The Many Colors of Iron

A tour through seven iron compounds — from grey metal to blue-black ink and orange rust

Difficulty: Easy–Medium | Time: 45–60 min | Visual Impact: High

Historical Context

Iron has colored the human story since the Bronze Age gave way to the Iron Age around 1200 BCE. The yellowy-brown of limonite, the deep red of hematite, the greenish glint of siderite — ancient miners recognized these iron ores by sight. The ochre pigments that prehistoric artists ground into cave paintings at Lascaux are iron oxides, the same compounds you will make in Sample 5.

The greenish crystals of “green vitriol” (ferrous sulfate, FeSO₄) were prized by medieval alchemists. Those same crystals, combined with oak gall tannins, produced iron gall ink — the dominant writing ink in the Western world for over a thousand years. The Magna Carta (1215), Leonardo da Vinci’s notebooks, Bach’s musical scores, and the original Declaration of Independence were all written with it.

What makes iron particularly interesting to a chemist is the sharp color difference between its two common oxidation states. Fe²⁺ gives pale, greenish hues; Fe³⁺ gives deep ambers and reds. One electron separates them, yet the visual contrast is dramatic. This tour shows why.

Materials

What How much
Ferrous sulfate (FeSO₄·7H₂O) 5 g
Ferric chloride (FeCl₃·6H₂O) 5 g
Sodium carbonate (Na₂CO₃) 4 g
Tannic acid 0.5 g
Ascorbic acid (vitamin C) 0.5 g
Iron nail or pinch of steel wool 1 piece
Water ~300 mL
Test tubes or small glass jars 6–8
Safety glasses

Optional: Strong black tea works as a substitute for tannic acid in Samples 6–7, and shows how natural tannin sources compare.

Colour Guide

Here is what you are building toward:

# Colour Compound Iron oxidation state
1 Silver-grey Fe⁰ metal Fe⁰
2 Pale blue-green FeSO₄ solution Fe²⁺ with water ligands
3 Amber-yellow FeCl₃ solution Fe³⁺ with water/chloride ligands
4 Pale green → orange-brown FeCO₃ oxidising in air Fe²⁺ → Fe³⁺ (rust in a beaker)
5 Red-brown Fe(OH)₃ precipitate Fe³⁺
6 Blue-gray → blue-black Fe(II)-tannate → Fe(III)-tannate Fe²⁺ oxidising to Fe³⁺
7 Amber → pale green FeCl₃ + vitamin C Fe³⁺ reduced back to Fe²⁺

Procedure

Work in order. Set up Samples 1–3 first as your permanent color references, then do Samples 4–7 in sequence, as some require watching over time.

Sample 1 — Silver-grey metal

  1. Hold an iron nail or pinch of steel wool. This is Fe⁰. The grey metallic surface is iron’s ground state. Set it aside as your reference.

Sample 2 — Pale blue-green Fe²⁺

  1. Dissolve 2 g of ferrous sulfate in 50 mL of water. The pale, slightly blue-green solution contains Fe²⁺ ions surrounded by water molecules. If your crystals are yellowish (partially oxidized), the solution will be tinged yellow rather than blue-green. Pour into a test tube and label it. Note: Fe²⁺ oxidizes slowly in air — work quickly and compare with Sample 3 before the color shifts.

Sample 3 — Amber Fe³⁺

  1. Dissolve 2 g of ferric chloride in 50 mL of water. The solution is a clear amber-yellow to orange-brown. This is Fe³⁺, iron one electron more oxidized than Sample 2. Place both test tubes side by side — the contrast between pale green and deep amber is striking for a one-electron difference.

Sample 4 — Rust formation (watch over 20 minutes)

  1. Dissolve 1 g of ferrous sulfate in 30 mL of water in a wide jar or beaker. Add 1 g of sodium carbonate dissolved in 10 mL of water and stir. A pale green precipitate (iron(II) carbonate, FeCO₃) settles immediately. Leave the jar uncovered and check every 5–10 minutes: the surface slowly turns orange-brown as atmospheric oxygen oxidizes Fe²⁺ to Fe³⁺, forming hydrated iron(III) oxide — rust. You are watching corrosion chemistry happen in real time.

Sample 5 — Red-brown Fe(OH)₃

  1. Dissolve 1 g of ferric chloride in 20 mL of water (amber). In a separate cup, dissolve 1 g of sodium carbonate in 10 mL of water. Pour the carbonate into the iron chloride while stirring. A voluminous reddish-brown precipitate forms immediately: iron(III) hydroxide, Fe(OH)₃ — the compound responsible for rust color and the ancient ochre earth pigments. Compare it directly with the slowly-oxidizing green precipitate from Sample 4.

Sample 6 — Iron gall ink developing

  1. In a small glass, dissolve 0.5 g of tannic acid in 20 mL of water — this gives a pale yellow solution. In a separate glass, dissolve 0.5 g of ferrous sulfate in 10 mL of water (pale green). Pour the iron solution into the tannic acid and stir. The mixture turns a striking grey-blue within seconds as Fe²⁺ coordinates with the tannin molecules. Leave it uncovered. Every 5 minutes, note the color: it will deepen progressively through blue-gray to a permanent blue-black as the iron is oxidized in air to Fe³⁺. After 15–20 minutes, dip a toothpick or pen and write on paper — this is iron gall ink.

Sample 7 — Reduction: turning amber back to green

  1. Pour 20 mL of ferric chloride solution (amber, Sample 3) into a fresh test tube. Dissolve 0.5 g of ascorbic acid (vitamin C) in 5 mL of water and add it slowly with swirling. Watch the amber color fade toward pale green as ascorbic acid donates electrons to Fe³⁺, reducing it to Fe²⁺. This is a real-time reduction reaction, and the same mechanism by which vitamin C acts as an antioxidant in the body.

Reactions

Sample 4 — Rust formation

\[\ce{FeSO4(aq) + Na2CO3(aq) -> FeCO3(s) + Na2SO4(aq)}\]

In air, the Fe²⁺ carbonate slowly oxidizes:

\[\ce{4FeCO3(s) + O2(g) + 3H2O(l) -> 4Fe(OH)3(s) + 4CO2(g)}\]

Sample 5 — Hydroxide precipitation

\[\ce{FeCl3(aq) + Na2CO3(aq) + 3H2O(l) -> 2Fe(OH)3(s) + 2NaCl(aq) + CO2(g)}\]

(Fe³⁺ is a strong enough Lewis acid that carbonate immediately hydrolyzes to give hydroxide.)

Sample 6 — Iron gall ink

Initial complexation (blue-gray, soluble):

\[\ce{Fe^{2+}(aq) + tannin -> [Fe^{II}\text{-tannate}](aq)}\]

Air oxidation gives the permanent blue-black, insoluble form:

\[\ce{[Fe^{II}\text{-tannate}](aq) + \tfrac{1}{4}O2(g) + \tfrac{1}{2}H2O(l) -> [Fe^{III}\text{-tannate}](s) + OH^-(aq)}\]

Sample 7 — Reduction by vitamin C

\[\ce{2Fe^{3+}(aq) + C6H8O6(aq) -> 2Fe^{2+}(aq) + C6H6O6(aq) + 2H^+(aq)}\]

Ascorbic acid (C₆H₈O₆) is oxidized to dehydroascorbic acid (C₆H₆O₆).

The Science

Iron is a d-block transition metal whose two common oxidation states, Fe²⁺ (d⁶) and Fe³⁺ (d⁵), produce strikingly different colors. Understanding why requires looking at how electrons absorb light.

Fe²⁺ — pale blue-green

In aqueous solution, Fe²⁺ is surrounded by six water molecules. It has six d-electrons and can undergo d–d transitions when it absorbs light. However, these transitions are symmetry-forbidden (Laporte’s rule), so the absorption is weak. The result: a pale, washed-out blue-green.

Fe³⁺ — deep amber

Fe³⁺ has five d-electrons, one in each orbital with all spins parallel. Any d–d transition would require a spin flip — doubly forbidden. So why is Fe³⁺ so much more intensely colored? The amber comes primarily from ligand-to-metal charge transfer (LMCT): electrons jump from the oxygen or chloride ligands to the iron center, absorbing violet light and transmitting yellow-orange. Charge-transfer transitions are allowed by symmetry and tend to be far more intense than d–d bands.

This explains a broader pattern: Fe³⁺ compounds are often much more deeply colored than Fe²⁺ compounds, even though quantum mechanics predicts Fe²⁺ should have more d-orbital transitions available.

Species Color origin Color
Fe²⁺ (aq) Weak d–d transition Pale blue-green
Fe³⁺ (aq) LMCT (O → Fe) Amber
Fe(OH)₃ LMCT + charge transfer Red-brown
Fe(III)-tannate Strong LMCT Near-black

The tannate colors

Tannic acid has many phenolic –OH groups that donate electrons to iron very effectively. The resulting charge-transfer absorption is so strong that even dilute solutions appear nearly black. The switch from Fe²⁺-tannate (blue-gray) to Fe³⁺-tannate (blue-black) is also a color-deepening on oxidation, because Fe³⁺ accepts electron density from the tannin more forcefully than Fe²⁺.

Iron and life

The Fe²⁺/Fe³⁺ redox couple is at the center of biology. Hemoglobin carries oxygen with Fe²⁺ at its heme center; oxidation to Fe³⁺ produces methemoglobin, which cannot bind oxygen and causes cyanosis. Vitamin C (Sample 7) keeps other iron-containing enzymes in their correct oxidation states — which is why its role as a biological reductant matters as much as its role as a vitamin.

Explore Further

Accelerate oxidation with H₂O₂: Instead of waiting for atmospheric oxygen to oxidize Fe²⁺, add a few drops of hydrogen peroxide to a fresh ferrous sulfate solution (Sample 2). Watch the color shift from pale green to amber within seconds — the same one-electron change that takes hours by air oxidation. Compare the speed to understand why H₂O₂ is such a potent oxidizer.

Accelerate the rust: Repeat Sample 4 but add a pinch of sodium chloride to the beaker. Does the orange color appear faster? Salt accelerates electrochemical corrosion because chloride ions are good conductors and disrupt the protective oxide layer. This is why iron corrodes faster near the sea.

Natural tannin sources: Brew very strong black tea and use it in place of tannic acid in Sample 6. Does the ink color match? Tea tannins (theaflavins, thearubigins) are structurally related to tannic acid and produce functional ink. Try red wine as well — the comparison shows how tannin concentration and type affect the color depth.

Reversibility of the ink: Take some of your blue-gray iron-tannate solution before it fully darkens. Add a pinch of ascorbic acid. The color lightens slightly as you reduce Fe³⁺ back toward Fe²⁺. This is the battle that manuscript conservators fight — iron gall ink oxidizes and destroys the paper it was written on, and stabilizing it requires careful redox chemistry.

Connect to Prussian blue: The Prussian Blue experiment uses both Fe²⁺ and Fe³⁺ ions together in the same crystal lattice, creating an entirely different color mechanism — intervalence charge transfer between adjacent iron centers. After completing this tour (pale green, amber, blue-black, orange rust), follow it with Prussian blue’s electric, saturated blue to see the full color range that a single element can produce.

Chemicals used in this experiment: