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Transition Metals & Complexes

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Edexcel IAL Chemistry · Inorganic

Transition Metals & Complexes

The Big Idea: Transition metal ions are magnets for electron pairs — molecules and ions called ligands dock onto them to form complexes. This bonding splits the metal's d-orbitals just enough that visible light gets absorbed, which is why transition metal compounds are colourful, why they can swap one ligand for another, and why some (like cisplatin) can be turned into life-saving drugs.

d-block chemistry Ligands & complexes Colour theory Isomerism Medicinal chemistry

Summary — The Whole Chapter at a Glance

A transition metal is an element that forms at least one stable ion with a partially filled d-subshell. This is why Sc and Zn are excluded, even though they sit in the d-block.
Transition metals lose their 4s electrons before 3d electrons when forming ions — always. Cr and Cu have "anomalous" ground-state configurations ([Ar]3d⁵4s¹ and [Ar]3d¹⁰4s¹) because a half-full or full d-subshell is extra stable.
Because the d-subshell can lose different numbers of electrons, transition metals show variable oxidation states (shown with Roman numerals).
A complex = central metal ion + ligands attached by dative covalent (co-ordinate) bonds, where the ligand donates both electrons.
Ligands are classed by how many dative bonds they can form: monodentate (1), bidentate (2), multidentate (many, e.g. EDTA⁴⁻ = 6).
The coordination number = total number of dative bonds to the central ion. This decides the shape: 6 → octahedral (90°), 4 → tetrahedral (109.5°) or square planar (90°).
Colour arises because ligands split the five degenerate 3d orbitals into two energy levels. An electron absorbs visible light to jump the gap (d–d transition); the colour you see is the complementary colour of what's absorbed.
Colour changes when you change the ligand, oxidation state, or coordination number — all four factors alter the size of the splitting energy, ΔE.
Ligand exchange (substitution) is reversible and swaps one ligand type for another, often changing colour, and sometimes changing coordination number/shape if the ligand sizes differ a lot.
Bidentate/multidentate ligands replacing monodentate ones is the chelate effect — favourable because it's driven by a big positive entropy change (more product particles than reactant particles).
Transition metal complexes can show cis-trans (geometric) isomerism and optical isomerism, even without a C=C double bond.
Cisplatin is a square planar complex used in cancer treatment; it works by ligand-exchanging its Cl⁻ ligands for water, then binding to DNA bases and blocking replication.
Haemoglobin uses an Fe²⁺ centre in a square-planar-like haem group to transport O₂; CO poisoning happens because CO binds more strongly than O₂ to the same site.

1. Transition Metals — Electron Configuration & Oxidation States

What actually makes something a "transition metal"?

Here's the precise definition, and it matters because exam questions love to test the exception cases: a transition metal is an element that forms at least one stable ion with an incomplete (partially filled) d-subshell.

Notice this definition is about the ion, not the atom. That's the trap. Look at the two elements that sit in the d-block but aren't classed as transition metals:

Scandium (Sc): only ever forms Sc³⁺, configuration [Ar] 3d⁰ — completely empty d-subshell. No partially-filled d ever exists.
Zinc (Zn): only ever forms Zn²⁺, configuration [Ar] 3d¹⁰ — completely full d-subshell. Also never partially filled.

So even though Sc and Zn are physically sitting in the d-block of the periodic table, they fail the "partially-filled d-subshell in a stable ion" test. That leaves titanium through to copper as the true first-row transition elements.

Analogy

Think of the d-subshell like a hotel with 10 rooms (5 orbitals × 2 electrons each). Sc's ion checks out with the hotel completely empty (0/10 rooms full), and Zn's ion checks out with the hotel completely full (10/10). Every element from Ti to Cu, though, always has some but not all rooms occupied in at least one of its stable ions — a partially-booked hotel. That partial occupancy is exactly what creates colour and variable oxidation states later on, so it's the whole reason the definition is written this way.

Filling order: 4s before 3d (but leaving 4s first!)

By the Aufbau principle, electrons fill the lowest-energy subshells first. For a neutral first-row transition metal atom, the 4s subshell is slightly lower in energy than 3d, so 4s fills before 3d.

Element	Full electron configuration
Ti	1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²
V	1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s²
Cr	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ (anomaly!)
Mn	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s²
Fe	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
Co	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s²
Ni	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁸ 4s²
Cu	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹ (anomaly!)

The Two Exceptions You MUST Know

Cr = [Ar] 3d⁵ 4s¹ (not [Ar] 3d⁴ 4s²)

Cu = [Ar] 3d¹⁰ 4s¹ (not [Ar] 3d⁹ 4s²)

In plain words: promoting one electron from 4s into 3d gives chromium a half-full d-subshell (5 electrons, one in each orbital) and gives copper a completely full d-subshell (10 electrons). Both of these arrangements are lower in energy — more stable — than the "expected" configuration, so nature prefers them.

Forming ions: 4s empties FIRST

This is the single most commonly tested idea in this section, and it trips students up because it seems backwards. When an atom fills its subshells, 4s goes in before 3d. But when the atom then loses electrons to become an ion, the 4s electrons leave first, not the 3d ones.

Why the flip? Once the 3d subshell starts filling up, the extra electron density in 3d repels the 4s electrons and pushes their energy above that of 3d. So in an ion (or actually in the filled atom too, once you look closely), 4s becomes the outermost, highest-energy, "first-to-leave" shell — even though it was the first to fill.

Worked Example

State the full electron configuration of the manganese(III) ion, Mn³⁺.

Practice Question

Write the full electron configuration of Fe²⁺ and explain why Sc does not count as a transition metal.

Variable Oxidation States

Because the d and s electrons are so close in energy, transition metals can lose different numbers of electrons depending on the reaction conditions — giving multiple stable oxidation states for the same element. This is why we need Roman numerals in complex ion names: Fe²⁺ = iron(II), Fe³⁺ = iron(III).

Element	Common oxidation states
Ti	+2, +3, +4
V	+2, +3, +4, +5
Cr	+2, +3, +6
Mn	+2, +4, +6, +7
Fe	+2, +3
Co	+2, +3
Ni	+2, +3, +4
Cu	+1, +2

Quick Note

+2 is the single most common oxidation state across the whole row — that's a direct consequence of the 4s electrons always going first. Losing just the 4s² gives +2 with minimal energy cost.

2. Transition Metal Complexes — Ligands & Naming

What is a complex, really?

A complex is a central metal ion surrounded by ligands attached via dative covalent (co-ordinate) bonds — bonds where both shared electrons come from the same atom (the ligand), not one from each.

A ligand is any molecule or ion that can donate a lone pair of electrons into an empty space on the metal ion. This should sound familiar from organic chemistry — a ligand is functionally the same idea as a nucleophile (an electron-pair donor, a Lewis base). The only difference is the "electrophile" here is a metal ion instead of an electrophilic carbon.

Key Definition

Ligand = a species with a lone pair (or negative charge with a lone pair) that forms a dative bond to a metal ion, donating BOTH electrons.

In plain words: the ligand does all the "giving" — the metal ion just has an empty orbital ready to receive.

Monodentate, bidentate, and multidentate — "how many hands does the ligand have?"

Some ligands can only form one dative bond because they only have one accessible lone pair. Others have two or more separated donor atoms, so they can grip the metal ion in multiple places at once — like holding a ball with one finger versus holding it with a whole hand.

Monodentate (1 bond)

H₂O (water)
NH₃ (ammonia)
Cl⁻ (chloride)
CN⁻ (cyanide)
OH⁻ (hydroxide)

Bidentate (2 bonds)

1,2-diaminoethane H₂NCH₂CH₂NH₂ (aka "en") — two N atoms each with a lone pair
Ethanedioate ion C₂O₄²⁻ (aka "ox") — two O atoms donate

Multidentate Star Example

EDTA⁴⁻ is a hexadentate ligand — it has 6 separate lone-pair-donating atoms (4 oxygens from carboxylate groups + 2 nitrogens) and wraps itself all the way around a single metal ion, forming all 6 dative bonds by itself. That's why EDTA is used to "mop up" heavy metal ions — one EDTA molecule locks up one whole metal ion.

Coordination number

The coordination number is the total number of dative bonds formed to the central metal ion — not the number of ligands. This distinction matters a lot: six monodentate ligands give coordination number 6, but so does one hexadentate ligand like EDTA⁴⁻, and so does three bidentate ligands (3 × 2 = 6).

Rule

Coordination number = (number of monodentate ligands × 1) + (number of bidentate ligands × 2) + (number of hexadentate ligands × 6)

Always count bonds, not ligand molecules.

Naming complexes

The naming pattern is: prefix (number of ligands) + ligand name + metal name/oxidation number. If the overall complex ion is negatively charged (an anion), the metal name changes to end in "-ate".

Complex	Name
[CuCl₄]²⁻	tetrachlorocuprate(II)
[Fe(H₂O)₆]²⁺	hexaaquairon(II)
[Co(NH₃)₆]²⁺	hexaamminecobalt(II)
[Ni(CN)₄]²⁻	tetracyanonickelate(II)

Spelling Trap

Ammonia as a ligand takes the prefix "ammine" (double-m) — not "amine" (the organic functional group, single-m). Examiners specifically check for this. "Ammine" = ligand; "amine" = —NH₂ functional group.

Practice Question

What is the coordination number of [Co(C₂O₄)₃]³⁻, and what shape would you predict?

3. Colour in Aqueous Ions

Why are transition metal compounds coloured at all?

This is genuinely one of the most elegant ideas in the whole A-Level course, so let's build it up slowly.

In a free transition metal ion (no ligands attached), all five 3d orbitals have exactly the same energy — chemists say they're degenerate. But the moment ligands bond to the metal ion, their approaching electron pairs repel the electrons sitting in the d orbitals unevenly (some d orbitals point more directly at the incoming ligands than others). This uneven repulsion splits the five d orbitals into two slightly different energy levels — three lower-energy orbitals and two higher-energy ones (for an octahedral complex).

┌───┐ ┌───┐ │dx²-y²│ dz² │ ← higher energy set (2 orbitals) └───┘ └───┘ 3d (all same energy) ──► ΔE (energy gap = "the split") ┌───┐┌───┐┌───┐ │dyz││dxz││dxy│ ← lower energy set (3 orbitals) └───┘└───┘└───┘

An electron sitting in the lower set can absorb a photon of visible light and jump ("get promoted") up into the higher set. This is called d–d transition / promotion / excitation. Crucially, only certain frequencies of light have exactly the right energy to cause this jump — so only those frequencies get absorbed. The rest of the visible spectrum passes straight through (or reflects), and that's the mixture of colours your eye actually perceives.

The Golden Rule of Complementary Colour

Colour you SEE = complementary (opposite) colour of the light that was ABSORBED

A solution that absorbs red light looks green (its complementary colour on the colour wheel), and vice versa. You never see the colour of the light it absorbed — you see everything else that got through.

Why Zn²⁺ and Sc³⁺ solutions are colourless Zn²⁺ has a completely full 3d¹⁰ subshell — there's no empty higher-energy orbital for an electron to jump into, so no d–d transition can happen, so no visible light is absorbed. Sc³⁺ has a completely empty 3d⁰ subshell — there's no electron there to be promoted in the first place. Either way, no absorption = no colour = colourless solution. This is the deep reason those two elements are excluded from the "transition metal" definition — colour and variable oxidation states are two sides of the same "partially-filled d-subshell" coin.

Four factors that change the size of the split (ΔE) — and therefore the colour

Since colour depends entirely on the size of the energy gap ΔE, anything that changes ΔE changes the colour. There are four such factors:

Factor	Effect	Example
1. Ligand type	Ligands with higher charge density interact more strongly with the metal ion, causing a bigger split	[Cu(H₂O)₆]²⁺ is light blue; [Cu(NH₃)₄(H₂O)₂]²⁺ is dark blue — same metal, same oxidation state, different ligands
2. Oxidation state	Higher oxidation state = stronger pull on ligand electrons = bigger split	[Fe(H₂O)₆]²⁺ is green (absorbs red); [Fe(H₂O)₆]³⁺ is orange (absorbs blue)
3. Coordination number / shape	Changing from 6-coordinate to 4-coordinate changes the geometry of orbital-ligand interaction	[Cu(H₂O)₆]²⁺ light blue → [CuCl₄]²⁻ yellow (6-coordinate octahedral → 4-coordinate)
4. Identity/nuclear charge of the metal	Different metals have different nuclear charges pulling on the d electrons, changing the baseline split	Compare the colours across the whole Ti→Cu row — all different even with the same ligand

Practice Question

A student adds concentrated HCl to a pink solution of [Co(H₂O)₆]²⁺. The solution turns blue. Explain this colour change in terms of coordination number and ΔE.

4. Ligand Exchange (Substitution)

What is it?

Ligand exchange is when one ligand attached to the central metal ion is swapped for another — partially or entirely. It happens because the new complex formed is more stable (thermodynamically favoured) than the old one. These reactions are reversible (equilibria), shown with a ⇌ arrow.

Classic Example — Copper

[Cu(H₂O)₆]²⁺(aq) + 4Cl⁻(aq) ⇌ [CuCl₄]²⁻(aq) + 6H₂O(l)

Add concentrated HCl slowly to blue copper(II) sulfate solution: blue → green (mixture) → yellow. All six water ligands are replaced by four chloride ligands, and coordination number drops from 6 to 4. Note the overall charge flips from +2 to −2, but the copper's oxidation state stays +2 throughout — ligand exchange does NOT change oxidation state on its own.

Common Mistake

Students often think dissolving solid CuCl₂ in water gives the yellow [CuCl₄]²⁻ complex. It doesn't! Water is in huge excess, so it outcompetes chloride and forms the blue [Cu(H₂O)₆]²⁺ complex instead. You only get [CuCl₄]²⁻ when chloride ions are added in high concentration (e.g. concentrated HCl or saturated NaCl).

Key points to remember about ligand exchange:

If the ligands being swapped are a similar size (e.g. H₂O ↔ NH₃), coordination number and shape stay the same.
If the ligands are a very different size (e.g. H₂O ↔ Cl⁻), coordination number and shape can change (Cl⁻ is bigger, so fewer fit around the metal).
Ligand exchange reactions can change the charge of the complex ion, but never on their own change the metal's oxidation state.

The Chelate Effect — why bidentate/multidentate ligands "win"

When a bidentate or multidentate ligand (like "en" or EDTA⁴⁻) replaces several monodentate ligands, the reaction is called the chelate effect, and it's always energetically favourable — meaning ΔG is negative.

The interesting part is why. Using the Gibbs equation:

Gibbs Free Energy Equation

ΔG = ΔH_reaction − TΔS_system

For chelation reactions, ΔH is usually small (similar bonds are being made and broken — dative M–O and M–N bonds either way). But ΔS is strongly POSITIVE, because one multidentate ligand kicks out several small monodentate ligands, meaning the total number of free particles increases. That big positive ΔS dominates, driving ΔG negative — the reaction happens because of entropy, not enthalpy.

Example

[Co(H₂O)₆]²⁺(aq) + EDTA⁴⁻(aq) → [CoEDTA]²⁻(aq) + 6H₂O(l)

Count the particles: 2 reactant species become 7 product species (1 complex + 6 free water molecules). That's a huge increase in disorder/entropy — this is the real driving force of the chelate effect.

Practice Question

Explain, using ideas about entropy, why the reaction below is spontaneous even though the enthalpy change is small:
[Ni(H₂O)₆]²⁺(aq) + 3en(aq) → [Ni(en)₃]²⁺(aq) + 6H₂O(l)

5. Shapes of Complexes — Octahedral, Tetrahedral & Square Planar

Same rule as VSEPR, just relabelled

Good news: you already know how to predict these shapes from AS Chemistry's electron-pair repulsion theory. The only twist is that you ignore the metal's own 3d electrons and the overall charge on the complex, and just count how many electron pairs the ligands are donating.

Coordination number	Shape	Bond angle	Typical cause
6	Octahedral	90°	6 small monodentate ligands, OR 3 bidentate, OR 1 hexadentate (EDTA⁴⁻)
4	Tetrahedral	109.5°	4 large ligands (e.g. Cl⁻) — too big to fit 6 around the metal
4	Square Planar	90°	Less common; typical for CN⁻ ligands and famously for platinum(II) complexes like cisplatin

Why does ligand size decide tetrahedral vs octahedral?

Small ligands like H₂O and NH₃ can pack six around a central metal ion without bumping into each other, giving octahedral geometry. Large ligands like Cl⁻ physically don't have room for six — only four fit — giving tetrahedral geometry instead. Size, not charge, is usually the deciding factor here.

Isomerism in complexes

Even though transition metal complexes don't contain a C=C double bond, they can still show geometric (cis-trans) isomerism — this happens in square planar and octahedral complexes that have two pairs of different ligands.

Cis isomer: the two "different" ligands are next to each other (adjacent).
Trans isomer: the two "different" ligands are opposite each other.

Octahedral complexes with bidentate ligands (like [Ni(en)₃]²⁺) can also show optical isomerism — the complex and its mirror image are non-superimposable, just like a left and right hand. These two mirror-image forms (called enantiomers) rotate plane-polarised light in opposite directions but are otherwise chemically identical.

Practice Question

[Co(NH₃)₄(H₂O)₂]²⁺ is octahedral. Sketch (in words) the difference between its cis and trans isomers, and state whether either shows optical isomerism.

6. Square Planar Complexes & Cisplatin

Cisplatin — chemistry that treats cancer

Cisplatin, [Pt(NH₃)₂Cl₂], is a square planar complex discovered in the 1960s that became a landmark cancer drug — effective against testicular, ovarian, cervical, breast, lung, and brain cancers. The mechanism is a beautiful real-world application of everything you've just learned about ligand exchange.

How Cisplatin Actually Works — Step by Step

1. Cisplatin passes through the cancer cell membrane.
2. Inside the cell, chloride ion concentration is much lower than outside — so ligand exchange occurs: the two Cl⁻ ligands are gradually replaced by water molecules.
3. Nitrogen atoms on the DNA bases (adenine/guanine) are better ligands than water, so a second ligand exchange happens: the water ligands are displaced and the platinum forms dative covalent bonds directly to the DNA bases.
4. This distorts the shape of the DNA double helix, which blocks the cell's ability to replicate its DNA — and a cell that can't replicate its DNA can't divide, so the cancer cell dies.

Crucially, only the cis isomer works this way. The trans isomer, transplatin, has the same formula but its geometry means it can't bind to DNA in the same double-anchoring way, so it has no useful anticancer activity — a striking example of how geometric isomerism alone can be the difference between a life-saving drug and an inactive compound.

The Trade-off

Cisplatin can't tell the difference between fast-dividing cancer cells and other fast-dividing healthy cells (like hair follicles), which is exactly why hair loss is a common side effect of chemotherapy. This is a genuine ethical/practical talking point examiners like to ask about: society accepts cisplatin's side effects because its therapeutic benefit (killing cancer cells) outweighs the harm, and newer drug delivery research aims to target cancer cells more selectively.

Practice Question

Explain why transplatin (the trans isomer) is not effective as an anticancer drug, while cisplatin is.

7. Haemoglobin & Ligand Exchange

Nature's own transition metal complex

Haemoglobin is built around an Fe²⁺ ion sitting at the centre of a haem group — a multidentate ligand made mostly of carbon and hydrogen, where four nitrogen atoms form dative covalent bonds to the iron in a roughly square planar arrangement. A fifth dative bond comes from the protein chain (globin) itself, and the sixth open coordination site is where oxygen binds.

Why oxygen transport works at all

Oxygen is actually a fairly weak ligand — it only bonds loosely to the Fe²⁺ ion. That weakness is a feature, not a bug: it means the O₂ can bind in the lungs (where oxygen concentration is high) and then break off easily once haemoglobin reaches a tissue cell that needs it (where oxygen concentration is lower). A strong bond would trap the oxygen and haemoglobin could never "let go" of it.

Why carbon monoxide is so dangerous

Carbon monoxide (CO) competes for exactly the same sixth binding site as oxygen — but CO forms a much stronger dative covalent bond to Fe²⁺ than O₂ does. This means CO effectively "wins" the ligand exchange competition and displaces oxygen from haemoglobin, forming carboxyhaemoglobin (which has a distinctive darker red colour, a diagnostic sign of CO poisoning). Since the CO bond is so strong, it's very slow to be displaced back off again, so the haemoglobin becomes unavailable for oxygen transport — the body effectively suffocates even while breathing normal air.

Practice Question

Using the idea of ligand exchange, explain why carbon monoxide poisoning is dangerous even in low concentrations.

What to Memorise

Definitions

Transition metal: forms ≥1 stable ion with a partially filled d-subshell
Ligand: donates a lone pair to form a dative bond to a metal ion
Complex: central metal ion + attached ligands
Coordination number: total number of dative bonds to the central ion
Chelate effect: multidentate ligands replacing monodentate ones, driven by +ΔS

Shapes & Angles

6-coordinate → Octahedral, 90°
4-coordinate (small area) → Tetrahedral, 109.5° (large ligands, e.g. Cl⁻)
4-coordinate (flat) → Square planar, 90° (e.g. CN⁻, Pt(II) complexes)

Cr & Cu Exceptions

Cr = [Ar] 3d⁵ 4s¹
Cu = [Ar] 3d¹⁰ 4s¹
Reason: half-full/full d-subshell = extra stable
Ions ALWAYS lose 4s electrons before 3d

4 Factors Affecting Colour

1. Type/size of ligand
2. Oxidation state of metal
3. Coordination number/shape
4. Identity of the metal ion

Key Equations

[Cu(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CuCl₄]²⁻ + 6H₂O (blue→yellow, 6→4 coord.)
[Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O (pink→blue)
ΔG = ΔH − TΔS (chelate effect, driven by +ΔS)

Applications

Cisplatin (square planar Pt) → binds DNA, blocks replication → cancer treatment
Haemoglobin (Fe²⁺, haem) → transports O₂; CO binds more strongly → poisoning

Concepts Checklist

I can explain why Sc and Zn are excluded from the "transition metal" definition
I can write full electron configurations for transition metal atoms and ions, including the Cr and Cu exceptions
I know that 4s electrons are always removed before 3d electrons when forming ions
I can explain why transition metals show variable oxidation states
I can define "ligand" and "complex" precisely, using the language of dative covalent bonding
I can classify ligands as monodentate, bidentate, or multidentate and count coordination number correctly
I can name complex ions correctly, including the "-ate" ending for anionic complexes
I can explain colour in transition metal complexes using d-orbital splitting and complementary colours
I can list and explain all four factors that change the splitting energy ΔE (and hence colour)
I can write and balance ligand exchange equilibrium equations
I can explain the chelate effect using entropy and the Gibbs free energy equation
I can predict complex shapes (octahedral, tetrahedral, square planar) from coordination number
I can identify cis/trans and optical isomers in transition metal complexes
I can explain how cisplatin works and why the trans isomer doesn't
I can explain haemoglobin's structure and why carbon monoxide poisoning occurs

Exam Tips — Traps & What Examiners Actually Want

Trap 1: Electron configuration order

When writing an ion's configuration, DO NOT just remove electrons from the end of the string you wrote for the atom. Always remove from 4s first, regardless of where 4s sits in your written order. Writing Fe²⁺ as [Ar]3d⁴4s² (removing from 3d) instead of [Ar]3d⁶ is one of the most common lost marks in this topic.

Trap 2: Coordination number ≠ number of ligands

If a question gives you 3 bidentate ligands, do not write "coordination number = 3." Count bonds: 3 × 2 = coordination number 6. Examiners specifically design questions with bidentate/multidentate ligands to catch students who forget this.

Trap 3: Ligand exchange ≠ oxidation state change

Swapping [Cu(H₂O)₆]²⁺ for [CuCl₄]²⁻ changes the overall ion charge from +2 to −2, but copper is still Cu²⁺ throughout — its oxidation state hasn't changed. Only redox reactions (like Fe²⁺ → Fe³⁺ + e⁻) change oxidation state. Don't confuse "charge on the complex" with "oxidation state of the metal."

Trap 4: "Explain the colour" answers need the full chain

A full-mark colour explanation must include: (1) ligands cause the d-orbitals to split into two energy levels, (2) an electron is promoted/excited from the lower to the higher level by absorbing a specific frequency of visible light, (3) the colour observed is complementary to the light absorbed. Missing any one of these three links loses marks — don't just say "it absorbs light and looks a colour."

Trap 5: The chelate effect is about ENTROPY, not enthalpy

If asked why a chelation reaction is thermodynamically favourable, the enthalpy change is usually small (similar bonds broken/formed). The answer examiners want is about the increase in the number of free particles (more ligand molecules released than consumed), giving a large positive ΔS_system, which makes ΔG negative via ΔG = ΔH − TΔS.

What examiners are actually checking for

Precise use of key terms: "dative covalent bond," "partially filled d-subshell," "degenerate orbitals"
Correctly balanced ligand exchange equilibrium equations with state symbols
Clear cause-and-effect chains (not just isolated facts) — especially for colour and cisplatin questions
Correct Roman numeral oxidation states in complex naming

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Also in the full note

1. Transition Metals — Electron Configuration & Oxidation States
2. Transition Metal Complexes — Ligands & Naming
5. Shapes of Complexes — Octahedral, Tetrahedral & Square Planar
6. Square Planar Complexes & Cisplatin
7. Haemoglobin & Ligand Exchange
Exam Tips — Traps & What Examiners Actually Want
Shapes & Angles
Cr & Cu Exceptions

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