Welcome to the World of Ions in Solution!
In this chapter, we explore what happens when metal salts dissolve in water. You'll learn how to identify different metal ions just by looking at how they react in a test tube. This is like being a chemical detective! Understanding these reactions is vital for everything from industrial processing to understanding how minerals move through our environment.
1. Metal-Aqua Ions: What Happens When Salts Dissolve?
When you dissolve a transition metal salt (like Copper Sulfate) in water, the metal ion doesn't just float around alone. It gets surrounded by water molecules. These water molecules act as ligands, donating a lone pair of electrons to the metal ion to form co-ordinate bonds.
According to your AQA syllabus, you need to know four specific metal-aqua ions. In all of these, the metal is surrounded by six water molecules, creating an octahedral shape:
The 2+ Ions:
1. \([Fe(H_2O)_6]^{2+}\) - Iron(II) (Pale green solution)
2. \([Cu(H_2O)_6]^{2+}\) - Copper(II) (Blue solution)
The 3+ Ions:
3. \([Al(H_2O)_6]^{3+}\) - Aluminium(III) (Colourless solution)
4. \([Fe(H_2O)_6]^{3+}\) - Iron(III) (Purple/Yellow-brown solution)
Note: Pure \([Fe(H_2O)_6]^{3+}\) is actually purple, but due to some hydrolysis, it usually looks yellow-brown in the lab!
Quick Review Box:
Remember: "Aqua" means water. A "hexaaqua" ion simply means a metal ion with 6 water molecules stuck to it.
2. Why are 3+ Ions more Acidic than 2+ Ions?
If you test the pH of these solutions, you'll find the 3+ ions are more acidic than the 2+ ions. But why? This is a favorite exam question!
It all comes down to the charge/size ratio (sometimes called charge density).
3+ ions are smaller and have a higher charge than 2+ ions. This means they are much better at "pulling" electrons towards themselves.
The "Tug-of-War" Analogy
Imagine the metal ion is playing tug-of-war with the electrons in the \(O-H\) bond of the water ligand. Because the 3+ ion is so "greedy" for electrons (high charge density), it pulls the electrons in the water molecule's \(O-H\) bond very strongly toward the oxygen and the metal. This weakens the O-H bond, making it much easier for a hydrogen ion (\(H^+\)) to break off and pop into the solution.
The Equation for Hydrolysis:
\([M(H_2O)_6]^{3+}(aq) \rightleftharpoons [M(H_2O)_5(OH)]^{2+}(aq) + H^+(aq)\)
Key Takeaway:
Higher charge density (3+) = More electron pulling = Weaker \(O-H\) bond = More \(H^+\) released = Lower pH (more acidic).
3. Reactions with Bases (Sodium Hydroxide and Ammonia)
When we add a base like Sodium Hydroxide (\(NaOH\)) or Ammonia (\(NH_3\)), we are essentially removing protons (\(H^+\)) from the water ligands until a neutral, insoluble precipitate forms.
A) Adding OH⁻ (Sodium Hydroxide)
In all four cases, adding $OH^-$ step-by-step eventually results in a solid metal hydroxide precipitate.
1. Iron(II): \([Fe(H_2O)_6]^{2+} + 2OH^- \rightarrow [Fe(H_2O)_4(OH)_2](s) + 2H_2O\)
Observation: Green precipitate (turns brown on standing as it oxidises).
2. Copper(II): \([Cu(H_2O)_6]^{2+} + 2OH^- \rightarrow [Cu(H_2O)_4(OH)_2](s) + 2H_2O\)
Observation: Blue precipitate.
3. Iron(III): \([Fe(H_2O)_6]^{3+} + 3OH^- \rightarrow [Fe(H_2O)_3(OH)_3](s) + 3H_2O\)
Observation: Brown precipitate.
4. Aluminium(III): \([Al(H_2O)_6]^{3+} + 3OH^- \rightarrow [Al(H_2O)_3(OH)_3](s) + 3H_2O\)
Observation: White precipitate.
B) The "Amphoteric" Exception (Aluminium)
Don't worry if this sounds complex! "Amphoteric" just means it can act as both an acid and a base.
If you add excess \(NaOH\) to the Aluminium precipitate, it dissolves!
\([Al(H_2O)_3(OH)_3](s) + OH^-(aq) \rightarrow [Al(OH)_4]^-(aq) + 3H_2O\)
Observation: The white precipitate dissolves to form a colourless solution.
C) Adding Ammonia (NH₃)
Ammonia acts as a base initially, producing the same precipitates as \(OH^-\). However, Copper has a special trick called ligand substitution if you add excess Ammonia.
Copper(II) in Excess NH₃:
\([Cu(H_2O)_4(OH)_2](s) + 4NH_3 \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+}(aq) + 2OH^- + 2H_2O\)
Observation: The blue precipitate dissolves to form a deep blue solution.
Common Mistake to Avoid:
Students often think all precipitates dissolve in excess. Only Aluminium dissolves in excess \(NaOH\), and only Copper dissolves in excess \(NH_3\).
4. Reactions with Carbonates (CO₃²⁻)
This is where the difference between 2+ and 3+ ions becomes really obvious. Carbonates like to fizz with acids.
Reaction with 2+ Ions (Fe²⁺ and Cu²⁺)
These are not acidic enough to make the carbonate "fizz." They simply swap ions to form a metal carbonate solid.
\([M(H_2O)_6]^{2+} + CO_3^{2-} \rightarrow MCO_3(s) + 6H_2O\)
Observations: FeCO₃ (Green precipitate); CuCO₃ (Blue-green precipitate).
Reaction with 3+ Ions (Fe³⁺ and Al³⁺)
Because these solutions are acidic, they react with the carbonate to produce Carbon Dioxide gas. You will see bubbles!
\(2[M(H_2O)_6]^{3+} + 3CO_3^{2-} \rightarrow 2[M(H_2O)_3(OH)_3](s) + 3CO_2(g) + 3H_2O\)
Observations: Bubbling (effervescence) and a precipitate (Brown for Fe, White for Al).
Memory Aid:
"3+ is a G": 3+ ions produce Gas (effervescence) with carbonates. 2+ ions do not!
5. Summary Cheat Sheet
Use this table to quickly review the observations for the four main ions:
Ion: \([Fe(H_2O)_6]^{2+}\)
+ \(NaOH\): Green ppt
+ \(NH_3\): Green ppt
+ \(Na_2CO_3\): Green ppt (No fizz)
Ion: \([Cu(H_2O)_6]^{2+}\)
+ \(NaOH\): Blue ppt
+ \(NH_3\): Blue ppt (Dissolves in excess to Deep Blue Solution)
+ \(Na_2CO_3\): Blue-green ppt (No fizz)
Ion: \([Fe(H_2O)_6]^{3+}\)
+ \(NaOH\): Brown ppt
+ \(NH_3\): Brown ppt
+ \(Na_2CO_3\): Brown ppt + Fizzing
Ion: \([Al(H_2O)_6]^{3+}\)
+ \(NaOH\): White ppt (Dissolves in excess to colourless solution)
+ \(NH_3\): White ppt
+ \(Na_2CO_3\): White ppt + Fizzing
Final Tip for Success:
When writing equations for these reactions, always count your charges! The total charge on the left must equal the total charge on the right. If you struggle with the big "aqua" formulas, practice drawing them out first to visualize the water molecules being replaced by \(OH\) groups.