Transition Metals: How d Orbitals Create Color and Catalysis

The broad central section of the periodic table, groups 3 to 12, holds the elements usually called the transition metals. Steel bridges, copper wiring, gold rings and titanium hip replacements all come from here; most of what people picture when they hear the word "metal" lives in this block. What these elements share is that their atoms are in the middle of filling a set of d orbitals.

What counts as a transition metal

IUPAC defines a transition element as one whose atom has an incomplete d subshell, or which can form cations with an incomplete d subshell. Applied strictly, that excludes zinc, cadmium and mercury in group 12, whose d subshells are full with ten electrons in both the atom and its common ions. For that reason some textbooks distinguish between "d-block elements" (all of groups 3 to 12) and "transition metals" (the narrower set). This guide covers the whole block for convenience and points out the difference where it matters.

The name dates from the early days of the periodic table, when these elements were seen as a bridge, a transition, between the main-group elements on either side.

How the d orbitals fill

In period 4, electrons first fill the 4s orbital and then, starting with scandium, fill the 3d orbitals of the shell beneath. Scandium is [Ar]3d¹4s², iron is [Ar]3d⁶4s² and zinc is [Ar]3d¹⁰4s². Period 5 fills 4d, and periods 6 and 7 fill 5d and 6d.

There are five d orbitals holding at most ten electrons, which is why each period's d-block contains ten elements. The order is reversed when ions form: a transition metal loses its outer 4s electrons before any 3d electrons. So Fe²⁺ is [Ar]3d⁶, not [Ar]3d⁴4s².

A few elements, notably chromium ([Ar]3d⁵4s¹) and copper ([Ar]3d¹⁰4s¹), have configurations that break the simple filling order. The usual explanation is that a half-filled or completely filled d subshell is energetically favorable. The electron configuration guide covers these exceptions in detail.

Many oxidation states

Group 1 and 2 metals almost always show a single oxidation state, +1 or +2. Transition metals show several. Because the 4s and 3d energies are close, the number of electrons that can be removed varies with the chemical surroundings.

Manganese is a classic example, with common oxidation states running from +2 to +7.

Iron moves between +2 (iron(II)) and +3 (iron(III)). Both the transport of oxygen by hemoglobin and the rusting of steel depend on changes in iron's oxidation state. This ability to switch oxidation states easily is also the foundation of catalytic activity, described below.

Complexes and color

Transition metal ions readily form coordination complexes, surrounding themselves with molecules or ions that carry lone pairs of electrons, such as water, ammonia or chloride. These surrounding species are called ligands. Each ligand donates a pair of electrons to the metal ion, forming a coordinate bond. The Swiss chemist Alfred Werner worked out the structure of such compounds and received the 1913 Nobel Prize in Chemistry.

The color of complexes comes from splitting of the d orbitals. In an isolated atom the five d orbitals have the same energy. When ligands approach, the orbitals that point toward the ligands are pushed up in energy and the rest drop, so the set divides into higher and lower groups. An electron in a lower orbital can absorb visible light of a particular energy and jump to a higher one. We see the light that is left over, the complementary color.

• The blue of copper(II) sulfate solution: the copper ion absorbs orange and red light.
• The red of ruby: traces of chromium ions in an aluminium oxide crystal.
• The green of emerald: again chromium ions, and in some stones vanadium.

Changing the ligands changes the size of the energy gap, and so the color. Adding ammonia to blue copper sulfate solution turns it a deep royal blue. Ions with empty d orbitals (Sc³⁺, Ti⁴⁺) or completely filled ones (Zn²⁺) cannot undergo these transitions and are usually colorless, which is why zinc compounds are typically white.

Transition metals as catalysts

Many of the most important industrial catalysts are transition metals or their compounds. They can exchange electrons by switching oxidation state, and their surfaces can adsorb molecules and weaken their bonds.

• Iron: the Haber–Bosch process, which makes ammonia from nitrogen and hydrogen. Fertilizer produced this way supports a large share of the world's food supply.
• Vanadium: vanadium(V) oxide catalyzes the contact process for making sulfuric acid.
• Nickel: hydrogenation of vegetable oils, used to make margarine and other solid fats.
• Platinum, palladium and rhodium: catalytic converters in car exhausts, turning carbon monoxide, unburned hydrocarbons and nitrogen oxides into less harmful gases.
• Cobalt and rhodium: a wide range of homogeneous catalysts in organic synthesis.

Living things use transition metals as catalysts too. Iron, copper, zinc and molybdenum at the active sites of enzymes drive key reactions in respiration, photosynthesis and nitrogen fixation.

Notable transition metals

• Iron (Fe): the fourth most abundant element in Earth's crust and the main component of steel. Few materials have shaped civilization as much.
• Copper (Cu): the second-best electrical conductor after silver, used in wiring, circuit boards, bronze and brass.
• Titanium (Ti): light, strong and corrosion-resistant. Used in aircraft, medical implants and, as titanium dioxide, in white pigment.
• Chromium (Cr): added to steel to make stainless steel, where it forms a thin protective oxide layer. Also familiar as shiny chrome plating.
• Nickel (Ni) and cobalt (Co): battery cathodes, heat-resistant alloys and magnets.
• Silver (Ag), gold (Au), platinum (Pt): the precious metals. Gold hardly corrodes at all, which makes it valuable for jewelry and electrical contacts.
• Tungsten (W): the highest melting point of any metal, once used for incandescent bulb filaments and still used in cutting tools.
• Mercury (Hg): the only metal that is liquid at room temperature. Because of its toxicity it is being phased out of thermometers and many other products.

Magnetism and physical properties

Iron, cobalt and nickel are ferromagnetic at room temperature, because the spins of their unpaired d electrons line up in the same direction across whole regions of the metal. Most transition metals are dense, hard and high-melting, which is attributed to d electrons taking part in metallic bonding. Osmium and iridium are the densest elements: a cube of either measuring 10 cm on a side would weigh well over 20 kg.

Transition metals in the body, and safety

The human body needs several transition metals. Iron binds oxygen in hemoglobin and myoglobin, zinc is part of the structure and function of hundreds of enzymes, copper supports energy and iron metabolism, and cobalt sits at the center of vitamin B12. These are trace elements: too little causes problems, and so does too much.

Other transition metals are harmful even in small amounts, including cadmium, mercury and chromium(VI) compounds. Chromium is a good illustration of how much oxidation state matters. Chromium(III) is comparatively benign, while chromium(VI) compounds are known to be toxic and carcinogenic and are tightly regulated.

The d-block is the largest region of the table, and its elements are correspondingly varied. Click through the d-block elements and compare their oxidation states, melting points and densities to see how the properties shift across each period.