Chemistry | Transition Metals and their Chemistry

🎨 Transition Metals and Their Chemistry: Comprehensive Study Notes 💡

Hello future Chemists! Welcome to one of the most colourful and exciting chapters in Inorganic Chemistry: Transition Metals. These elements are the chameleons of the Periodic Table—they change colours, change their mind (oxidation state), and are essential for everything from life processes to industrial manufacturing.

Don't worry if this chapter seems tricky at first. We will break down complex ideas like complex ions and d-d transitions into simple, manageable steps. By the end, you will understand why transition metals are the superheroes of the d-block!

Section 1: What Makes a Transition Metal "Transitional"?

1.1 Definition and Location

Transition metals are the elements found in the d-block of the Periodic Table (Groups 3 to 12).

The official definition, according to chemists, is specific:

Transition Metal: An element that forms at least one stable ion with a partially filled d-subshell.

This definition usually excludes Scandium (Sc) and Zinc (Zn).

• Scandium only forms $\text{Sc}^{3+}$, which has zero d-electrons (empty d-subshell).
• Zinc only forms $\text{Zn}^{2+}$, which has ten d-electrons (full d-subshell).

1.2 Electron Configuration: The Rule Breakers

Transition metals fill the 4s orbital before the 3d orbital (following the Aufbau Principle). However, when they form ions, the electrons are always removed from the 4s orbital first, because the 4s orbital becomes higher in energy than the 3d orbital once electrons are present.

Example: Iron (Fe, atomic number 26)

Neutral Fe: \( \text{[Ar] } 3d^6 4s^2 \)
$\text{Fe}^{2+}$ Ion: \( \text{[Ar] } 3d^6 \) (Two electrons removed from 4s)
$\text{Fe}^{3+}$ Ion: \( \text{[Ar] } 3d^5 \) (Two from 4s, one from 3d)

🚨 Common Mistakes to Avoid: Remember that when you ionise, you *always* remove 4s electrons before 3d electrons!

The Exceptions (Chromium and Copper)

To achieve a more stable configuration (half-filled or fully-filled subshell), Chromium and Copper break the usual filling rule for their neutral atoms:

• Chromium (Cr): Instead of $3d^4 4s^2$, it is $3d^5 4s^1$ (stable half-filled d-subshell).
• Copper (Cu): Instead of $3d^9 4s^2$, it is $3d^{10} 4s^1$ (stable fully-filled d-subshell).

1.3 Key Characteristics

The partially filled d-subshell is the reason transition metals possess their unique chemistry:

1. Variable Oxidation States: They can easily lose different numbers of electrons (e.g., Iron can be +2 or +3).
2. Formation of Coloured Ions and Compounds: Due to d-d electron transitions (explained in Section 4).
3. Ability to Form Complex Ions: They act as central metal atoms bonded to ligands.
4. Catalytic Activity: They accelerate chemical reactions (both homogeneous and heterogeneous).

Section 2: The World of Complex Ions

2.1 What is a Complex Ion?

A Complex Ion (or coordination complex) consists of a central metal ion (the transition metal) bonded to one or more molecules or ions called ligands.

The bonds formed between the ligand and the central metal ion are coordinate (dative covalent) bonds, where both electrons in the bond come from the ligand.

Analogy: Think of the central metal ion as a hungry spider waiting for flies (ligands). The ligands (flies) donate their electron pairs (food) into the empty orbitals of the metal ion.

2.2 Ligands and Coordination Number

Ligands: Species (molecules or ions) that have at least one lone pair of electrons they can donate to the central metal ion. Common examples include: $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{Cl}^-$, and $\text{CN}^-$.

Coordination Number (CN): This is simply the total number of coordinate bonds formed between the ligands and the central metal atom.

Example: In the complex ion \( \text{[Cu}(\text{H}_2\text{O})_6\text{]}^{2+} \), the central ion is $\text{Cu}^{2+}$, the ligand is $\text{H}_2\text{O}$, and the Coordination Number is 6.

Types of Ligands

• Monodentate (Unidentate) Ligands: Donate one lone pair, forming one bond (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{Cl}^-$).
• Bidentate Ligands: Donate two lone pairs, forming two bonds (e.g., ethane-1,2-diamine (en) or ethanedioate ($\text{C}_2\text{O}_4^{2-}$)).
• Polydentate Ligands: Donate multiple lone pairs (more than two). The most important example is EDTA$^{4-}$ (ethylenediaminetetraacetate), which is hexadentate (forms 6 bonds!).

The Chelate Effect: Complexes formed with bidentate or polydentate ligands are called chelates. These complexes are significantly more stable than complexes formed with only monodentate ligands. This enhanced stability is driven by a favourable increase in entropy (more particles/disorder in solution).

2.3 Shapes of Complex Ions

The shape is dictated by the coordination number (CN):

1. CN = 6 (Most Common): Octahedral shape. Ligands are placed at the vertices of an octahedron (e.g., \( \text{[Fe}(\text{H}_2\text{O})_6\text{]}^{2+} \)). Bond angles are $90^\circ$.
2. CN = 4 (Common): Two possible shapes:
   • Tetrahedral: Found mostly with larger ligands or when the metal ion is not highly charged (e.g., $\text{Zn}^{2+}$ or $\text{Cu}^{+}$ complexes like \( \text{[CuCl}_4\text{]}^{2-} \)). Bond angles are $109.5^\circ$.
   • Square Planar: Typically found with metals having $d^8$ configuration (often $\text{Pt(II)}$ and $\text{Pd(II)}$). Bond angles are $90^\circ$. (Cisplatin is a key example here!)
3. CN = 2 (Rare): Linear shape. (e.g., some $\text{Ag(I)}$ complexes). Bond angle $180^\circ$.

Section 3: Isomerism in Complex Ions

Isomers are compounds that have the same chemical formula but different arrangements of atoms. Transition metal complexes can exhibit two main types of stereoisomerism: geometrical and optical.

3.1 Geometrical Isomerism (Cis/Trans)

This type of isomerism occurs in square planar and octahedral complexes when two identical ligands can be arranged in different spatial positions.

• Cis Isomer: Identical ligands are next to each other (at $90^\circ$).
• Trans Isomer: Identical ligands are opposite each other (at $180^\circ$).

Example: Square Planar Complex \( \text{[Pt(NH}_3\text{)}_2\text{Cl}_2\text{]} \)

The cis isomer, Cisplatin, is a crucial anti-cancer drug, while the trans isomer has no medicinal value. This shows how crucial spatial arrangement is!

3.2 Optical Isomerism (Enantiomers)

Optical isomerism occurs when a complex ion is non-superimposable on its mirror image. These pairs are called enantiomers and they rotate plane-polarised light in opposite directions.

This is most common in octahedral complexes containing bidentate ligands (like 'en').

Analogy: Your hands are optical isomers. They are mirror images of each other, but you cannot perfectly superimpose your right hand onto your left hand (the thumb is always in the wrong place!).

Section 4: The Origin of Colour

4.1 The Role of the d-Orbitals

Most transition metal compounds are coloured because their ions have partially filled d-orbitals.

In an isolated metal atom, the five d-orbitals all have the same energy (they are degenerate). However, when ligands approach the central metal ion to form coordinate bonds, the electrons (which are negative) repel the electrons in the metal's d-orbitals.

This repulsion causes the five d-orbitals to split into two different energy levels. This is called d-orbital splitting.

4.2 The Colour Mechanism (d-d Transitions)

1. Absorption: When white light (which contains all colours/wavelengths) hits the complex ion, an electron in the lower energy d-orbital absorbs a specific frequency/wavelength of visible light.
2. Excitation: This energy promotes the electron to the higher energy d-orbital (this jump is called a d-d transition).
3. Transmission/Reflection: The frequency of light that was absorbed is 'missing' from the light transmitted or reflected back to your eye.
4. Observed Colour: The colour you see is the complementary colour to the one absorbed.

Example: If a complex absorbs yellow light, you will see blue light (its complement).

4.3 Factors Affecting Colour

The colour of a complex depends on the size of the energy gap ($\Delta E$) between the split d-orbitals, which is determined by:

1. The Identity of the Ligand: Stronger ligands (like $\text{CN}^-$ or $\text{NH}_3$) cause a larger energy gap ($\Delta E$) than weaker ligands (like $\text{H}_2\text{O}$ or $\text{Cl}^-$). Changing the ligand changes the amount of energy absorbed, thus changing the colour.
2. The Oxidation State of the Metal Ion: A higher positive charge on the metal ion generally causes greater d-orbital splitting, leading to different colours (e.g., $\text{Fe}^{2+}$ is pale green, $\text{Fe}^{3+}$ is yellow/brown).
3. The Coordination Number/Geometry: Changing the shape (e.g., from octahedral to tetrahedral) changes how the ligands approach the metal ion, which fundamentally alters the way the d-orbitals split, resulting in a new colour.

Section 5: Reactions of Transition Metal Ions in Aqueous Solution

Aqueous transition metal ions exist as complexes where water molecules act as ligands, e.g., \( \text{[Fe}(\text{H}_2\text{O})_6\text{]}^{2+} \). Their reactions often involve exchanging these ligands or changing the oxidation state.

5.1 Ligand Substitution

This is a reaction where one set of ligands is replaced by another. This is often an equilibrium reaction, and the colour changes dramatically because the new ligand causes a different d-orbital split (Section 4.3).

A. Substitution with Ammonia ($\text{NH}_3$)

When concentrated ammonia solution is added to an aqueous copper(II) solution (light blue):

\( \text{[Cu}(\text{H}_2\text{O})_6\text{]}^{2+} \, (\text{aq, light blue}) + 4\text{NH}_3 \, (\text{aq}) \rightleftharpoons \text{[Cu}(\text{NH}_3\text{)}_4(\text{H}_2\text{O})_2\text{]}^{2+} \, (\text{aq, deep blue}) + 4\text{H}_2\text{O} \, (\text{l}) \)

Note: Only four water ligands are replaced by ammonia in this specific copper complex.

B. Substitution with Chloride Ions ($\text{Cl}^-$)

Chloride ions are larger than water molecules, so substitution often leads to a change in coordination number and shape (usually from octahedral CN=6 to tetrahedral CN=4).

Example: Cobalt(II) solution (pink) turning blue upon addition of concentrated $\text{HCl}$ or concentrated $\text{NaCl}$:

\( \text{[Co}(\text{H}_2\text{O})_6\text{]}^{2+} \, (\text{aq, pink}) + 4\text{Cl}^- \, (\text{aq}) \rightleftharpoons \text{[CoCl}_4\text{]}^{2-} \, (\text{aq, blue}) + 6\text{H}_2\text{O} \, (\text{l}) \)

5.2 Reaction with Aqueous Sodium Hydroxide ($\text{OH}^-$)

When aqueous sodium hydroxide or aqueous ammonia is added to solutions of transition metal ions, the hydroxide ion acts as a base and deprotonates the coordinated water ligands, forming an insoluble metal hydroxide precipitate.

General Equation (using $\text{M}^{2+}$): $$ \text{[M}(\text{H}_2\text{O})_6\text{]}^{2+} \, (\text{aq}) + 2\text{OH}^- \, (\text{aq}) \rightarrow \text{M}(\text{OH})_2(\text{H}_2\text{O})_4 \, (\text{s}) + 2\text{H}_2\text{O} \, (\text{l}) $$

Key Precipitation Observations (Must Know!)

• $\text{Fe}^{2+}$ (aq): Forms a dirty green precipitate ($\text{Fe(OH)}_2$).
• $\text{Fe}^{3+}$ (aq): Forms a rusty brown precipitate ($\text{Fe(OH)}_3$).
• $\text{Cu}^{2+}$ (aq): Forms a blue precipitate ($\text{Cu(OH)}_2$).
• $\text{Cr}^{3+}$ (aq): Forms a green precipitate ($\text{Cr(OH)}_3$). This precipitate is amphoteric—it redissolves in excess $\text{OH}^-$ or excess $\text{H}^+$ to form a soluble complex ion.

5.3 Redox Chemistry and Variable Oxidation States

Transition metals are great redox agents because the energy required to change their oxidation state is often small.

Key examples often tested: Iron and Chromium.

• Iron: The $\text{Fe}^{2+}$ ion is easily oxidised to the more stable $\text{Fe}^{3+}$ ion.
• Chromium: $\text{Cr}^{3+}$ is very stable, but it can be oxidised to $\text{Cr}^{6+}$ (in $\text{CrO}_4^{2-}$ or $\text{Cr}_2\text{O}_7^{2-}$ ions). Chromate ($\text{CrO}_4^{2-}$, yellow) and Dichromate ($\text{Cr}_2\text{O}_7^{2-}$, orange) exist in equilibrium, which is pH dependent.
  $$ 2\text{CrO}_4^{2-} \, (\text{yellow}) + 2\text{H}^+ \rightleftharpoons \text{Cr}_2\text{O}_7^{2-} \, (\text{orange}) + \text{H}_2\text{O} $$

Section 6: Transition Metals as Catalysts

A catalyst increases the rate of a reaction without being chemically changed itself. Transition metals and their compounds are excellent catalysts due to their ability to:

1. Use their partially filled d-orbitals to form weak bonds with reactants.
2. Easily access multiple oxidation states.

6.1 Heterogeneous Catalysis

The catalyst is in a different phase from the reactants (e.g., solid catalyst, gaseous reactants).

Mechanism: Reactants adsorb onto the catalyst surface, weakening their bonds. Reaction occurs, and products desorb.

Examples:

• Iron (Fe) in the Haber Process ($\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$).
• Vanadium(V) Oxide ($\text{V}_2\text{O}_5$) in the Contact Process ($2\text{SO}_2 + \text{O}_2 \rightarrow 2\text{SO}_3$).

6.2 Homogeneous Catalysis

The catalyst and reactants are in the same phase, typically in solution (liquid).

Mechanism: The catalyst acts via a cycle of redox changes, involving the formation of an intermediate compound.

Example: The reaction between Iodide ($\text{I}^-$) and Persulfate ($\text{S}_2\text{O}_8^{2-}$) ions, which is usually very slow. This is catalysed by $\text{Fe}^{2+}$ or $\text{Fe}^{3+}$ ions:

1. $\text{Fe}^{2+}$ reduces the persulfate ion: $$ \text{S}_2\text{O}_8^{2-} + 2\text{Fe}^{2+} \rightarrow 2\text{SO}_4^{2-} + 2\text{Fe}^{3+} $$
2. The $\text{Fe}^{3+}$ then oxidises the iodide ion, regenerating the catalyst: $$ 2\text{Fe}^{3+} + 2\text{I}^- \rightarrow 2\text{Fe}^{2+} + \text{I}_2 $$

The overall reaction is simply: $$ \text{S}_2\text{O}_8^{2-} + 2\text{I}^- \rightarrow 2\text{SO}_4^{2-} + \text{I}_2 $$

The $\text{Fe}^{2+}$ (or $\text{Fe}^{3+}$) allows the reaction to occur through two steps with lower activation energy than the single uncatalysed step.