Examples:
- Hydrogencarbonate ion $HCO_3^-$
- Dihydrogen phosphate ion $H_2PO_4^-$
- Hydrogen sulfate ion $HSO_4^-$
These species will behave as acids in the presence of stronger bases and as bases in the presence of stronger acids.
Amphoterism of Water
Water is the simplest and most important amphoteric substance in chemistry.
In the Brønsted sense:
- As an acid:
$$H_2O + NH_3 \rightarrow NH_4^+ + OH^-$$
Water donates a proton to ammonia. - As a base:
$$H_2O + HCl \rightarrow H_3O^+ + Cl^-$$
Water accepts a proton from hydrochloric acid.
Water is also amphiprotic: it can both donate and accept $H^+$. In pure water, this leads to the self-ionization (autoprotolysis) of water:
$$2H_2O \rightleftharpoons H_3O^+ + OH^-$$
This amphoterism of water is the basis for the pH scale and for acid–base behavior in aqueous solutions.
Amphoteric Hydroxides and Oxides of Metals
A particularly important group of amphoteric substances are certain metal hydroxides and oxides, especially of some p-block and d-block metals. These compounds can react both with acids (as bases) and with strong bases (as acids, typically via formation of complex anions).
General Pattern
For a metal hydroxide $M(OH)_n$:
- With acid (acting as a base), it typically dissolves by forming metal salts:
$$M(OH)_n + nH^+ \rightarrow M^{n+} + nH_2O$$ - With strong base (acting as an acid), it can dissolve by forming complex anions:
$$M(OH)_n + OH^- \rightarrow [M(OH)_{n+1}]^- \quad \text{(simplified)}$$
Similarly for oxides $M_2O_n$:
- With acid: formation of salts and water.
- With base: formation of complex anions or “oxoanions.”
Which metals form amphoteric hydroxides/oxides depends on their position in the periodic table and oxidation state. Common examples are:
- Aluminum: $Al(OH)_3$, $Al_2O_3$
- Zinc: $Zn(OH)_2$, $ZnO$
- Tin(II): $Sn(OH)_2$, $SnO$
- Lead(II): $Pb(OH)_2$, $PbO$
- Chromium(III): $Cr(OH)_3$
These often correspond to metals that are not purely “metallic” or purely “nonmetallic” in character and often lie near the diagonal border between metals and nonmetals or in particular oxidation states.
Example: Amphoterism of Aluminum Hydroxide
Aluminum hydroxide, $Al(OH)_3$, is amphoteric.
- Reaction with acids (acting as a base):
With hydrochloric acid:
$$Al(OH)_3 + 3HCl \rightarrow AlCl_3 + 3H_2O$$
In aqueous solution, the hydrated ion is often written:
$$Al(OH)_3 + 3H^+ \rightarrow Al^{3+} + 3H_2O$$
- Reaction with strong bases (acting as an acid):
In excess sodium hydroxide solution, aluminum hydroxide dissolves, forming a complex aluminate ion:
$$Al(OH)_3 + OH^- \rightarrow [Al(OH)_4]^-$$
Often in practical terms:
$$Al(OH)_3 + NaOH \rightarrow Na[Al(OH)_4]$$
In acidic solution, $Al(OH)_3$ behaves as a base (neutralizing acid); in concentrated basic solution, it behaves as an acid (giving up hydroxide ligands to form a negatively charged complex).
Example: Amphoterism of Zinc Hydroxide
Zinc hydroxide, $Zn(OH)_2$, is another classic amphoteric hydroxide.
- With acids (acting as a base):
$$Zn(OH)_2 + 2HCl \rightarrow ZnCl_2 + 2H_2O$$
or
$$Zn(OH)_2 + 2H^+ \rightarrow Zn^{2+} + 2H_2O$$
- With strong base (acting as an acid):
$$Zn(OH)_2 + 2OH^- \rightarrow [Zn(OH)_4]^{2-}$$
In sodium hydroxide solution:
$$Zn(OH)_2 + 2NaOH \rightarrow Na_2[Zn(OH)_4]$$
This dual behavior is exploited in qualitative inorganic analysis when separating and identifying metal cations: amphoteric hydroxides dissolve both in strong acid and in excess strong base, which distinguishes them from purely basic hydroxides (which only dissolve readily in acid).
Amphoterism vs. Acid–Base Strength and pH
Whether an amphoteric substance behaves as an acid or base in a given situation depends on:
- The strengths of other acids and bases present.
- The pH of the solution.
- The stability of possible products (e.g., complex ions, salts).
For amphiprotic species, you can think in terms of possible acid–base equilibria:
- As acid:
$$HA^- \rightleftharpoons H^+ + A^{2-}$$ - As base:
$$H_2A \rightleftharpoons H^+ + HA^-$$
At low pH (high $[H^+]$), equilibria are shifted toward more protonated forms; the substance tends to behave as a base (accepting protons). At high pH (low $[H^+]$), equilibria shift toward more deprotonated forms; the same substance tends to behave as an acid (donating protons).
For amphoteric metal hydroxides, solubility often shows a characteristic “U-shaped” dependence on pH:
- Low pH: soluble as metal cations due to reaction with acid.
- Intermediate pH: low solubility as neutral hydroxide precipitates.
- High pH: soluble again as complex anions (e.g., $[Al(OH)_4]^-$, $[Zn(OH)_4]^{2-}$).
This behavior is important in precipitation and dissolution processes and is widely used in separation techniques and in controlling metal solubility in environmental and industrial contexts.
Examples of Amphiprotic Anions
A number of common ions in aqueous chemistry show amphiprotic behavior:
- Hydrogencarbonate ion:
- As an acid:
$$HCO_3^- \rightleftharpoons CO_3^{2-} + H^+$$ - As a base:
$$HCO_3^- + H^+ \rightleftharpoons H_2CO_3$$ - Dihydrogen phosphate ion:
- As an acid:
$$H_2PO_4^- \rightleftharpoons HPO_4^{2-} + H^+$$ - As a base:
$$H_2PO_4^- + H^+ \rightleftharpoons H_3PO_4$$ - Hydrogen sulfate ion:
- As an acid:
$$HSO_4^- \rightleftharpoons SO_4^{2-} + H^+$$ - As a base:
$$HSO_4^- + H^+ \rightleftharpoons H_2SO_4$$
In practice, which direction is relevant depends on the acids and bases present and the pH of the solution.
Significance and Applications of Amphoterism
Amphoterism is not just a theoretical curiosity; it has practical consequences:
- Buffer systems: Amphiprotic species (like $H_2PO_4^-/HPO_4^{2-}$ or $HCO_3^-$/related species) are integral components of many buffer solutions, including biological buffers in blood and cells.
- Qualitative inorganic analysis: Amphoteric hydroxides (Al, Zn, Cr, etc.) are separated and identified based on their ability to dissolve in both acids and bases.
- Materials chemistry and corrosion: The amphoteric behavior of aluminum oxide and hydroxide influences how aluminum reacts with acidic or basic environments, affecting corrosion resistance and surface treatments.
- Environmental chemistry: The solubility and mobility of metal ions in soils and waters depend on their amphoteric hydroxides and oxides, which dissolve or precipitate at different pH values.
Understanding amphoterism allows you to predict and control the behavior of substances that can switch roles between acid and base, depending on their chemical environment.