Transition metals: transition metals periodic table facts

What are transition metals? definition and classification

A transition metal is a metallic element whose atoms fill the d orbitals of the next-to-outermost shell. The IUPAC 2011 nomenclature rules count every element in groups 3 to 12 as part of the family, which lines up exactly with the d-block. A good page on the tin boiling point covers stable isotopes, oxidation states and why stannous and stannic compounds behave so differently.

An older IUPAC wording is stricter: an element qualifies only if its atom, or one of its common ions, has a partially filled d subshell. That narrower test is why a few elements sit on the fence, as you'll see below.

In practice you can spot a member of the family on the periodic table by shared traits: they are lustrous, dense, hard conductors with ions of more than one charge, brightly coloured compounds and a talent for speeding up reactions.

Why they are called transition elements

The English chemist Charles Bury coined the term in 1921 to describe a run of elements in which an inner layer of electrons changes from a stable group of 8 to a stable group of 18. The name also fits their position: transition elements sit between the reactive alkaline earth metals of group 2 on their left and the p-block elements that lead across to the noble gases on the right, and their chemistry is a bridge between the two.

Inner transition metals: the F-block lanthanide and actinide series

The two rows printed under the main table are the inner transition metals, also called inner transition elements. They fill the f orbitals two shells below the outermost one, which is why they form the separate f-block.

The lanthanide row runs from lanthanum (57) onward and the actinide row from actinium (89). Some textbooks fold them into a combined count of about 68 metals, but most chemists treat them as a family of their own because their chemistry is governed by f electrons rather than d electrons.

Where transition metals sit on the periodic table

On the standard 18-column periodic table, the d-block occupies the ten columns between the s-block and the p-block, starting in period 4. Scandium (atomic number 21) opens the first series straight after calcium, and each step to the right adds one more electron to the same d subshell until it is full at zinc (atomic number 30). Directly to the right of group 12 come the post-transition metals such as gallium, tin and lead, which are softer and have lower melting points. For nuclear medicine imaging, technetium-99m and pertechnetate are the focus of the technetium density guide for readers.

Early and late transition metals

Chemists often split the block by position. Early transition metals in groups 3 to 7 hold few d electrons and tend to reach high oxidation states and bind tightly to oxygen. Late transition metals in groups 8 to 11 (or 12) carry fuller d shells, favour lower charges and include the noble metals.

A three-way scheme adds a middle set for groups 6 to 8. In group 3 itself, a 2021 IUPAC report supports placing lutetium and lawrencium below yttrium rather than lanthanum and actinium.

The group 12 question: zinc, cadmium and mercury

Group 12 is the most disputed column. Zinc, cadmium and mercury keep a full d¹⁰ shell in every common compound, so they fail the "partially filled" test and are often grouped as post-transition metals.

Even so, many textbooks keep them in because they sit in the d-block. Their full d shells also explain why their metallic bonding is weak:

• Zinc melts at only 420 °C, against 1,085 °C for copper next door.
• Mercury is a liquid at room temperature, melting at −38.83 °C, partly because of relativistic effects on its 6s electrons.
• Copernicium, the superheavy member, is predicted to use its 6d electrons, though it has barely been studied.

Transition metal list for groups 3-12

Use this transition metal list to find any member by group and period. Each cell gives the symbol and atomic number. Groups 3-12 across periods 4 to 7 hold 40 positions; the often-quoted total of 38 leaves out the two superheavy elements whose chemistry is still unconfirmed.

Everything in period 7 is synthetic, made a few atoms at a time in accelerators, so the bulk properties described on this page come from periods 4 to 6.

Transition metal examples in everyday life

You handle these elements constantly, usually as alloys:

• Iron is the base of steel, and chromium and nickel turn it into stainless steel.
• Titanium gives aircraft frames and hip implants strength at low weight.
• Copper and silver carry current in wiring and electronics technology.
• Gold, platinum and palladium resist corrosion, which suits jewellery and electrical contacts.
• Tungsten survives drill bits and furnace parts, while osmium and iridium harden pen nibs and spark plug tips.
• Molybdenum, vanadium, manganese and cobalt are added in small amounts to toughen tool steels and magnets.

Electronic configuration of the D-block elements

The general electronic configuration of the d-block elements is [noble gas] (n−1)d^1–10 ns^1–2, where n is the outermost shell. The noble gas core is the configuration of the noble gas that ends the period before. Because the electrons being added go into an inner d subshell, the outer valence shell barely changes from one element to the next, which is why neighbours in the transition metal series look so alike.

These d orbitals hold up to ten electrons and drive almost every property in the rest of this guide. The table covers the first row transition metals, whose valence electrons sit in 3d and 4s.

One detail catches many students: when first-row transition elements form ions, the 4s electrons leave before the 3d ones. Fe²⁺ is therefore [Ar] 3d⁶, not [Ar] 3d⁴ 4s², a result that makes sense once quantum mechanics shows the 3d level dropping below 4s as nuclear charge builds.

Electron configuration exceptions in chromium and copper

The Madelung rule predicts 3d⁴ 4s² for chromium and 3d⁹ 4s² for copper, but the measured electron configuration is 3d⁵ 4s¹ and 3d¹⁰ 4s¹. Moving one electron gives a half-filled or completely filled d shell, which lowers the total energy through exchange stabilisation. The gap between the two arrangements is small, so the exceptions rarely change how the elements react.

Properties of transition metals

Nearly every property in this family traces back to delocalised d electrons. Strong metallic bonding holds the atoms together, so cohesion rises as more electrons are shared toward the middle of each row and falls away again at group 12.

Physical properties: density, hardness and melting points

High density

Osmium is the densest naturally occurring element at about 22.6 g/cm³, and most members outweigh aluminium several times over.

Melting points and boiling points

Mostly high; tungsten melts at 3,422 °C, the highest of any metal, while the group 12 trio sits far lower.

Hardness

Chromium is among the hardest pure metals, while copper, silver and gold stay soft and malleable.

Conduction

Good electrical and thermal conductivity, with silver the best conductor of any element.

Appearance

Lustrous and mostly silvery-grey, with copper and gold the coloured exceptions.

Chemical properties and reactivity

Reactivity varies widely: iron rusts in damp air while gold and platinum barely react at all. In inorganic chemistry terms, transition metals share moderate electronegativity (roughly 1.3 to 2.5 on the Pauling scale), relatively low first ionization energy compared with nonmetals, and a weak electron affinity.

They lose electrons to form cations, and with nonmetals such as oxygen or chlorine they build ionic compounds, although higher charges make the bonding more covalent. Each oxidation state a metal adopts brings its own chemistry, as the next section shows.

Oxidation states of the transition elements

The trait that most clearly separates these elements from main-group metals is their variable oxidation states. Because the 4s and 3d energy levels sit close together, an atom can give up different numbers of electrons for a similar energy cost. Most members therefore show multiple oxidation states that usually differ by one, whereas p-block elements tend to jump by two.

• The highest oxidation state across the first row equals the total of 3d and 4s electrons, from titanium (+4) up to manganese (+7), then it falls away toward zinc.
• Ruthenium and osmium reach +8, as in OsO₄, and iridium has been observed at +9.
• The lowest states, zero or negative, appear in metal carbonyl complexes such as Cr(CO)₆, which follow the 18-electron rule.

Ionic compounds mostly hold transition metal ions in the +2 and +3 states. Each change of charge also shifts colour and magnetic behaviour, which is why the same metal can produce several very different-looking salts.

Coloured compounds of the transition series

Ask anyone who has studied chemistry to name one feature of the transition series and they'll probably say colour. Most transition metal compounds absorb part of visible light and show you the complementary colour. Classic laboratory examples include copper sulfate (blue), nickel chloride (green), cobalt nitrate (red), potassium dichromate (orange) and potassium permanganate (deep purple).

Within transition metal complexes, the metal is bound to ligands such as water, ammonia or chloride, forming coordination complexes. Swapping the ligand changes the colour, so dilute copper(II) solution turns from pale blue to deep blue when ammonia is added. Intense colours like permanganate and dichromate come from charge transfer, where an electron moves between ligand and metal, rather than from d-d jumps.

Crystal field splitting and d-d transitions

Crystal field theory explains the rest. In an octahedral complex the five d orbitals split into a lower set of three and an upper set of two; in a tetrahedral complex the pattern flips and the gap is smaller. When light of matching energy lifts an electron across that gap, the crystal field splitting decides which colour is absorbed.

A large gap favours low spin arrangements with electrons paired up, while a small gap gives high spin arrangements with as many unpaired electrons as possible. Ligand field theory refines the model by allowing for covalent bonding to the ligands.

Magnetism and unpaired d electrons

Magnetism follows directly from the electron count. Any ion with unpaired d electrons is paramagnetic and is drawn into a magnetic field, while low-spin d⁶ and square-planar d⁸ complexes pair every electron and are diamagnetic. You can estimate the effect with the spin-only formula, where n is the number of unpaired electrons:

$$\mu_{s} = \sqrt{n(n+2)}\ \text{BM}$$

A high-spin Mn²⁺ ion (d⁵, n = 5) gives \(\sqrt{35} \approx 5.92\) BM, while Ni²⁺ (d⁸, n = 2) gives \(\sqrt{8} \approx 2.83\) BM. When neighbouring atomic moments line up in a solid, the material becomes ferromagnetic, which is why iron, cobalt and nickel are the only elements that stay magnetic at room temperature.

Transition element catalysts and industrial uses

A transition element can switch between oxidation states and offer empty d orbitals, so it can hold reactants on its surface, weaken their bonds and release the product unchanged. That makes transition metals some of the most important catalysts in industry and a central subject in inorganic chemistry.

Catalytic properties in industry

Their catalytic properties drive several of the largest chemical processes in the world:

1. Haber process: finely divided iron combines nitrogen and hydrogen into ammonia for fertiliser.
2. Contact process: vanadium(V) oxide oxidises sulfur dioxide on the way to sulfuric acid.
3. Catalytic hydrogenation: nickel adds hydrogen to vegetable oils and other unsaturated compounds.
4. Car catalytic converters use platinum, palladium and rhodium to clean exhaust gases.

Homogeneous catalysis uses dissolved complexes instead, such as palladium in carbon-carbon coupling reactions; each catalyst still works by cycling the metal between oxidation states.