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⚛️ Transition Elements Master Cheatsheet (CIE A-Level 9701)

Overview

This document contains high-yield, mark-scheme-approved definitions, equations, and explanations. Words in bold are standard marking points (AW - alternative wording usually not accepted for these).

1. Core Definitions & Electronic Configurations

Definition: A transition element is a d-block element that forms one or more stable ions with an incomplete (partially filled) d-subshell.

• Exceptions (The Traps):
  • Scandium (Sc): Only forms . Electronic configuration is . (Empty d-subshell).
  • Zinc (Zn): Only forms . Electronic configuration is . (Full d-subshell).
  • Neither are classified as transition elements.
• Filling/Emptying Rule: The 4s subshell fills first and empties first. When writing ion configurations, ALWAYS remove 4s electrons before 3d electrons.
• Anomalies: Chromium () and Copper (). They move one 4s electron to the 3d subshell to achieve energetically stable half-filled or fully-filled 3d subshells.

2. 3d Orbital Shapes (Syllabus Requirement)

You must be able to sketch and describe specific 3d orbitals.

• , , orbitals: Contain four lobes that point between the axes.
• orbital: Contains four lobes that point directly along the x and y axes.
• orbital: Has a dumbbell shape pointing directly along the z-axis, with a “doughnut” ring around the center.

3. Key Chemical Properties

Transition metals exhibit four characteristic properties due to the similar energies of the 3d and 4s subshells:

1. Variable Oxidation States: Due to the small energy difference between 3d and 4s electrons, varying numbers of electrons can be lost.
2. Formation of Complex Ions: Possess energetically accessible vacant d-orbitals that can accept lone pairs from ligands.
3. Colored Compounds: Due to d-d electron transitions between non-degenerate d-orbitals.
4. Catalytic Activity:
   • Homogeneous Catalysis: Transition elements have variable oxidation states, allowing them to transfer electrons, act as reaction intermediates, and easily switch between states (e.g., in the reaction).
   • Heterogeneous Catalysis: Vacant d-orbitals allow reactants to adsorb onto the solid metal surface. This weakens bonds in the reactant molecules, lowering the activation energy, holding them in the correct orientation, and then they desorb (e.g., solid Fe in the Haber process).

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4. Ligands & Complex Ions

Definitions:

• Complex Ion: A central metal atom/ion surrounded by ligands bonded via dative covalent (coordinate) bonds.
• Ligand: A molecule or ion with at least one lone pair of electrons capable of forming a dative covalent bond to a metal ion.
• Coordination Number: The total number of dative covalent bonds formed between the central metal ion and the ligands (NOT necessarily the number of ligands).

Types of Ligands

Type	Definition	Common Examples
Monodentate	Donates 1 lone pair	, , , , ,
Bidentate	Donates 2 lone pairs	en (1,2-diaminoethane: ), ox (ethanedioate: )
Polydentate	Donates >2 lone pairs	(Hexadentate - donates 6 lone pairs)

Geometry & Shapes (Mark Scheme Standard)

Coord. No.	Shape	Bond Angle	Typical Complexes
6	Octahedral		, ,
4	Tetrahedral		Large ligands (high repulsion): ,
4	Square Planar		Platinum complexes: (Cisplatin)
2	Linear		Silver complexes: (Tollens’),

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5. The “Holy Grail” Correction: Deprotonation vs. Ligand Exchange

EXAMINER TRAP: Deprotonation vs Ligand Exchange

Many textbooks incorrectly label the addition of aqueous or dilute aqueous to a hexaaqua ion as “ligand exchange”. It is NOT. It is an acid-base reaction (deprotonation / precipitation) where the or acts as a base, removing protons from the water ligands to form a neutral precipitate.

Ligand exchange ONLY occurs when bonds between the metal and the ligand are broken and replaced (e.g., adding excess concentrated or ).

A. Acidity of Hexaaqua Ions (Hydrolysis)

• Aqueous ions (like ) are highly charged and polarize the bonds in the water ligands.
• They act as Brønsted-Lowry acids, releasing into solution (creating acidic solutions):

B. Copper(II) Chemistry

• Aqueous Ion: (Pale blue solution)
• Add dilute or limited (DEPROTONATION):
  • Equation:
  • Observation: Pale blue precipitate.
• Add EXCESS conc. (LIGAND EXCHANGE):
  • Equation:
  • Observation: Precipitate dissolves to form a deep blue solution. (Coordination number remains 6, shape is octahedral).
• Add conc. (LIGAND EXCHANGE):
  • Equation:
  • Observation: Pale blue solution turns yellow (appears green in transition due to blue + yellow mixing).
  • Note: Coordination number changes from 6 to 4 because is a larger ligand and they repel each other more strongly. Geometry changes from Octahedral to Tetrahedral.

C. Cobalt(II) Chemistry

• Aqueous Ion: (Pink solution)
• Add dilute or limited (DEPROTONATION):
  • Observation: Blue precipitate (which turns red/pink on warming or standing). Formula: .
• Add EXCESS conc. (LIGAND EXCHANGE):
  • Observation: Precipitate dissolves to form a brown solution ().
  • Redox note: The complex is unstable and oxidizes in air to form the complex (yellow/orange).
• Add conc. (LIGAND EXCHANGE):
  • Observation: Pink solution turns into a blue solution (). Tetrahedral shape.

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6. Origin of Colour in Complexes (5-Mark Standard Answer)

Orbital Splitting Patterns (Frequently Tested)

Octahedral complexes: The 6 ligands approach along the x, y, and z axes. Therefore, electrons in the and orbitals experience greater repulsion, splitting the 3d subshell so 2 orbitals are at a higher energy and 3 orbitals are at a lower energy. Tetrahedral complexes: The 4 ligands approach between the axes. The orbitals experience greater repulsion, splitting the 3d subshell so 3 orbitals are at a higher energy and 2 orbitals are at a lower energy.

If asked “Explain why transition element complexes are colored” write exactly this sequence:

1. In an isolated atom, the five 3d orbitals are degenerate (equal in energy).
2. In a complex, ligands donate lone pairs. The repulsion from these lone pairs causes the 3d orbitals to split into two sets of non-degenerate orbitals with an energy gap, .
3. Electrons in the lower energy d-orbitals absorb energy from the visible light spectrum.
4. This exact energy promotes an electron to a higher energy d-orbital (d-d transition). Equation to cite: (where is Planck’s constant, is frequency).
5. The color observed is the complementary color (the frequencies of visible light that are not absorbed / are transmitted).

Factors changing the color (changing ):

1. Oxidation state of the metal. 2. Identity of the metal. 3. Identity of the ligand. 4. Coordination number / Geometry.

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7. Stability Constants ()

Definition: The equilibrium constant for the formation of a complex ion in a solvent from its constituent ions or molecules.

• Expression Rule: Written exactly like . is omitted from the expression because it is the solvent and its concentration is effectively constant/in excess.
• Interpretation: A higher value (or more positive ) means the equilibrium lies further to the right the complex is more stable.

The Chelate Effect (Entropy)

Replacing monodentate ligands with bidentate/polydentate ligands causes a massive increase in stability. Example:

• Explanation: 4 reacting particles generate 7 product particles. There is a massive increase in the number of particles in solution, leading to a large increase in entropy (). Therefore, the reaction is highly feasible, making polydentate complexes incredibly stable.

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8. Stereoisomerism in Complexes

A. Geometrical (cis-trans) Isomerism

• Square Planar: .
  • Cisplatin: Identical ligands are adjacent (). Used as an anti-cancer drug (binds to DNA, preventing replication). Polar.
  • Transplatin: Identical ligands are opposite (). Inactive in cancer treatment. Non-polar (dipoles cancel).
• Octahedral: Requires complexes of type (e.g., ) or .
  • Cis: The two identical ligands are apart.
  • Trans: The two identical ligands are apart (Non-polar due to symmetry).

B. Optical Isomerism

• Occurs primarily in octahedral complexes containing bidentate ligands.
• Criteria: Must have non-superimposable mirror images (enantiomers) and lack a plane of symmetry.
• Classic Example: or .
• Property: Rotate the plane of plane-polarized light in opposite directions.

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9. Redox Reactions & Feasibility

We use standard electrode potentials () to predict feasibility.

• Rule: The species with the more positive will be reduced (gains electrons / acts as the oxidizing agent).
• Reaction is feasible if . (Note: Very large activation energies can stop feasible reactions from occurring in reality).

System 1: Permanganate () and Iron(II) ()

• Acidic Conditions required (usually ).
• Half equations:
• Overall:
• Observation: Titrate until a permanent pale pink color appears (self-indicating).

System 2: Permanganate () and Oxalate () - AUTOCATALYSIS

• Overall:
• Examiner Gold: This reaction exhibits autocatalysis. The reaction is initially slow (high activation energy due to repulsion between two negative ions, and ). As is produced, it acts as a catalyst, and the reaction rate increases exponentially.
• Mechanism: changes oxidation state between and to transfer electrons between the negative ions.

System 3: Copper(II) and Iodide (Iodometric Titration)

• Used to find the concentration of in solution.
• Step 1: Add excess to .
  • Observation: A white precipitate () forms in a brown solution (iodine).
• Step 2: Titrate the liberated iodine against standard Sodium Thiosulfate ().
  • Indicator: Add starch near the end-point (when solution fades to pale yellow). Solution turns blue-black, titrate until colorless.
• Ratio: 1 mole of produces 0.5 moles of , which reacts with 1 mole of . Therefore, .

System 4: Iron(III) and Iodide (Common Exam Reaction)

• Overall Equation:
• Observation: The yellow/orange solution of turns brown due to the formation of aqueous . Can also be titrated against Thiosulfate.