Introduction to Transition Metals
Transition elements are those which have incomplete d-subshell in their most stable ionic states. They are called ‘transition’ because they show gradual change from highly reactive group 2 elements to less reactive group 13 elements.
Remember! Zinc is a d-block element, but not a transition element because it forms only one colourless ion, Zn²⁺, with a complete 3d sub-shell.
Definition
A transition element is defined as a d-block element which forms one or more stable ions with incomplete d orbitals.
Tip: To identify transition elements, check if they form ions with incomplete d-orbitals. Zinc (Zn²⁺) and Scandium (Sc³⁺) have empty or full d-orbitals, so they don’t qualify as true transition metals.
Importance and Applications
- Copper: Used in wiring, coins, and plumbing
- Iron: Essential for bridges, vehicle parts, and construction
- Chromium: Enhances plumbing fixtures
- Gold and Silver: Valuable for jewelry and electronics
- Platinum: Aids in catalytic converters
- Titanium: Used in bicycles, aircraft, and artificial joints
- Nickel, Vanadium, Molybdenum, Tantalum: Various industrial and medical purposes
- Alloys: Brass (copper and zinc) and bronze (copper and tin)
Note on Scandium: Scandium was initially controversial as its common +3 oxidation state (Sc³⁺) has an empty d sub-shell. However, it was later included when compounds with +1 and +2 oxidation states were synthesized.
First Row Transition Elements
The first period of transition elements starts from scandium to zinc, but we mainly focus on elements from titanium through to copper.
Electron Configurations of D-Block Elements
The electronic configuration of transition metals generally follows the Aufbau principle, with exceptions for chromium and copper.
Key Information: When forming cations, electrons are always removed from the higher energy sub-shell, i.e., 4s in first row of d-block elements.
Exceptions to Aufbau Principle
Chromium and copper break the Aufbau principle because atoms or ions with:
- Half-filled 3d sub-level (3d⁵) – relatively stable
- Filled 3d sub-level (3d¹⁰) – relatively stable
These configurations avoid inter-electronic repulsion in the 4s orbital and have symmetrical 3d electron clouds.
Electron Configurations Table
| Element | Electronic Configuration | Ion | Electronic Configuration |
|---|---|---|---|
| Sc | [Ar] 3d¹ 4s² | Sc²⁺ | [Ar] 3d¹ |
| Ti | [Ar] 3d² 4s² | Ti²⁺ | [Ar] 3d² |
| V | [Ar] 3d³ 4s² | V⁴⁺ | [Ar] 3d¹ |
| Cr | [Ar] 3d⁵ 4s¹ | Cr²⁺ | [Ar] 3d⁴ |
| Mn | [Ar] 3d⁵ 4s² | Mn²⁺ | [Ar] 3d⁵ |
| Fe | [Ar] 3d⁶ 4s² | Fe²⁺ | [Ar] 3d⁶ |
| Co | [Ar] 3d⁷ 4s² | Co²⁺ | [Ar] 3d⁷ |
| Ni | [Ar] 3d⁸ 4s² | Ni²⁺ | [Ar] 3d⁸ |
| Cu | [Ar] 3d¹⁰ 4s¹ | Cu²⁺ | [Ar] 3d⁹ |
| Zn | [Ar] 3d¹⁰ 4s² | Zn²⁺ | [Ar] 3d¹⁰ |
Memory Trick: Remember Cr and Cu have exceptional configurations: Cr is [Ar]3d⁵4s¹ (not 3d⁴4s²) and Cu is [Ar]3d¹⁰4s¹ (not 3d⁹4s²). This gives them extra stability.
Relative Energies of 4s and 3d Sub-shells
The energy difference between 3d and 4s orbitals changes across the period:
- In calcium, the 4s orbital has lower energy than the 3d orbital
- Along the 3d period, after calcium, the 4s electrons are shielded by the 3d electrons
- The 3d electrons face higher effective nuclear charge due to less shielding
3d Orbital Shapes
The five d orbitals have different orientations in space, which is important for understanding complex formation and d-orbital splitting.
Properties of Transition Elements
Physical Properties
- High melting points due to strong metallic bonding
- High density due to close packing of atoms
- Good conductors of electricity and heat
- Shiny metallic luster
- Malleable and ductile
- Hard and strong
- Show magnetic properties (paramagnetism and ferromagnetism)
Trends in Physical Properties
Density
Density generally increases from scandium to copper due to:
- Increasing atomic mass
- Decreasing atomic radius due to increasing nuclear charge
Atomic Radii
Atomic radii decrease across the period but less steeply than in main group elements due to poor shielding by d electrons.
Interesting Information: The atomic radii of vanadium (23) and zinc (30) is the same because the shielding offered by addition of 3d electrons counter the effect of increasing nuclear charge.
Melting and Boiling Points
Generally high with two minima at manganese and zinc due to stability of electronic configurations.
Magnetic Properties
Paramagnetic Substances
Weakly attracted to magnetic fields due to unpaired electrons (e.g., aluminium, sodium).
Ferromagnetic Substances
Strongly attracted to magnetic fields and retain magnetism (e.g., iron, cobalt, nickel).
Diamagnetic Substances
Weakly repelled by magnets due to paired electrons (e.g., copper).
Tip: To determine if a transition metal complex is paramagnetic, count the number of unpaired electrons. More unpaired electrons = stronger paramagnetism.
Alloys
Transition metals form alloys due to similar atomic sizes. Examples include:
- Steel (iron, chromium, nickel, manganese)
- Brass (copper and zinc)
- Bronze (copper and tin)
Variable Oxidation States
Transition metals can exhibit multiple oxidation states due to the similar energy levels of 4s and 3d orbitals, allowing electrons to be removed from both sub-shells.
Tip: The maximum oxidation state for a transition metal is usually equal to the sum of its 4s and 3d electrons. For example, manganese (3d⁵4s²) can reach +7 oxidation state.
Oxidation States of First Row Transition Metals
Do You Know? Simple ions of transition metals exist in low oxidation states (Mn²⁺) while their complex ions (MnO₄⁻) are stable in high oxidation states.
Stability of Oxidation States
The stability of +2 oxidation state compared to +3 increases with atomic number from left to right. The stability of higher oxidation states decreases across the series.
The greater stability of +2 oxidation state of manganese compared to its +3 oxidation state and +3 oxidation state of iron relative to its +2 oxidation state can be explained by the stability of half-filled 3d sub-shell.
Memory Trick: Remember that Cr³⁺ and Mn²⁺ have half-filled d-orbitals (d⁵), making them particularly stable. Similarly, Cu²⁺ has d⁹ configuration which is also relatively stable.
Factors Affecting Stability of Oxidation States
- Electronic Configuration: Half-filled (d⁵) and fully filled (d¹⁰) configurations are particularly stable
- Nature of Ligands: Strong field ligands stabilize higher oxidation states
- pH of Solution: Acidic conditions favor higher oxidation states
- Complex Formation: Complex ions can stabilize unusual oxidation states
Catalytic Activity
Transition elements or their compounds can be used as catalysts due to:
- Ability to exist in more than one stable oxidation state
- Vacant d orbitals capable of forming temporary dative bonds with ligands
Mechanism of Catalysis
Transition metals can change between different oxidation states (e.g., Fe²⁺/Fe³⁺ or Cu⁺/Cu²⁺). They temporarily change oxidation states during reactions and help convert reactants to products, then regain their original state.
Their incomplete 3d-orbitals can accept electron pairs from ligands, forming temporary dative bonds that hold reactants on the surface and bring them closer, weakening bonds within reactant particles.
Examples of Catalytic Processes
Heterogeneous Catalysis
- Haber process: Iron catalyst for ammonia synthesis
- Contact process: Vanadium pentoxide (V₂O₅) catalyst with reversible change in oxidation state from +5 to +4
Homogeneous Catalysis
Where reactants and catalysts are in same physical state.
Example: Oxidation of iodide ion by peroxodisulphate ion catalyzed by Fe²⁺ ions:
Catalyzed by Fe²⁺ which changes between +2 and +3 oxidation states:
Tip: Remember that transition metals can form intermediate complexes with reactants, lowering the activation energy of reactions.
Formation of Complex Compounds
Ligands
A ligand is an atom or group of atoms which is electron rich and can donate lone pairs of electrons to the transition metal ions, forming dative covalent bonds.
Types of Ligands
- Monodentate: Form one dative covalent bond (e.g., H₂O, NH₃, Cl⁻, CN⁻)
- Bidentate: Form two dative covalent bonds (e.g., 1,2-diaminoethane, ethanedioate ion)
- Polydentate: Form multiple dative covalent bonds (e.g., EDTA – hexadentate)
Ligand Structures
Why Transition Metals Form Complexes
- Small size and high charge density, enabling strong attraction to ligands
- 3d sub-shell has low energy and can bond easily to p-orbitals of ligands
- Vacant d orbitals undergo hybridization before complex formation
Do You Know? The transition metal ions have high charge density, so they are highly polarizing compared to s-block metals. This is why they have higher tendency to form complexes.
Examples of Complex Formation
Copper(II) Complexes
Cu²⁺ ([Ar]3d⁹) forms:
- Octahedral [Cu(H₂O)₆]²⁺ with sp³d² hybridization
- Tetrahedral [CuCl₄]²⁻ with sp³ hybridization
Cobalt(II) Complexes
Co²⁺ ([Ar]3d⁷) forms:
- Octahedral [Co(H₂O)₆]²⁺ with sp³d² hybridization
- Tetrahedral [CoCl₄]²⁻ with sp³ hybridization
Tip: To determine the geometry of a complex, count the coordination number. Coordination number 6 = octahedral, coordination number 4 = tetrahedral or square planar.
Naming Complex Compounds
IUPAC rules for naming complexes:
- Name cation before anion
- Name ligands first in alphabetical order, then the transition metal ion
- Add suffix “ate” to the metal name if the coordination sphere is anionic
- Use prefixes di, tri, tetra for ligands; bis, tris, tetrakis for polydentate ligands
- Anionic ligands use suffix “o” (e.g., chloro, hydroxo)
- Neutral ligands retain their names (e.g., aqua, ammine)
- Oxidation state of central ion is represented by Roman numerals in parentheses
Examples of Complex Compounds
Tetraaminebromochloroplatinum(IV) chloride
Diamminecopper(II) ion
Potassium hexacyanoferrate(II)
Coloured Complexes
Transition metal ions in solution or solid form are coloured because they absorb part of visible light and transmit the remaining part.
Why Are Complexes Coloured?
Light is absorbed by transition metal complexes with incomplete 3d sub-shells. For example:
- Copper(II) ([Ar]3d⁹) – coloured compounds
- Zinc(II) ([Ar]3d¹⁰) – colourless compounds
Color Absorption and Transmission
When a complex absorbs a specific color, we see its complementary color:
| Complex | Color Absorbed | Color Observed |
|---|---|---|
| [Ni(H₂O)₆]²⁺ | Red | Green |
| [Cu(H₂O)₆]²⁺ | Orange | Blue |
| [Ti(H₂O)₆]³⁺ | Yellow-green | Red-violet |
Tip: Remember that colour in transition metal complexes arises from d-d transitions. If there are no electrons in d-orbitals (d⁰) or if d-orbitals are completely filled (d¹⁰), no d-d transitions are possible and the complex is colourless.
d-Orbital Splitting in Complexes
In an isolated transition metal ion, all five d orbitals are degenerate (same energy). When ligands approach, the d orbitals split into sets with different energies.
Octahedral Complexes
In octahedral complexes, d orbitals split into:
- e₉ orbitals: Higher energy (dx²-y² and dz²) – point toward ligands
- t₂₉ orbitals: Lower energy (dxy, dyz, dxz) – point between ligands
d-Orbital Splitting in Octahedral Field
The energy difference (ΔE) between these orbitals corresponds to the energy of photons in the visible region.
Tetrahedral Complexes
In tetrahedral complexes, the splitting is reversed:
- t₂ orbitals: Higher energy
- e orbitals: Lower energy
Tip: The splitting in tetrahedral complexes is smaller (about 4/9) than in octahedral complexes with the same ligands.
Effect of Ligands on Splitting
Different ligands cause different splitting energies (ΔE). Ligands follow the spectrochemical series:
Strong field ligands cause greater splitting than weak field ligands.
Memory Trick: Remember the spectrochemical series with the phrase: “Can Nancy Have Orange Clothes, Brother?” (CN⁻, NH₃, H₂O, OH⁻, Cl⁻, Br⁻)
Ligand Substitution Reactions
Reactions in which one ligand is substituted by another ligand. These reactions occur if the new complex formed is more stable than the original.
Copper(II) Complexes
With Hydroxide Ions
Blue solution → Pale blue precipitate
With Ammonia
Blue solution → Violet-blue solution
With Chloride Ions
Blue solution → Yellow solution (reversible)
Cobalt(II) Complexes
With Hydroxide Ions
Pink solution → Blue precipitate
With Ammonia
Pink solution → Brown solution
With Chloride Ions
Pink solution → Blue solution
Tip: Remember that ligand substitution reactions often result in colour changes, which can be used as simple chemical tests.
Redox Reactions
Transition elements can exist in various oxidation states and undergo redox reactions where their oxidation states change.
Key Reactions
1. Acidified Manganate(VII) with Ethanedioate Ions
E° = +1.03 V (feasible reaction)
2. Manganate(VII) with Iron(II) Ions
E° = +0.75 V (feasible reaction)
3. Copper(II) with Iodide Ions
Forms white precipitate of CuI and red-brown solution of I₂
Tip: To predict if a redox reaction is feasible, calculate the overall E° value. If E° is positive, the reaction is feasible.
Redox Titration
Used to determine the amount of Fe²⁺ ions in a sample by titrating with potassium manganate(VII) solution.
Redox Titration Setup
Stereoisomerism in Coordination Compounds
Stereoisomerism results from different arrangements of atoms/ligands in space.
Geometric Isomerism
Exhibited by complexes with ligands arranged differently in space relative to the central metal ion.
Example: Diamminedichloroplatinum(II)
- cis-isomer: Identical ligands next to each other
- trans-isomer: Identical ligands opposite to each other
Geometric Isomers of [Pt(NH₃)₂Cl₂]
Medical Application: cis-Platin is used as an anticancer drug, while trans-platin has no medicinal properties. Cis-platin forms bridges within DNA strands, disrupting replication in cancerous cells.
Octahedral Complexes
Octahedral complexes can also show geometric isomerism (e.g., [Co(NH₃)₄(H₂O)₂]²⁺).
Note: Tetrahedral complexes cannot show geometric isomerism because all four ligands are adjacent to each other.
Optical Isomerism
Exhibited by molecules that have non-superimposable mirror images (enantiomers).
Optical isomers rotate the plane of polarized light in opposite directions.
Optical Isomers of [Ni(en)₃]²⁺
Polarity of Complexes
cis-isomers are generally polar while trans-isomers are non-polar due to symmetry.
Tip: To determine if a complex can show stereoisomerism, check if it has at least two different types of ligands and the appropriate geometry.
Multiple Choice Questions
Concept Assessment Exercise 6.1
1. Determine the oxidation states of transition metals in the following complexes:
- [Co(NH₃)₆]Cl₃
- Na₂[MnCl₅]
- [Co(en)₂(H₂O)₂]
- [CuCl₄]²⁻
- Na₄[Fe(CN)₂(OH)₄]
2. Name the following complex ions:
- [Co(NH₃)₄Cl₂]Cl
- [Zn(OH)₄]²⁻
- [CrCl₂(NH₃)₄]⁺
Chapter Exercise MCQs
1. Which one of the following species has d¹⁰ sub-shell?
2. Which one is correct statement about zinc element?
- It has complete 3d sub-shell
- It forms colourless complexes
- It is transition metal
- It can exist in one stable oxidation state
3. The electronic configuration of iron cation in the complex, [Fe(CN)₆]³⁻, is
4. Why is the hexaaquacopper(II) ion blue in colour?
5. The cyanide ligand (CN⁻) can form two complexes with iron with formulae, [Fe(CN)₆]³⁻ and [Fe(CN)₆]⁴⁻. What is the oxidation state of iron in them?
6. Transition metals are different from alkali metals and alkaline earth metals in many ways. Which one is incorrect statement about transition metal when compared with group 1 and group 2 metals?
7. The species that can act as ligand in transition metal complex is
8. Transition metals can show different oxidation states. Which one is the most common oxidation state in first row transition elements?
9. Which one of the following properties of transition metals is held responsible for their catalytic behaviour?
10. The coordination number of nickel in the complex [Ni(en)₂(OH₂)₂]²⁺ is
Short Answer Questions
Chapter Exercise Questions
1. The melting point of titanium is higher than calcium in the same period. Justify this statement?
2. The first-row elements of d-block exist in more than one oxidation states. However, zinc shows only +2 oxidation state in its complexes. Why?
3. Define the terms:
- ligand
- coordination number
- complex ion
4. Explain why hydrated Ti⁴⁺ complexes are colourless and hydrated Ti³⁺ complexes are coloured?
5. Why do the melting points of first row transition metals increase upto the middle and then decrease? Comment
6. Why are transition metals used as catalysts in industries for performing different reactions? Give one example.
7. The blue cobalt chloride paper is used to test the presence of water. If water is present, the paper turns pink. This is because six water ligands exchange for four chloride ligands present in cobalt chloride, [CoCl₄]²⁻. Write the equation to show the ligand substitution reaction when the test is positive.
8. Octahedral complexes of copper(II) ion have different colours. Explain why?
Key Points
- Transition metals are those elements that have incomplete d sub-shell.
- The elements of the 3d block have high melting points, boiling points and densities.
- The first and second ionisation energies increase only slightly across the block from scandium to zinc.
- A ligand is a molecule or ion with one or more lone pairs of electrons available to donate to a transition metal ion.
- Transition elements form complexes by combining with ligands.
- Most transition elements form M²⁺ ions by loss of the 4s electrons.
- The elements vanadium, chromium and manganese have a maximum oxidation number equal to the sum of the numbers of 3d and 4s electrons.
- Transition metal ions accept electrons from ligands so metal ions act as Lewis acid while ligands as Lewis bases.
- The colours of complexes are due to d-d electronic transitions.
- The colour of complex observed depend on the colour absorbed.
- The complexes of d block elements which have complete or empty d sub-shell are colourless.
- Transition elements can exist in several oxidation states because of the involvement of both 4s and 3d orbitals.
- Ligand exchange reactions involve exchange of ligands in a complex resulting in change in colours of complexes.
- A strong ligand can displace a weak ligand.
- Ligand exchange can be described in terms of competing equilibria.
- The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion.
- The higher the value of the stability constant, the more stable is the complex ion formed.
- The splitting pattern is different in octahedral and tetrahedral complexes.
- Different ligands will split the d orbitals by different amounts of ΔE, resulting in differently coloured complexes.
Exam Tips:
- Remember that Zn and Sc are not true transition metals as they don’t form ions with incomplete d-orbitals
- For naming complexes, always list ligands in alphabetical order
- For calculating oxidation states in complexes, remember that neutral ligands have zero charge
- Colour in complexes is due to d-d transitions, so d⁰ and d¹⁰ complexes are colourless
- The common oxidation state for first row transition metals is +2
References for Further Information
- Chemistry by Brain Ratcliff, Helen Eccles, John Raffan, John Nicholson, David Johnson and John Newman.
- Chemistry by George Facer
- Chemistry-The molecular nature of matter and change by Silberberg.
- Chemistry by Peter Cann and Peter Hughes.
- Chemistry by Blackman, Bottle, Schmid, Mocerino and Wille.
- Chemistry by Cliff Curtis, Jason Murgatroyd and Divid Scott.
- Chemistry by Christopher Talbot, Richard Harwood and Christopher Coates.