Reactions of Ions in Aqueous Solution (International A2)

Hello future chemists! This chapter dives into the fascinating world of transition metal ions and how they behave when dissolved in water. These reactions are fundamental to identifying these metals in the lab, so mastering this topic is essential for both theory and practical work. Don't worry if the formulas look long—we'll break down exactly what is happening step-by-step!

1. The Formation of Metal-Aqua Ions

When a transition metal salt dissolves in water, the positive metal ion ($\text{M}^{n+}$) attracts surrounding water molecules ($\text{H}_2\text{O}$). Water molecules act as ligands, forming dative covalent bonds (where the oxygen atom donates a lone pair of electrons) to the central metal ion.

Structure of Hexaaqua Ions

In aqueous solution, the metal ions covered in this syllabus typically form hexaaqua complexes, meaning six water molecules surround the central metal ion in an octahedral shape.

  • Divalent Ions (2+ charge): $\text{[M(H}_2\text{O)}_6\text{]}^{2+}$
    Examples: $\text{[Fe(H}_2\text{O)}_6\text{]}^{2+}$ (pale green), $\text{[Cu(H}_2\text{O)}_6\text{]}^{2+}$ (pale blue)
  • Trivalent Ions (3+ charge): $\text{[M(H}_2\text{O)}_6\text{]}^{3+}$
    Examples: $\text{[Fe(H}_2\text{O)}_6\text{]}^{3+}$ (yellow/violet), $\text{[Al(H}_2\text{O)}_6\text{]}^{3+}$ (colourless)
Quick Review: Key Terms
  • Ligand: A molecule or ion that donates a lone pair of electrons to a central metal ion to form a coordinate bond.
  • Complex Ion: A central metal atom or ion surrounded by ligands.
  • Co-ordination Number: The number of coordinate bonds attached to the central metal ion (here, usually 6).

2. Acidity of Metal-Aqua Ions

These complex ions behave as weak acids in water because they are capable of donating a proton ($\text{H}^+$). This process is called hydrolysis.

The Hydrolysis Reaction (Acting as an Acid)

The highly charged metal ion pulls electron density away from the oxygen atoms in the coordinated water ligands. This weakens the $\text{O-H}$ bonds, making it easier for an outside water molecule ($\text{H}_2\text{O}$) to act as a base and remove an $\text{H}^+$ proton from the ligand.

The general reaction for the first proton loss is: $$ \text{[M(H}_2\text{O)}_6\text{]}^{n+} + \text{H}_2\text{O} \rightleftharpoons \text{[M(H}_2\text{O)}_5\text{(OH)}\text{]}^{(n-1)+} + \text{H}_3\text{O}^+ $$

Why Trivalent Ions ($\text{3+}$) are More Acidic

The acidity of the hexaaqua ion depends heavily on the central metal ion's ability to weaken the $\text{O-H}$ bonds.

  • High Acidity means a strong positive pull on the electrons.
  • The strength of this pull depends on the Charge/Size Ratio of the metal ion.

Compare $\text{[Fe(H}_2\text{O)}_6\text{]}^{3+}$ (trivalent) with $\text{[Fe(H}_2\text{O)}_6\text{]}^{2+}$ (divalent):

The $\text{Fe}^{3+}$ ion has a higher charge ($+3$ vs $+2$) and is usually smaller in size than $\text{Fe}^{2+}$ (because $\text{Fe}^{3+}$ has lost more electrons and its nucleus pulls the remaining electrons tighter).

Key Point: A high charge/size ratio (or high charge density) means the metal ion pulls electron density much more strongly, making the $\text{O-H}$ bonds in the water ligands weaker. Therefore, $\text{[M(H}_2\text{O)}_6\text{]}^{3+}$ ions are significantly more acidic and cause a greater drop in $\text{pH}$ than $\text{[M(H}_2\text{O)}_6\text{]}^{2+}$ ions.

3. Reactions of Aqua Ions with Bases (Precipitation)

When a base is added to a solution of a metal-aqua ion, it removes protons ($\text{H}^+$) from the water ligands. This process continues until the complex becomes neutral, forming an insoluble metal hydroxide precipitate.

The neutral, insoluble precipitate is represented as $\text{[M(H}_2\text{O)}_{6-x}\text{(OH)}_x\text{]}$, where $x$ is equal to the positive charge of the metal ion. For $\text{M}^{2+}$, $x=2$ (forming $\text{M(OH)}_2$). For $\text{M}^{3+}$, $x=3$ (forming $\text{M(OH)}_3$).

3.1 Reaction with Hydroxide Ions ($\text{OH}^-$) (Strong Base)

Adding aqueous sodium hydroxide ($\text{NaOH}$) or potassium hydroxide ($\text{KOH}$) causes an immediate precipitate.

A. Reactions of $\text{M}^{2+}$ (Fe and Cu): $$ \text{[M(H}_2\text{O)}_6\text{]}^{2+} \text{(aq)} + \text{2OH}^- \text{(aq)} \rightarrow \text{M(OH)}_2\text{(H}_2\text{O)}_4\text{(s)} + \text{2H}_2\text{O}\text{(l)} $$

  • $\text{Fe}^{2+}$ (Pale green solution) $\rightarrow$ Green precipitate ($\text{Fe(OH)}_2$).
  • $\text{Cu}^{2+}$ (Pale blue solution) $\rightarrow$ Blue precipitate ($\text{Cu(OH)}_2$).

B. Reactions of $\text{M}^{3+}$ (Fe and Al): $$ \text{[M(H}_2\text{O)}_6\text{]}^{3+} \text{(aq)} + \text{3OH}^- \text{(aq)} \rightarrow \text{M(OH)}_3\text{(H}_2\text{O)}_3\text{(s)} + \text{3H}_2\text{O}\text{(l)} $$

  • $\text{Fe}^{3+}$ (Yellow/Brown solution) $\rightarrow$ Brown/Rust-coloured precipitate ($\text{Fe(OH)}_3$).
  • $\text{Al}^{3+}$ (Colourless solution) $\rightarrow$ White precipitate ($\text{Al(OH)}_3$).


3.2 Reaction with Aqueous Ammonia ($\text{NH}_3$) (Weak Base)

Ammonia acts as a base and removes $\text{H}^+$ ions, causing precipitation in the same way as $\text{OH}^-$.

  • $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$ form the same coloured precipitates ($\text{Fe(OH)}_2$ and $\text{Fe(OH)}_3$) as with $\text{OH}^-$.
  • $\text{Al}^{3+}$ forms the white precipitate $\text{Al(OH)}_3$.

The Copper Exception: Ligand Substitution with excess $\text{NH}_3$

When aqueous ammonia is added to a $\text{Cu}^{2+}$ solution, a pale blue precipitate of $\text{Cu(OH)}_2$ forms initially. However, if excess ammonia is added, the ammonia starts acting as a ligand (instead of a base), displacing some of the water ligands.

This is an important reaction for identifying $\text{Cu}^{2+}$ ions, as the precipitate redissolves to form a deep blue solution: $$ \text{Cu(OH)}_2\text{(H}_2\text{O)}_4\text{(s)} + \text{4NH}_3\text{(aq)} \rightarrow \text{[Cu(NH}_3)_4\text{(H}_2\text{O)}_2\text{]}^{2+} \text{(aq)} + \text{2H}_2\text{O}\text{(l)} + \text{2OH}^- \text{(aq)} $$

Did you know? This deep blue ion, the tetraamminecopper(II) complex, is often used as a confirmatory test for copper in qualitative analysis.


3.3 Reaction with Carbonate Ions ($\text{CO}_3^{2-}$)

Carbonate ions are strong bases. Their reaction with metal-aqua ions is used to distinguish between the divalent ($\text{2+}$) and trivalent ($\text{3+}$) ions.

A. Trivalent Ions ($\text{M}^{3+}$ - Fe, Al): Effervescence Occurs

The $\text{3+}$ ions are highly acidic (high charge density) and are strong enough acids to react fully with the carbonate ion ($\text{CO}_3^{2-}$), releasing carbon dioxide gas ($\text{CO}_2$), which you observe as fizzing (effervescence).

Equation (using $\text{Fe}^{3+}$): $$ \text{2[Fe(H}_2\text{O)}_6\text{]}^{3+} \text{(aq)} + \text{3CO}_3^{2-} \text{(aq)} \rightarrow \text{2Fe(OH)}_3\text{(H}_2\text{O)}_3\text{(s)} + \text{3CO}_2\text{(g)} + \text{3H}_2\text{O}\text{(l)} $$

B. Divalent Ions ($\text{M}^{2+}$ - Fe, Cu): No Effervescence

The $\text{2+}$ ions are not acidic enough to decompose the carbonate ion fully. Instead, they just form an insoluble metal carbonate precipitate, $\text{MCO}_3$, or sometimes the hydroxide, but without the $\text{CO}_2$ effervescence observed for $\text{3+}$ ions.

Memory Aid: Carbonate Test

Think of the $\text{3+}$ ions as being more aggressive (more acidic). They attack the carbonate fully, releasing $\text{CO}_2$. The $\text{2+}$ ions are less aggressive, so they only form a simple precipitate.

  • M$^{3+}$: Precipitate AND Fizzing ($\text{CO}_2$ evolved)
  • M$^{2+}$: Precipitate ONLY (No $\text{CO}_2$ evolved)

4. Amphoteric Character of Metal Hydroxides

Most metal hydroxides are insoluble solids. However, some metal hydroxides show amphoteric character.

Amphoteric means a substance can react and dissolve in both acids (acting as a base) and in excess strong bases (acting as an acid).

Aluminium Hydroxide: The Amphoteric Example

The syllabus requires you to understand that hydroxides of $\text{Al}^{3+}$ are amphoteric.

Initially, adding $\text{OH}^-$ to $\text{Al}^{3+}$ forms the white precipitate $\text{Al(OH)}_3$: $$ \text{[Al(H}_2\text{O)}_6\text{]}^{3+} \text{(aq)} + \text{3OH}^- \text{(aq)} \rightarrow \text{Al(OH)}_3\text{(H}_2\text{O)}_3\text{(s)} + \text{3H}_2\text{O}\text{(l)} $$

A. $\text{Al(OH)}_3$ reacting with Acid (Acting as a Base):

The precipitate dissolves readily in any acid to reform the original aqua ion: $$ \text{Al(OH)}_3\text{(H}_2\text{O)}_3\text{(s)} + \text{3H}^+ \text{(aq)} \rightarrow \text{[Al(H}_2\text{O)}_6\text{]}^{3+} \text{(aq)} $$

B. $\text{Al(OH)}_3$ reacting with Excess Strong Base (Acting as an Acid):

If you add excess strong base (like concentrated $\text{NaOH}$), the precipitate dissolves because the $\text{Al(OH)}_3$ acts as an acid and reacts further, forming a soluble complex ion, usually the tetrahydroxoaluminate ion.

This step is key to identifying $\text{Al}^{3+}$: the precipitate redissolves in excess $\text{OH}^-$. $$ \text{Al(OH)}_3\text{(H}_2\text{O)}_3\text{(s)} + \text{OH}^- \text{(aq)} \rightarrow \text{[Al(OH)}_4\text{]}^{-} \text{(aq)} + \text{3H}_2\text{O}\text{(l)} $$

Key Takeaway: Identifying Ions (Required Practical 9)

Mastering these reactions allows you to identify the ions in the lab. Remember the distinct color changes and the exceptions:

  • $\text{Al}^{3+}$: White precipitate ($\text{Al(OH)}_3$) that redissolves in excess $\text{OH}^-$ (amphoteric). Also fizzes with $\text{CO}_3^{2-}$.
  • $\text{Cu}^{2+}$: Pale blue precipitate ($\text{Cu(OH)}_2$) that redissolves in excess $\text{NH}_3$ (turning deep blue). No fizzing with $\text{CO}_3^{2-}$.
  • $\text{Fe}^{3+}$: Brown/Rust precipitate ($\text{Fe(OH)}_3$). Fizzes with $\text{CO}_3^{2-}$.
  • $\text{Fe}^{2+}$: Green precipitate ($\text{Fe(OH)}_2$). No fizzing with $\text{CO}_3^{2-}$.