Introduction to Transition Metals
Welcome to one of the most colourful and dynamic topics in A2 Chemistry! In this chapter of A2 2 (Analytical, Transition Metals, Electrochemistry and Organic Nitrogen Chemistry), we will explore the fascinating world of the transition metals. From the brilliant blues of copper solutions to the catalysts driving industrial processes like the Haber and Contact processes, transition metals are essential to both laboratory chemistry and global industry.
Don't worry if electron configurations or complex ions felt a little daunting in earlier units. We will break down every single concept step-by-step with clear definitions, visualised structures, and exam-proven memory tips so you can tackle any CCEA exam question with confidence!
1. What is a Transition Metal?
The IUPAC & CCEA Definition
Many students assume that every element in the d-block is automatically a transition metal, but this is a classic exam trap! Make sure you learn the exact CCEA definition word-for-word:
Transition Metal: A d-block element that forms at least one stable ion with a partially filled (incomplete) d-subshell.
Electronic Configurations of Atoms
As you move across Period 4 from Scandium (\(\text{Sc}\)) to Zinc (\(\text{Zn}\)), electrons generally fill the \(4\text{s}\) subshell before the \(3\text{d}\) subshell. However, there are two crucial exceptions you must memorise:
• Chromium (\(\text{Cr}\), \(Z=24\)): \([\text{Ar}]\,3\text{d}^5\,4\text{s}^1\)
Why? Having a half-filled \(3\text{d}\) subshell (\(3\text{d}^5\)) provides extra electronic stability, so one electron is promoted from \(4\text{s}\) to \(3\text{d}\).
• Copper (\(\text{Cu}\), \(Z=29\)): \([\text{Ar}]\,3\text{d}^{10}\,4\text{s}^1\)
Why? Having a completely filled \(3\text{d}\) subshell (\(3\text{d}^{10}\)) provides extra electronic stability.
Forming Transition Metal Ions: The "\(4\text{s}\) First" Rule
When transition metal atoms lose electrons to form positive ions, electrons are always removed from the \(4\text{s}\) subshell before the \(3\text{d}\) subshell.
• \(\text{Fe}\): \([\text{Ar}]\,3\text{d}^6\,4\text{s}^2 \implies \text{Fe}^{2+}\): \([\text{Ar}]\,3\text{d}^6\)
• \(\text{Fe}^{3+}\): \([\text{Ar}]\,3\text{d}^5\)
• \(\text{Cu}\): \([\text{Ar}]\,3\text{d}^{10}\,4\text{s}^1 \implies \text{Cu}^{2+}\): \([\text{Ar}]\,3\text{d}^9\)
Why Scandium and Zinc are NOT Transition Metals
CCEA examiners love asking why \(\text{Sc}\) and \(\text{Zn}\) belong to the d-block but fail the definition of a transition metal:
• Scandium (\(\text{Sc}\)): Forms only one stable ion, \(\text{Sc}^{3+}\). Its electronic configuration is \([\text{Ar}]\,3\text{d}^0\). Because its d-subshell is completely empty, it does not have a partially filled d-subshell.
• Zinc (\(\text{Zn}\)): Forms only one stable ion, \(\text{Zn}^{2+}\). Its electronic configuration is \([\text{Ar}]\,3\text{d}^{10}\). Because its d-subshell is completely full, it also does not have a partially filled d-subshell.
Section Key Takeaway: Always remove \(4\text{s}\) electrons before \(3\text{d}\) electrons when forming ions. A transition metal must have an ion with a \(3\text{d}^1\) to \(3\text{d}^9\) configuration.
2. Characteristic Properties of Transition Metals
Transition metals share four distinctive chemical properties that set them apart from s-block and p-block metals:
1. Variable Oxidation States: Unlike Group 1 or Group 2 metals (which only form \(+1\) or \(+2\) ions), transition metals can form ions with multiple oxidation states (e.g. \(\text{Fe}^{2+}\) and \(\text{Fe}^{3+}\), or \(\text{Mn}\) from \(+2\) up to \(+7\)). This happens because the \(3\text{d}\) and \(4\text{s}\) energy levels are very close in energy, allowing a variable number of electrons to participate in bonding.
2. Formation of Complex Ions: Their small, highly charged ions readily attract molecules or ions with lone pairs to form coordinate bonds.
3. Formation of Coloured Compounds: Hydrated ions and coordination complexes exhibit vibrant colours.
4. Catalytic Activity: Transition metals and their compounds speed up chemical reactions both in industry and in solution.
3. Coordination Chemistry: Complexes and Ligands
Essential Definitions
• Complex (or Complex Ion): A central metal atom or ion surrounded by ligands bonded by coordinate (dative covalent) bonds.
• Ligand: An ion or neutral molecule possessing at least one lone pair of electrons available to donate into vacant orbitals of a central metal ion.
• Coordinate Bond: A covalent bond where both shared electrons originate from the same atom (the ligand).
• Coordination Number: The total number of coordinate bonds formed to the central metal ion (not necessarily the number of ligands!).
Ligand Denticity (How Ligands Bind)
Ligands are classified by the number of coordinate bonds each individual molecule or ion can form:
• Monodentate Ligands: Donate one lone pair to form 1 coordinate bond per ligand.
Examples: \(\text{H}_2\text{O}\), \(\text{NH}_3\), \(\text{Cl}^-\), \(\text{OH}^-\), \(\text{CN}^-\).
• Bidentate Ligands: Donate two lone pairs from two separate atoms on the same molecule/ion to form 2 coordinate bonds.
Examples:
1. 1,2-diaminoethane (\(\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2\), often abbreviated as "en") — each nitrogen atom donates a lone pair.
2. Ethanedioate ion (\(\text{C}_2\text{O}_4^{2-}\) or oxalate) — two oxygen atoms donate lone pairs.
• Polydentate / Hexadentate Ligands (Chelating Agents): Donate multiple lone pairs. The star example is \(\text{EDTA}^{4-}\) (ethylenediaminetetraacetate). One single \(\text{EDTA}^{4-}\) ion forms 6 coordinate bonds to a central metal ion, wrapping around it in a \(1:1\) ratio with ions like \(\text{Cu}^{2+}\), \(\text{Ca}^{2+}\), and \(\text{Mg}^{2+}\).
Geometries and Shapes of Complexes
The shape of a complex depends on its coordination number:
• 6-Coordinate (Octahedral): Bond angles are \(90^\circ\).
Examples: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\), \([\text{Fe}(\text{H}_2\text{O})_6]^{2+}\), \([\text{Cr}(\text{NH}_3)_6]^{3+}\).
• 4-Coordinate (Tetrahedral): Bond angles are \(109.5^\circ\). Typically formed with large ligands like chloride (\(\text{Cl}^-\)) due to steric hindrance.
Examples: \([\text{CuCl}_4]^{2-}\), \([\text{CoCl}_4]^{2-}\).
• 4-Coordinate (Square Planar): Bond angles are \(90^\circ\).
Example: Cisplatin, \([\text{Pt}(\text{NH}_3)_2\text{Cl}_2]\).
• 2-Coordinate (Linear): Bond angle is \(180^\circ\).
Example: Tollens' reagent, \([\text{Ag}(\text{NH}_3)_2]^+\).
Section Key Takeaway: Coordination number equals the number of coordinate bonds, not simply the count of ligand molecules. A complex with 3 bidentate ligands has a coordination number of 6 and an octahedral geometry.
4. The Origin of Colour in Transition Metal Complexes
Why are Transition Metal Ions Coloured?
To understand why these compounds are colourful, let's follow the step-by-step process:
1. Degenerate Orbitals: In an isolated, unbonded gaseous transition metal ion, all five \(3\text{d}\) orbitals have identical energy (they are degenerate).
2. Ligand Repulsion and Splitting: When ligands approach the central metal ion to form coordinate bonds, lone pairs of electrons on the ligands repel the electrons in the metal's \(3\text{d}\) orbitals. Because the d-orbitals point in different directions in 3D space, they experience different amounts of repulsion. This causes the five \(3\text{d}\) orbitals to split into two distinct energy levels separated by an energy gap, \(\Delta E\).
3. Absorption of Light: When visible white light shines on the complex, an electron in a lower-energy d-orbital absorbs a photon of light with a specific frequency (\(\nu\)) or wavelength (\(\lambda\)) that matches the energy gap:
\(\Delta E = h\nu = \frac{hc}{\lambda}\)
This promotes the electron to a higher-energy d-orbital (known as a \(\text{d}\to\text{d}\) electron transition).
4. Transmission of Complementary Colour: The frequencies of light that are not absorbed are transmitted (or reflected) and pass to our eyes. The observed colour is the complementary colour to the wavelength of light that was absorbed.
Why are \(\text{Sc}^{3+}\), \(\text{Ti}^{4+}\), \(\text{Zn}^{2+}\), and \(\text{Cu}^+\) Colourless?
• \(\text{Sc}^{3+}\) and \(\text{Ti}^{4+}\) have a \(3\text{d}^0\) configuration: there are no d-electrons to promote.
• \(\text{Zn}^{2+}\) and \(\text{Cu}^+\) have a \(3\text{d}^{10}\) configuration: the d-subshell is completely full, so there is no vacant higher d-orbital for an electron to move into.
Without a \(\text{d}\to\text{d}\) transition, visible light cannot be absorbed in this way, making these solutions colourless and their solid salts white.
Examiner Warning: Never state that colour is produced because light is emitted when electrons drop back down! That is the mechanism for atomic emission flame tests. For transition metal complexes, colour arises from the absorption of specific frequencies of visible light, leaving the transmitted complementary light to be seen.
5. Core Aqueous Ions and Complex Colours
You are expected to know the colours of specific ions and oxidation states for the CCEA examination:
Vanadium (\(\text{V}\)) Oxidation States
When ammonium vanadate(V) is reduced by zinc metal in acidic solution (\(\text{Zn}/\text{H}^+\)), it passes through four distinct oxidation states in sequence:
• \(\text{VO}_2^+\) (Vanadium(V)): Yellow
• \(\text{VO}^{2+}\) (Vanadium(IV)): Blue
• \(\text{V}^{3+}\) (Vanadium(III)): Green
• \(\text{V}^{2+}\) (Vanadium(II)): Violet / Lilac
Memory Aid for the reduction sequence: You Better Get Vanadium (Yellow \(\to\) Blue \(\to\) Green \(\to\) Violet).
Chromium (\(\text{Cr}\))
• Dichromate(VI) ion, \(\text{Cr}_2\text{O}_7^{2-}\): Orange
• Chromate(VI) ion, \(\text{CrO}_4^{2-}\): Yellow
• Aqueous chromium(III), \([\text{Cr}(\text{H}_2\text{O})_6]^{3+}\): Green (ruby-violet when pure hexaaqua, but green is standard in aqueous solution and exam answers)
• Chromium(II), \(\text{Cr}^{2+}\): Blue
Manganese (\(\text{Mn}\))
• Manganate(VII) ion, \(\text{MnO}_4^-\): Deep Purple
• Manganese(II) ion, \(\text{Mn}^{2+}\): Very pale pink (often appears virtually colourless in dilute solution)
Iron (\(\text{Fe}\))
• Hexaaquairon(II), \([\text{Fe}(\text{H}_2\text{O})_6]^{2+}\): Pale green
• Hexaaquairon(III), \([\text{Fe}(\text{H}_2\text{O})_6]^{3+}\): Yellow / Orange-brown
Cobalt (\(\text{Co}\)) & Nickel (\(\text{Ni}\))
• \([\text{Co}(\text{H}_2\text{O})_6]^{2+}\): Pink
• \([\text{CoCl}_4]^{2-}\): Blue
• \([\text{Ni}(\text{H}_2\text{O})_6]^{2+}\): Green
Copper (\(\text{Cu}\))
• \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\): Pale blue
• \([\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}\): Deep blue
• \([\text{CuCl}_4]^{2-}\): Yellow-green / Yellow (appears olive-green if mixed with unreacted blue hydrated ions)
6. Ligand Substitution Reactions
A ligand substitution (replacement) reaction occurs when one ligand in a complex ion is replaced by a different ligand.
1. Copper(II) with Concentrated Hydrochloric Acid (\(\text{HCl}\))
Adding concentrated \(\text{HCl}\) (a source of \(\text{Cl}^-\) ions) to pale blue aqueous copper(II) sulfate causes a reversible substitution:
\([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{Cl}^-(\text{aq}) \rightleftharpoons [\text{CuCl}_4]^{2-}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\)
• Colour change: Pale blue solution turns yellow-green / yellow.
• Change in coordination number: Decreases from 6 (octahedral) to 4 (tetrahedral).
Why does coordination number change? Chloride ions (\(\text{Cl}^-\)) are significantly larger than neutral water molecules (\(\text{H}_2\text{O}\)). Only four \(\text{Cl}^-\) ions can fit around the central \(\text{Cu}^{2+}\) ion due to steric hindrance and mutual electrostatic repulsion.
2. Copper(II) with Aqueous Ammonia (\(\text{NH}_3\))
This reaction occurs in two distinct stages depending on how much ammonia is added:
• Stage 1 (Dropwise addition of \(\text{NH}_3\)): Ammonia acts as a Brønsted-Lowry base, removing \(\text{H}^+\) ions from water ligands to form a pale blue precipitate of copper(II) hydroxide:
\([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 2\text{NH}_3(\text{aq}) \to \text{Cu}(\text{OH})_2(\text{s}) + 2\text{NH}_4^+(\text{aq}) + 4\text{H}_2\text{O}(\text{l})\)
• Stage 2 (Addition of excess \(\text{NH}_3\)): The precipitate dissolves as ammonia acts as a ligand, replacing four water molecules to form a deep blue solution:
\([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{NH}_3(\text{aq}) \rightleftharpoons [\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}(\text{aq}) + 4\text{H}_2\text{O}(\text{l})\)
The coordination number remains 6 (octahedral geometry).
3. The Chelate Effect
When monodentate ligands (such as \(\text{H}_2\text{O}\)) are replaced by multidentate ligands (such as \(\text{EDTA}^{4-}\) or 1,2-diaminoethane), the reaction is strongly favoured thermodynamically:
\([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + \text{EDTA}^{4-}(\text{aq}) \rightleftharpoons [\text{Cu}(\text{EDTA})]^{2-}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\)
• Notice the stoichiometry: 2 species on the left produce 7 species on the right!
• This massive increase in the number of free particles leads to a large positive change in system entropy (\(\Delta S_{\text{system}} > 0\)).
• Because \(\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ\), a positive \(\Delta S^\circ\) makes the free energy change \(\Delta G^\circ\) much more negative, driving the equilibrium far to the right. This phenomenon is known as the chelate effect.
7. Hydroxide Precipitation Reactions (Qualitative Analysis)
When aqueous sodium hydroxide (\(\text{NaOH}\)) is added dropwise and then in excess to aqueous transition metal ions, distinctive hydroxide precipitates form:
• Chromium(III), \(\text{Cr}^{3+}\):
Forms a green precipitate of \(\text{Cr}(\text{OH})_3\). It is amphoteric and dissolves in excess \(\text{NaOH}\) to form a dark green solution of hexahydroxochromate(III), \([\text{Cr}(\text{OH})_6]^{3-}\).
• Iron(II), \(\text{Fe}^{2+}\):
Forms a green precipitate of \(\text{Fe}(\text{OH})_2\). Insoluble in excess \(\text{NaOH}\). On standing exposed to air, the surface turns brown because atmospheric oxygen oxidises \(\text{Fe}(\text{OH})_2\) to \(\text{Fe}(\text{OH})_3\).
• Iron(III), \(\text{Fe}^{3+}\):
Forms a rust-brown precipitate of \(\text{Fe}(\text{OH})_3\). Insoluble in excess \(\text{NaOH}\).
• Manganese(II), \(\text{Mn}^{2+}\):
Forms an off-white / pale brown precipitate of \(\text{Mn}(\text{OH})_2\). Insoluble in excess \(\text{NaOH}\). It darkens to brown on contact with air due to oxidation.
• Copper(II), \(\text{Cu}^{2+}\):
Forms a pale blue precipitate of \(\text{Cu}(\text{OH})_2\). Insoluble in excess \(\text{NaOH}\).
8. Catalysis
Transition metals make exceptional catalysts because of their ability to change oxidation states readily and adsorb molecules onto their d-orbital active surfaces.
A. Heterogeneous Catalysis
Definition: The catalyst is in a different physical state (phase) from the reactants (usually a solid catalyst with gaseous reactants).
How it works (Surface Adsorption Mechanism):
1. Adsorption: Reactant molecules form weak bonds with active sites on the solid catalyst surface.
2. Reaction: Bonds within the reactant molecules weaken and break; new bonds form between reactant fragments.
3. Desorption: The product molecules leave the surface, freeing up the active sites for new reactant molecules.
Key Examples:
• Haber Process: Solid iron (\(\text{Fe}(\text{s})\)) catalyses the synthesis of ammonia:
\(\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})\)
• Contact Process: Solid vanadium(V) oxide (\(\text{V}_2\text{O}_5(\text{s})\)) catalyses the oxidation of sulfur dioxide to sulfur trioxide:
\(2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g})\)
Two-step mechanism involving variable oxidation states of vanadium:
Step 1: \(\text{SO}_2 + \text{V}_2\text{O}_5 \to \text{SO}_3 + \text{V}_2\text{O}_4\)
Step 2: \(\text{V}_2\text{O}_4 + \frac{1}{2}\text{O}_2 \to \text{V}_2\text{O}_5\)
B. Homogeneous Catalysis
Definition: The catalyst is in the same physical state (phase) as the reactants (usually all in aqueous solution).
Key Examples:
• Peroxodisulfate and Iodide Reaction:
\(\text{S}_2\text{O}_8^{2-}(\text{aq}) + 2\text{I}^-(\text{aq}) \to 2\text{SO}_4^{2-}(\text{aq}) + \text{I}_2(\text{aq})\)
Why uncatalysed is slow: Both \(\text{S}_2\text{O}_8^{2-}\) and \(\text{I}^-\) are negatively charged anions, so high electrostatic repulsion creates a very high activation energy.
Catalysed by \(\text{Fe}^{2+}\) (or \(\text{Fe}^{3+}\)):
Step 1: \(\text{S}_2\text{O}_8^{2-} + 2\text{Fe}^{2+} \to 2\text{SO}_4^{2-} + 2\text{Fe}^{3+}\) (opposite charges attract!)
Step 2: \(2\text{Fe}^{3+} + 2\text{I}^- \to 2\text{Fe}^{2+} + \text{I}_2\)
Because every step involves oppositely charged ions, the activation energy is significantly lower.
• Autocatalysis in the Manganate(VII) / Ethanedioate Titration:
Reaction: \(2\text{MnO}_4^- + 16\text{H}^+ + 5\text{C}_2\text{O}_4^{2-} \to 2\text{Mn}^{2+} + 8\text{H}_2\text{O} + 10\text{CO}_2\)
Autocatalysis occurs when a product of the reaction acts as a catalyst for that same reaction.
The reaction is initially very slow because both \(\text{MnO}_4^-\) and \(\text{C}_2\text{O}_4^{2-}\) are negative anions. However, as soon as some \(\text{Mn}^{2+}\) ions are produced, they act as an efficient homogeneous catalyst, and the rate of reaction accelerates dramatically.
Quick Review: Top 5 Examiner Pitfalls to Avoid
1. Definition of a Transition Metal: Do not simply say "an element in the d-block." You must specify that it forms at least one stable ion with a partially filled d-subshell.
2. Electron Configuration of Ions: Always empty the \(4\text{s}\) orbital before the \(3\text{d}\) orbitals (e.g. \(\text{Fe}^{2+}\) is \([\text{Ar}]\,3\text{d}^6\), never \([\text{Ar}]\,3\text{d}^4\,4\text{s}^2\)).
3. Why Complexes are Coloured: State clearly that visible light is absorbed to promote a d-electron to a higher energy d-orbital (\(\text{d}\to\text{d}\) transition), and the transmitted/complementary colour is seen. Never say light is emitted!
4. Coordination Number: Remember that coordination number counts coordinate bonds, not the number of ligand molecules (e.g. \([\text{Cu}(\text{EDTA})]^{2-}\) has 1 ligand, but its coordination number is 6).
5. Ammonia Reactions: Distinguish between dropwise \(\text{NH}_3\) (acting as a base to form a hydroxide precipitate) and excess \(\text{NH}_3\) (acting as a ligand to form a soluble deep-blue complex).