The d and f Block Elements

Positions of d and f-Block Elements in the Periodic Table

**Definition and Location:**

The periodic table contains two major blocks of transition elements:

**d-block (Transition Metals)**: Groups 3–12, consisting of elements where d orbitals are progressively filled across four long periods (3d, 4d, 5d, 6d)

**f-block (Inner Transition Metals)**: Two series where f orbitals are progressively filled (4f series: Ce to Lu called lanthanoids; 5f series: Th to Lr called actinoids)

**IUPAC Definition of Transition Elements:**

According to IUPAC, transition metals are defined as **metals which have incomplete d subshells either in their neutral atoms or in their common ions**. This distinguishes them from s and p-block elements.

**Important Exception — Zinc, Cadmium, Mercury, and Copernicium (Zn, Cd, Hg, Cn):**

Electronic configuration: **(n-1)d¹⁰ns²** with **completely filled d orbitals** in both ground state and common oxidation states

**NOT classified as transition elements** because their d orbitals are fully occupied

However, their chemistry is studied alongside transition metals as they are end members of 3d, 4d, and 5d series respectively

**Example:** Silver (Z = 47) has configuration [Ar]3d¹⁰4s¹ in ground state (completely filled d orbitals), but it is still classified as a transition element because in its common oxidation states (Ag⁺: 3d¹⁰), the d orbital remains incompletely occupied in the ion before ionization occurs, or because historically it shows variable valency and complex formation characteristic of transition metals.

Electronic Configurations of d-Block Elements

**General Electronic Configuration:**

The outer electron configuration of d-block elements follows the pattern: **(n-1)d¹⁻¹⁰ ns¹⁻²**

(n-1) represents penultimate d orbitals (inner d electrons)

ns represents outermost s orbital

Electrons can range from 1–10 in d orbitals and 1–2 in s orbitals

**Exceptions to the General Configuration:**

Due to **small energy difference between (n-1)d and ns orbitals** and the relative stability of **half-filled (d⁵) and completely filled (d¹⁰) subshells**, several elements show anomalous configurations:

**Chromium (Cr, Z = 24):**

Expected: 3d⁴4s²

Actual: **3d⁵4s¹**

Reason: d⁵ (half-filled) and s¹ (singly occupied) configurations are more stable

**Copper (Cu, Z = 29):**

Expected: 3d⁹4s²

Actual: **3d¹⁰4s¹**

Reason: d¹⁰ (completely filled) configuration provides extra stability

**Palladium (Pd, Z = 46):**

Expected: 4d⁸5s²

Actual: **4d¹⁰5s⁰**

Reason: d¹⁰ completely filled configuration is highly stable, so even the s orbital is left unfilled

**First Series (3d) Electronic Configurations:**

| Element | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn |

|---|---|---|---|---|---|---|---|---|---|---|

| Configuration | 3d¹4s² | 3d²4s² | 3d³4s² | 3d⁵4s¹ | 3d⁵4s² | 3d⁶4s² | 3d⁷4s² | 3d⁸4s² | 3d¹⁰4s¹ | 3d¹⁰4s² |

**Similar configurations in second (4d) and third (5d) series with corresponding anomalies at d⁵ and d¹⁰ positions.**

General Characteristics of Transition Elements (3d Series Focus)

Physical Properties

**Metallic Character:**

Nearly all transition elements display **typical metallic properties**:

High tensile strength, ductility, malleability

High thermal and electrical conductivity

Metallic lustre and characteristic metallic bonding

**Exceptions:** Zn, Cd, Hg, and Mn have atypical properties; Zn, Cd, Hg lack typical metallic structures at normal temperatures.

**Melting and Boiling Points:**

**Very high** compared to s and p-block elements (except Zn, Cd, Hg)

**Trend in 3d series:** Increase from Sc to Mn (maximum at d⁵), then decrease towards Zn

**Cause:** Large number of unpaired electrons in d orbitals participate in **interatomic metallic bonding** in addition to ns electrons

**Anomalies:** Mn and Tc show lower-than-expected melting points despite d⁵ configuration

**Enthalpy of Atomisation (ΔₐH°):**

**Definition:** Energy required to convert one mole of solid metal into gaseous atoms

**Trend:** Shows maximum at **d⁵ configuration** (Mn in 3d series) with one unpaired electron per d orbital being most favorable for strong interatomic interaction

**Significance:** Indicates strength of metallic bonding; metals with very high ΔₐH° tend to be **noble in reactions** (less easily oxidized)

**Series comparison:** **4d and 5d series > 3d series** in ΔₐH° values, explaining greater frequency of metal–metal bonding in heavier transition metal compounds

Variation in Atomic and Ionic Radii

**Trend Within a Series (3d):**

**Atomic and ionic radii decrease** progressively from Sc to Zn with increasing atomic number

**Reason:** Each new electron enters d orbital while nuclear charge increases (+1), but d electrons provide weak shielding (d electrons are not effective at shielding each other due to diffuse orbital shape)

**Net effect:** Progressive increase in effective nuclear charge causes radius contraction

**Magnitude:** Variation is **small** within a series (~120–140 pm for M⁰, ~70–85 pm for M²⁺)

**Variation Between Series (3d → 4d → 5d):**

**3d to 4d:** Significant **increase** in atomic radius (e.g., Zr > Cr, Ti > V, etc.)

**4d to 5d:** Radii **remain virtually the same** (e.g., Zr ≈ Hf, Mo ≈ W) despite increased atomic number

**Cause of 4d ≈ 5d phenomenon — Lanthanoid Contraction**

**Lanthanoid Contraction Definition and Cause:**

**Definition:** Regular and gradual **decrease in ionic and atomic radii of lanthanoids** (4f series) from La to Lu

**Cause:** Filling of 4f orbitals **before** 5d orbitals begins; 4f electrons provide **poor shielding** (worse than d electrons)

As atomic number increases across 4f series, nuclear charge increases but 4f electrons do not shield each other effectively

Result: Each successive 4f electron experiences stronger effective nuclear charge

The entire 4f^n orbital becomes progressively smaller with increasing n

**Compensation Effect:** The **contraction from 4f filling** (~25 pm total) **almost exactly compensates** for the expected **increase in size from 3d to 4d** (e.g., Zr: 160 pm ≈ Hf: 159 pm)

**Chemical Consequence:** 4d and 5d transition elements of same group have **remarkably similar** physical and chemical properties (e.g., Zr and Hf are nearly inseparable), much more similar than corresponding 3d and 4d pairs

**Density Variation:**

**Trend:** **Increases significantly** from Ti (Z=22, ρ=4.11) to Cu (Z=29, ρ=8.9 g cm⁻³) in 3d series

**Cause:** Decrease in metallic radius combined with **increase in atomic mass** results in greater packing density

**Ionic Radius for M²⁺ and M³⁺:**

**M²⁺ ions:** 3d^n configuration; decrease from Sc²⁺ (not easily formed) to Zn²⁺ (~75 pm)

**M³⁺ ions:** [Ar]3d^(n-3) configuration; decrease from Sc³⁺ (73 pm) to Fe³⁺ (~65 pm)

**Stability:** M²⁺ more common than M³⁺ for later 3d elements

Ionisation Enthalpies

**First Ionisation Enthalpy (ΔᵢH° I):**

**Trend:** General increase from Sc to Zn (631 to 906 kJ mol⁻¹), but with **small magnitude of increase** (~275 kJ mol⁻¹ total)

**Comparison:** Much smaller variation than s and p-block elements

**Reason for small variation:**

New electrons enter inner d orbitals, not outermost shell

3d electrons provide moderate shielding of outermost 4s electrons

Atomic radii decrease only slightly, reducing effect on ionization energy

**Slight Irregularities:** Cr and Cu show slightly higher ΔᵢH° than neighbors

**Cr (3d⁵4s¹):** Half-filled d⁵ provides extra stability; easier to remove 4s¹ electron than might be expected

**Cu (3d¹⁰4s¹):** Filled d¹⁰ provides extra stability; slightly harder to remove 4s¹ than expected

**Second Ionisation Enthalpy (ΔᵢH° II):**

**Trend:** Increases significantly along 3d series (1235 to 1734 kJ mol⁻¹)

**Magnitude:** Increases much more steeply than ΔᵢH° I because removing d electrons requires overcoming stronger nuclear attraction

**Anomaly at Cr and Cu:**

**Cr:** Very high ΔᵢH° II because Cr⁺ has d⁵ configuration (extra stable)

**Cu:** Exceptionally high (1958 kJ mol⁻¹) because Cu⁺ has d¹⁰ configuration (highly stable)

**Significance:** High ΔᵢH° II for Cr and Cu explains why +1 oxidation state is relatively stable (e.g., CuCl, CuBr, but not easily CuCl₂ from Cu⁺)

**Third Ionisation Enthalpy (ΔᵢH° III):**

**Trend:** Generally increases from Sc³⁺ to Zn³⁺

**Very High Values:** Typically 2400–3800 kJ mol⁻¹ (two electrons already removed)

**Anomalies — Breaks at d⁵ and d¹⁰:**

**Mn²⁺ (d⁵) to Mn³⁺:** Very high ΔᵢH° III (~3260 kJ mol⁻¹) because d⁵ is extra stable; Mn⁺³ is rare

**Fe²⁺ (d⁶) to Fe³⁺:** Shows break in smooth trend; ΔᵢH° III of Fe lower than expected because Fe²⁺ has d⁶ (no special stability compared to d⁵)

**Zn²⁺ (d¹⁰) to Zn³⁺:** Extremely high (3837 kJ mol⁻¹); Zn³⁺ is virtually non-existent

**Chemical Implication:** Third ionisation is so difficult that transition metals rarely show oxidation states +3 or higher, except for late 3d metals

**Exchange Energy and Stability:**

**Definition:** Stabilization energy arising when electrons occupy degenerate orbitals with parallel spins (Hund's rule)

**Mechanism:** Parallel spin electron pairs in same orbital minimize electron repulsion by keeping electrons in different spatial regions

**Effect on Ionisation:** Removal of electron from half-filled (d⁵) or filled (d¹⁰) configuration requires loss of exchange energy, making ionization harder

**Example:**

Mn²⁺ (d⁵ with all parallel spins) → Mn³⁺: requires loss of exchange energy; ΔᵢH° III very high

Fe²⁺ (d⁶ with one paired electron) → Fe³⁺: no net loss of exchange energy; ΔᵢH° III relatively lower than Mn

**Formation of M²⁺ Ions:**

**Energy requirement:** ΔₐH° (atomisation) + ΔᵢH° I + ΔᵢH° II

**Dominant term:** ΔᵢH° II (removal of d electron requires strong nuclear attraction)

**Trend:** Generally decreases across 3d series except anomalies at Cr and Cu

**Result:** M²⁺ is most common oxidation state for transition metals

Oxidation States of Transition Metals

**Variable Oxidation States — Key Characteristic:**

Transition metals exhibit **multiple oxidation states**, unlike s and p-block elements which typically show one or two. The range and stability of oxidation states depend on:

Nuclear charge and effective nuclear charge

Exchange energy considerations

Stabilization by ligands (in complexes)

Hydration energy (in aqueous solutions)

**Range of Oxidation States:**

**3d series:** Range from 0 to +8

**Sc:** +3 only

**Ti:** +2, +3, +4

**V:** +2, +3, +4, +5

**Cr:** +2, +3, +6 (skip +4, +5)

**Mn:** +2, +3, +4, +6, +7

**Fe:** +2, +3, +6

**Co:** +2, +3 (rarely +4, +5)

**Ni:** +2, +3 (rarely)

**Cu:** +1, +2

**Zn:** +2 only

**Factors Stabilizing Particular Oxidation States:**

**1. Exchange Energy (Half-Filled d⁵):**

d⁵ configuration with all parallel spins is exceptionally stable

**Example:** Mn²⁺ has 3d⁵ and is highly stable; Mn³⁺ (removing one electron from d⁵) requires overcoming exchange energy

**Result:** Mn²⁺ much more stable than Mn³⁺

**2. Filled d¹⁰ Configuration:**

Completely filled d¹⁰ is stable

**Example:** Zn²⁺ (3d¹⁰) is the only stable oxidation state; Zn³⁺ virtually non-existent

**3. Electrode Potential (E° values):**

**Definition:** Reduction potential indicates ease of reduction; relates to oxidation state stability

**Interpretation:**

**Negative E°** (e.g., Mn²⁺/Mn = –1.18 V): M²⁺ is stable; M is easily oxidized

**Positive E°** (e.g., Cu²⁺/Cu = +0.34 V): M²⁺ is less easily reduced; Cu is hard to oxidize

**Very Positive E°** (e.g., Fe³⁺/Fe²⁺ = +0.77 V): Fe³⁺ is stable relative to Fe²⁺

**Common Oxidation States for 3d Series:**

**Lowest oxidation state:** Generally **+2** (loss of 4s electrons first)

**Exceptions:** Cu⁺ (+1) is stable due to d¹⁰ configuration

**Higher oxidation states:** Increase from left to right in series (Sc: +3 only; Cr: +6; Mn: +7)

**Trend in stability:** Higher oxidation states become relatively more stable as we move across the series, but are still generally unstable and require oxidizing conditions

**Stability of Oxidation States in Aqueous Solution:**

**Mn³⁺:** Unstable in aqueous solution (strong oxidizing agent, decomposes to Mn²⁺ and O₂)

**Fe³⁺:** Stable in acidic solution; precipitates in basic solution

**Cu⁺:** Unstable in aqueous solution (disproportionates: 2Cu⁺ → Cu²⁺ + Cu), but stable in insoluble compounds (CuCl, Cu₂O) and in complexes

**Cr³⁺:** Stable, widely used

**Cr⁶⁺:** Strong oxidizing agent, stable only in alkaline/neutral solutions (CrO₄²⁻, Cr₂O₇²⁻)

Characteristic Properties of Transition Elements

1. Variable Oxidation States

Already discussed above. Results from:

Similar energies of (n-1)d and ns orbitals

Exchange energy effects

Ligand stabilization in complexes

2. Formation of Coloured Compounds and Ions

**Definition:** Transition metals and their compounds display characteristic colours (unlike most s and p-block compounds which are colourless).

**Cause of Colour:**

**Presence of partially filled d orbitals** with electrons capable of absorbing visible light

**Mechanism:**

d orbitals split into different energy levels in crystal field (CFT) or ligand field

d-d electronic transitions occur when photons of visible light energy are absorbed

Energy of visible light (400–700 nm, corresponding to 170–300 kJ mol⁻¹) matches energy differences between d orbital levels

Absorbed light removed from white light, and transmitted/reflected light appears colored

**Examples:**

**Sc³⁺:** Colorless (no d electrons, d⁰)

**Ti³⁺:** Purple ([Ar]3d¹; one d electron can absorb light)

**V³⁺:** Green (3d²)

**Cr³⁺:** Green (3d³)

**MnO₄⁻:** Intense purple (d⁵ in tetrahedral field)

**Fe³⁺:** Yellow-brown (3d⁵ in octahedral field)

**Cu²⁺:** Blue (3d⁹; one d hole in octahedral field)

**Zn²⁺:** Colorless (3d¹⁰, completely filled)

**Contrast with s and p-block:**

s-block elements: all electrons paired, not capable of d-d transitions → colorless

p-block elements: color typically results from charge-transfer transitions (not d-d transitions)

3. Complex Formation

**Definition:** Transition metals and their ions form **coordinate complexes** with a wide variety of ligands (Lewis bases) that donate electron pairs to form coordinate covalent bonds.

**Reasons for Complex Formation:**

**Availability of d orbitals:** Can expand coordination number beyond typical VSEPR predictions

**Moderate charge density:** Sufficient to attract ligands but not so strong as to cause hydrolysis

**Variable oxidation states:** Allows tuning of complex stability and reactivity

**Examples:**

**[Fe(CN)₆]⁴⁻:** Ferrocyanide (colorless because of very large d orbital splitting)

**[Cu(NH₃)₄]²⁺:** Deep blue complex

**[CrO₄]²⁻:** Chromate ion (yellow)

**[Cr₂O₇]²⁻:** Dichromate ion (orange)

**Contrast with s and p-block:**

Mg²⁺, Al³⁺: form limited complexes, only in specific conditions (not characteristic)

Zn²⁺: forms some complexes but less varied than 3d elements

4. Catalytic Properties

**Definition:** Transition metals and their compounds act as **catalysts** — substances that accelerate reaction rates without being consumed.

**Mechanism:**

Multiple oxidation states allow participation in redox cycles

Complex formation with reactants/intermediates

Variety of coordination environments enables activation of different types of bonds

**Industrial Examples:**

**Fe:** Haber process (N₂ + H₂ → NH₃) catalyst

**Fe³⁺:** Catalyzes decomposition of H₂O₂

**Ni:** Hydrogenation of unsaturated hydrocarbons

**V₂O₅:** Contact process for SO₃ production (2SO₂ + O₂ → 2SO₃)

**Pt:** Catalytic converters in automobiles (CO + NO oxidation)

**MnO₂:** Decomposition of KMnO₄ or H₂O₂

**Enzymes:** Contain transition metal ions (Fe²⁺ in cytochromes, Cu²⁺ in oxidases)

**Why effective catalysts:**

Can easily change oxidation states (allows electron transfer)

Form intermediate complexes with reactants

Provide alternative reaction pathway with lower activation energy

5. Magnetic Properties — Paramagnetism

**Definition:** Paramagnetic behavior — attraction to magnetic field due to unpaired electrons

**Cause:**

Presence of **unpaired d electrons** creates net magnetic moment

Each unpaired electron has spin magnetic moment; in paramagnetic species, these are not cancelled

**Magnetic Moment Calculation:**

**Formula:** μ = √[n(n+2)] Bohr magnetons (μB), where n = number of unpaired electrons

**Bohr magneton:** μB = 9.274 × 10⁻²⁴ J T⁻¹

**Examples:**

**Sc³⁺:** [Ar]3d⁰; diamagnetic (no unpaired electrons)

**Ti³⁺:** [Ar]3d¹; paramagnetic (1 unpaired electron); μ = √[1(3)] = 1.73 μB

**Mn²⁺:** [Ar]3d⁵; highly paramagnetic (5 unpaired electrons); μ = √[5(7)] = 5.92 μB

**Fe³⁺:** [Ar]3d⁵; highly paramagnetic (5 unpaired electrons); μ = √[5(7)] = 5.92 μB

**Cu²⁺:** [Ar]3d⁹; paramagnetic (1 unpaired electron); μ = √[1(3)] = 1.73 μB

**Zn²⁺:** [Ar]3d¹⁰; diamagnetic (no unpaired electrons)

**Experimental Methods:**

Magnetic susceptibility measurement

Gouy balance method

Vibrating sample magnetometer

**Contrast with s and p-block:**

Most are diamagnetic (paired electrons)

Some p-block radicals and ions can be paramagnetic, but less common

Important Compounds of 3d Transition Elements

Potassium Permanganate (KMnO₄)

**Physical Properties:**

Deep purple crystalline solid

Soluble in water (purple solution)

Melting point: 513 K

**Electronic Configuration of Mn:**

Ground state: [Ar]3d⁵4s²

In KMnO₄: Mn in **+7 oxidation state** (MnO₄⁻ ion)

Configuration of Mn⁷⁺: [Ar] (all d electrons removed)

**Structure of MnO₄⁻:**

**Tetrahedral geometry**

Mn at center bonded to 4 oxygen atoms

All Mn–O bonds are equivalent (resonance structures)

High oxidation state and small size of Mn creates strong polarizing power

**Preparation:**

**Industrial Method (Fusion Method):**

1. Oxidize solid MnO₂ with KClO₃ and KOH (molten):

2MnO₂ + 2KClO₃ + 4KOH → 2KMnO₄ + 2KCl + O₂ + 2H₂O

2. Or oxidize MnO₂ in alkaline solution:

MnO₂ + 2KOH + oxidizing agent → KMnO₄ + byproducts

**Laboratory Method (Fusion of MnO₂ with KOH):**

MnO₂ + 2KOH + [O] → KMnO₄ + H₂O (where [O] from heat or K₂S₂O₃)

**Properties — Redox Behavior:**

KMnO₄ is a **strong oxidizing agent** in both acidic and basic media. The reduction product (Mn) depends on pH:

**1. In Acidic Solution:**

**Half-reaction:** MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

**E° = +1.49 V** (very strong oxidant)

**Reduction product:** Mn²⁺ (colorless)

**Color change:** Purple MnO₄⁻ → colorless Mn²⁺

**Indicators:** Used as self-indicator in titrations (first drop of excess KMnO₄ imparts purple color)

**2. In Neutral Solution:**

**Half-reaction:** MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂O (or in neutral pH, more complex)

**Reduction product:** MnO₂ (brown/black precipitate)

**Color change:** Purple → brown/black

**3. In Alkaline Solution:**

**Half-reaction:** MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻

**Reduction product:** MnO₂ (brown/black)

**Color change:** Purple → colorless as MnO₂ precipitate forms, but solution may appear greenish-yellow due to MnO₄²⁻ (manganate, green)

**General Oxidation Reactions of KMnO₄:**

**Organic Substrates in Acidic Medium:**

1. **Alcohols → Carbonyl compounds:**

Primary alcohols: → Carboxylic acids

Secondary alcohols: → Ketones

Tertiary alcohols: Not oxidized (no H on C-OH)

Example: 3CH₃CH₂OH + 2KMnO₄ + 3H₂SO₄ → 3CH₃CHO + K₂SO₄ + 2MnSO₄ + 4H₂O

2. **Aldehydes → Carboxylic acids:**

Example: 3RCHO + 2KMnO₄ + H₂SO₄ → 3RCOOH + K₂SO₄ + 2MnSO₄ + H₂O

3. **Alkenes → Diols (cold dilute) or Cleavage (hot concentrated):**

Cold dilute: adds OH across C=C

Hot concentrated: cleaves C=C to carboxylic acids

Example: CH₂=CH₂ + KMnO₄ → CH₂(OH)CH₂(OH) (ethylene glycol, colorless)

4. **Alkanes:** Not oxidized by KMnO₄ (unless benzylic position)

**In Basic Solution:**

Similar oxidations but may produce different final products

Organic acids → Salts or CO₂

Aldehydes → Carboxylates

**Inorganic Substrates:**

1. **Fe²⁺ → Fe³⁺:**

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

2. **H₂O₂ → O₂ and H₂O (in acidic medium):**

2MnO₄⁻ + 5H₂O₂ + 6H⁺ → 2Mn²⁺ + 5O₂ + 8H₂O

3. **SO₃²⁻/SO₂ → SO₄²⁻:**

2MnO₄⁻ + 5SO₃²⁻ + 4H⁺ → 2Mn²⁺ + 5SO₄²⁻ + 2H₂O

4. **S²⁻/S → S⁰ or SO₄²⁻:**

2MnO₄⁻ + 5H₂S + 3H⁺ → 2Mn²⁺ + 5S + 8H₂O (yellow precipitate)

**Applications:**

**Analytical:** Volumetric analysis (permanganometry); self-indicating titrant

**Disinfectant:** Water purification

**Bleaching agent:** Color removal in textiles

**Medical:** Antiseptic (dilute solutions)

Potassium Dichromate (K₂Cr₂O₇)

**Electronic Configuration of Cr:**

Ground state: [Ar]3d⁵4s¹ (anomalous d⁵ configuration for stability)

In K₂Cr₂O₇: Cr in **+6 oxidation state**

Configuration: [Ar] (all d electrons removed from neutral Cr)

**Structure of Cr₂O₇²⁻:**

Two CrO₄ tetrahedra sharing one oxygen vertex

Central oxygen bridges two Cr atoms

All Cr–O bonds approximately equivalent due to resonance

**Color:**

Orange-red solid and solution

Intense due to charge-transfer transitions (O → Cr)

**Preparation:**

**Industrial Method:**

1. Oxidation of Cr ore (FeCr₂O₄) in molten alkali with oxidizing agent

2. Convert Cr₂O₃ to Na₂CrO₄ by fusion with soda ash and oxidant:

Cr₂O₃ + 4Na₂CO₃ + 3O₂ → 2Na₂CrO₄ + 4CO₂

3. Convert CrO₄²⁻ to Cr₂O₇²⁻ by acidification:

2CrO₄²⁻ + 2H⁺ → Cr₂O₇²⁻ + H₂O

**Laboratory Method:**

Heat Na₂CrO₄ or K₂CrO₄ with H₂SO₄ (acidify):

2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O (equilibrium shifts right with acid)

**Redox Properties:**

**Half-Reaction in Acidic Solution:**

**Reduction:** Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O

**E° = +1.33 V** (strong oxidant, but slightly weaker than MnO₄⁻)

**Reduction product:** Cr³⁺ (green)

**Color change:** Orange Cr₂O₇²⁻ → green Cr³⁺

**In Neutral or Basic Solution:**

Less stable; converts to CrO₄²⁻ (yellow) with OH⁻

Oxid