Chemistry: Atoms First 2e · Transition Metals and Coordination Chemistry
Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
On this page 9 sections
In 30 seconds
The transition metals are the elements in the d block of the periodic table — the ten columns (groups 3–12) between the s-block and p-block elements. Their defining feature is a partially filled d subshell: as electrons fill the d orbitals, the outermost s electrons are already in place, so chemistry is dominated by d-orbital electrons. This electronic arrangement produces the properties we associate with metals at their best — high melting points, high density, metallic luster, excellent conductivity — plus two signatures found almost nowhere else: variable oxidation states and colored compounds. Iron rusts in shades of orange, copper patinas green, chromium gives rubies their red, and manganese dioxide blackens battery cathodes; all of that color comes from d-electron transitions.
This topic covers where transition metals are found in nature (almost always as ores, never free except for gold, silver, and platinum), how they are extracted and purified (Roasting Heating a sulfide ore in air to convert it to an oxide, releasing SO2 Full entry →, Smelting Melting and chemically reducing an ore to free the metal Full entry →, Reduction A process in which a substance gains electrons (oxidation state decreases) Full entry →, and Electrolysis Using electric current to drive a nonspontaneous redox reaction Full entry →), and the physical and chemical properties that make them the structural and electronic backbone of modern technology — from steel beams to catalytic converters.
Why this matters
- Materials and infrastructure: Steel (iron plus carbon and other metals), titanium alloys for aircraft, and copper wiring carry modern civilization; their properties come straight from transition-metal chemistry.
- Catalysis: Iron catalyzes ammonia synthesis (Haber process), platinum and palladium clean car exhaust in catalytic converters, and vanadium(V) oxide is the workhorse catalyst of sulfuric acid production.
- Biology and health: Iron carries oxygen in hemoglobin, cobalt sits at the center of vitamin B12, and zinc is required by hundreds of enzymes; transition-metal ions are also used as MRI contrast agents and in chemotherapy (cisplatin).
- Everyday objects: Coinage metals (Cu, Ag, Au), stainless steel, rechargeable batteries (lithium cobalt oxide cathodes), and pigments in paints, glazes, and gems are all transition-metal chemistry.
- Exams: Electron configurations of d-block ions, oxidation-state calculations, and the chemistry of the ore-to-metal process are high-frequency topics.
The college version
Core Concepts
Where transition metals are found: ores, not free metals
Because most transition metals are reactive enough to combine with oxygen, sulfur, or carbon dioxide, they occur in nature as ores — minerals from which a metal can be profitably extracted. Important examples include:
- Oxides: hematite (Fe2O3), magnetite (Fe3O4), rutile (TiO2), and chromite (FeCr2O4).
- Sulfides: chalcopyrite (CuFeS2), galena (PbS), sphalerite (ZnS), and cinnabar (HgS).
- Carbonates and other salts: malachite (Cu2CO3(OH)2), siderite (FeCO3).
Gold, silver, and platinum are the exceptions — they are found free (native) because they are so unreactive. Even these are usually concentrated by physical separation, since ores are rarely rich enough to use directly.
Concentration and roasting: preparing the ore
Before reduction, ores are concentrated to remove rocky waste (gangue). Sulfide ores are then roasted — heated in air — to convert them to oxides and drive off sulfur dioxide:
\[2ZnS + 3O_2 \rightarrow 2ZnO + 2SO_2\]
Carbonate ores are heated (calcined) to drive off carbon dioxide. The oxide left behind is the starting material for reduction. (Roasting releases SO2, a pollutant, so modern smelters capture it and convert it to sulfuric acid — a textbook example of industrial waste becoming a product.)
Reduction to the free metal
The oxide is reduced to the metal, most commonly with carbon or carbon monoxide in a blast furnace (for iron):
\[Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2\]
or with a more reactive metal. Aluminum is too reactive for carbon reduction; it is produced by electrolysis of molten Al2O3 (the Hall–Héroult process). Titanium is made by the Kroll process: TiCl4 (from rutile) is reduced with molten magnesium:
\[TiCl_4 + 2Mg \rightarrow Ti + 2MgCl_2\]
The choice of method depends on how easily the metal's oxide can be reduced — the more reactive the metal, the stronger the reducing agent (or electricity) required.
Physical properties: what makes them "transition"
- High melting and boiling points: d electrons participate in metallic bonding, making these metals hard and refractory (tungsten melts at 3422 °C, the highest of any metal).
- High density and strength: atoms pack tightly; osmium and iridium are the densest elements known.
- Luster and conductivity: free-flowing d and s electrons reflect light and carry current.
- Variable oxidation states: because ns and (n-1)d electrons have similar energies, transition metals can lose different numbers of electrons — iron shows +2 and +3, manganese +2 through +7, chromium +2, +3, and +6. This is the single most important chemical difference from main-group metals, which almost always show one (or two closely related) oxidation states.
Chemical properties: colors, complexes, and catalysis
- Colored compounds: d-d electron transitions absorb visible light, so many compounds are intensely colored — K2Cr2O7 is orange, KMnO4 is deep purple, copper(II) sulfate is blue. (This theme returns in the spectroscopic topic later in the chapter.)
- Complex-ion formation: transition-metal ions accept electron pairs from ligands (water, ammonia, chloride) to form complex ions such as [Cu(NH3)4]2+ and [Fe(CN)6]4-.
- Catalytic activity: the ability to change oxidation state lets transition metals adsorb reactants, transfer electrons, and lower activation energies — the basis of the Haber, contact, and catalytic-converter processes.
- Magnetism: unpaired d electrons make many ions and compounds paramagnetic (attracted into a magnetic field), which is exploited in MRI contrast agents.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Transition metals | Main-group metals (e.g., Al, Sn, Pb) | Transition metals have partially filled d subshells → variable oxidation states, colored compounds, complexes; main-group metals show one main oxidation state |
| Roasting | Smelting | Roasting heats sulfide ores in air to make oxides (chemical change, may not melt); smelting melts and reduces ore to free the metal |
| Reduction with carbon | Reduction with electrolysis | Carbon works for moderately reactive metals (Fe, Cu); electrolysis is required for reactive metals (Al) whose oxides carbon cannot reduce |
| Oxidation state | Ionic charge | Oxidation state is a bookkeeping number from electronegativity rules; it can exceed the ionic charge (e.g., Mn in KMnO4 is +7, but no Mn⁷⁺ ion exists in aqueous solution) |
| Zinc, cadmium, mercury | Typical transition metals | Their d subshells are full (d10) in the atom and common ions, so they lack variable oxidation states and colored compounds — often treated separately |
| All colored solids in nature being transition-metal compounds | Color always arising from d-d transitions | Many colored materials (iodine vapor, organic dyes) are colored for other reasons; but the intense colors of most metal compounds do come from d-electron transitions |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Transition metals are the elements in the middle of the periodic table that love to change their "mood" — the same metal can lose different numbers of electrons, which is why iron can be rust-colored in one compound and black in another. They make strong, shiny, colorful materials: steel buildings, copper wires, gold rings, and the purple stuff in your chemistry lab. They're also super helpful as catalysts — little helpers that make reactions go faster without being used up, like in a car's catalytic converter.
Worked example
Example 1: Stoichiometry of iron production from hematite
A blast furnace reduces hematite with carbon monoxide: Fe2O3 + 3CO → 2Fe + 3CO2. How many grams of iron can be produced from 1.00 kg of pure hematite?
Step 1 — Molar masses. Fe2O3: 2(55.85) + 3(16.00) = 159.70 g/mol. Fe: 55.85 g/mol.
Step 2 — Convert kg → g → mol Fe2O3 → mol Fe → g Fe. The mole ratio from the equation is 1 mol Fe2O3 : 2 mol Fe.
\[m_{Fe} = 1000\ \text{g Fe2O3} \times \frac{1\ \text{mol Fe2O3}}{159.70\ \text{g Fe2O3}} \times \frac{2\ \text{mol Fe}}{1\ \text{mol Fe2O3}} \times \frac{55.85\ \text{g Fe}}{1\ \text{mol Fe}} = 699\ \text{g Fe}\]
Unit check: g Fe2O3 → mol Fe2O3 → mol Fe → g Fe. One kilogram of pure hematite yields 699 g of iron — about 70% of the ore's mass, matching the mass fraction of iron in Fe2O3 (111.7159.7 = 0.699).
Example 2: Assigning oxidation states in transition-metal compounds
Potassium dichromate, K2Cr2O7, is a common oxidizing agent. What is the oxidation state of chromium?
Step 1 — Rules. K is +1 (Group 1), O is −2 (usual), and the compound is neutral, so the sum of oxidation states is zero.
Step 2 — Set up the equation and solve for Cr:
\[2(+1) + 2x + 7(-2) = 0 \qquad \Rightarrow \qquad 2 + 2x - 14 = 0 \qquad \Rightarrow \qquad x = +6\]
So chromium is in the +6 state in K2Cr2O7. Check with potassium permanganate, KMnO4:
\[+1 + x + 4(-2) = 0 \qquad \Rightarrow \qquad x = +7\]
Manganese shows +7 in permanganate, the highest oxidation state of any element in a common laboratory reagent. The same algebra works for any compound: define the unknown, write the zero-sum equation, solve.
Example 3: Kroll process — titanium from titanium(IV) chloride
Titanium is produced by TiCl4 + 2Mg → Ti + 2MgCl2. What mass of titanium forms when 9.50 g of TiCl4 reacts with excess magnesium?
Step 1 — Molar masses. TiCl4: 47.87 + 4(35.45) = 189.67 g/mol. Ti: 47.87 g/mol.
Step 2 — Mole ratio: 1 mol TiCl4 : 1 mol Ti.
Step 3 — Convert:
\[m_{Ti} = 9.50\ \text{g TiCl4} \times \frac{1\ \text{mol TiCl4}}{189.67\ \text{g TiCl4}} \times \frac{1\ \text{mol Ti}}{1\ \text{mol TiCl4}} \times \frac{47.87\ \text{g Ti}}{1\ \text{mol Ti}} = 2.40\ \text{g Ti}\]
Unit check: g TiCl4 → mol TiCl4 → mol Ti → g Ti. Because the mole ratio is 1:1, the mass of titanium is just the mass of TiCl4 scaled by the ratio of molar masses (47.87/189.67 ≈ 0.252), giving 2.40 g.
Key takeaways
- Transition metals = d-block elements (groups 3–12) with partially filled d subshells; they lose s electrons first, then d electrons.
- Occurrence: mostly as oxide, sulfide, and carbonate ores; Au, Ag, Pt occur free (native).
- Ore-to-metal workflow: concentration → roasting (sulfides to oxides) → reduction (C, CO, more reactive metal, or electrolysis).
- Iron: blast furnace, Fe2O3 + 3CO → 2Fe + 3CO2. Aluminum: electrolysis (Hall–Héroult). Titanium: Kroll process, TiCl4 + 2Mg → Ti + 2MgCl2.
- Variable oxidation states: Fe +2/+3, Mn +2→+7, Cr +2/+3/+6 — the hallmark of transition-metal chemistry.
- Colored compounds arise from d-d transitions; complex ions form with ligands; many ions are paramagnetic.
- Transition metals are key industrial catalysts (Fe in Haber, Pt/Pd in catalytic converters, V2O5 in contact process).
Check yourself
6 review questions from the chapter. Try each one, then open the answer.
What electronic feature defines a Transition metal A d-block element with a partially filled d subshell (in the atom or a common ion) Full entry →, and what two chemical properties follow directly from it?
Show answer
A partially filled d subshell in the atom or a common ion. It produces variable oxidation states (different numbers of d electrons can be lost) and colored compounds (visible-light d-d transitions).
Name three types of ores with one example mineral each.
Show answer
Oxide: hematite Fe2O3. Sulfide: chalcopyrite CuFeS2. Carbonate: malachite Cu2CO3(OH)2 or siderite FeCO3.
Write the balanced equation for roasting zinc sulfide, and identify the two products.
Show answer
2ZnS + 3O2 → 2ZnO + 2SO2. Products: zinc oxide and sulfur dioxide.
Why is aluminum produced by electrolysis rather than by carbon reduction, but iron is not?
Show answer
Aluminum's oxide is too stable for carbon to reduce (Al is very reactive); electrolysis supplies the extra energy. Iron's oxide is reduced easily by carbon monoxide in the blast furnace.
Determine the oxidation state of manganese in KMnO4 and of chromium in Cr2O3, showing your equations.
Show answer
KMnO4: +1 + x + 4(-2) = 0 → x = +7. Cr2O3: 2x + 3(-2) = 0 → x = +3.
How many grams of copper can be obtained from 159.0 g of Cu2O (molar mass 143.10 g/mol) if reduction by hydrogen goes to completion (Cu2O + H2 → 2Cu + H2O; Cu = 63.55 g/mol)?
Show answer
159.0 g Cu2O × (1 mol/143.10 g) × (2 mol Cu/1 mol Cu2O) × (63.55 g/mol) = 141.2 g Cu.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Transition metal
- A d-block element with a partially filled d subshell (in the atom or a common ion)
- Ore
- A mineral from which a metal can be extracted profitably
- Roasting
- Heating a sulfide ore in air to convert it to an oxide, releasing SO2
- Smelting
- Melting and chemically reducing an ore to free the metal
- Reduction
- A process in which a substance gains electrons (oxidation state decreases)
- Electrolysis
- Using electric current to drive a nonspontaneous redox reaction
- Variable oxidation state
- The ability of an element to lose different numbers of electrons in different compounds
- Complex ion
- A metal ion surrounded by ligands (electron-pair donors) bonded through coordinate covalent bonds
- Native metal
- A metal found uncombined in nature
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
Educational content only. It is not medical, legal or professional advice. Found an error? Tell us.