19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds

Transition Metals and Their Compounds

Properties of transition metals

Transition metals are elements in the d-block of the periodic table that have partially filled d orbitals in their neutral or cationic states. This partially filled d subshell is what sets them apart from main-group metals and gives rise to most of their distinctive behavior.

Common physical properties of transition metals include:

• High melting and boiling points due to strong metallic bonding (tungsten has the highest melting point of any metal: 3422°C)
• High densities from close packing of atoms (osmium is the densest naturally occurring element at 22.59 g/cm³)
• High tensile strengths from strong metallic bonding (steel ranges from 400–2000 MPa depending on composition)
• Ductility and malleability, allowing them to be drawn into wires or hammered into sheets (copper wires, gold foil)
• Good electrical and thermal conductivity due to delocalized electrons (silver has the highest electrical conductivity of any element: 6.30 × 10⁷ S/m)

Variable oxidation states are one of the most important chemical features of transition metals. Because the energy differences between their d orbitals are small, these metals can lose different numbers of electrons and form stable compounds in multiple oxidation states. Manganese, for example, can exist as +2, +3, +4, +6, or +7 in different compounds.

Colored compounds form because unpaired d electrons can undergo d-d electronic transitions, absorbing specific wavelengths of visible light. The color you see is the complement of the wavelength absorbed. Copper(II) sulfate appears blue, for instance, because its d-d transitions absorb light in the orange-red region.

Coordination compounds form when transition metal ions act as Lewis acids, accepting electron pairs from ligands (ions or molecules that donate electron pairs) through coordinate covalent bonds. Hemoglobin is a biological example: an iron ion is complexed with a porphyrin ligand, enabling oxygen transport in your blood.

Extraction of iron, copper, and silver

These extraction processes are examples of metallurgy, the science of obtaining metals from their ores and refining them for use.

Iron is extracted from hematite ($Fe_2O_3$):

1. Concentrate the ore through froth flotation and magnetic separation
2. Reduce the concentrated ore with coke (carbon) in a blast furnace at high temperatures. Carbon monoxide acts as the actual reducing agent: $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$
3. Collect molten iron at the bottom of the furnace for further refining into steel

Copper is extracted from the sulfide ore chalcopyrite ($CuFeS_2$):

1. Concentrate the ore through froth flotation
2. Roast the concentrated ore to convert sulfides to oxides: $2CuFeS_2 + 4O_2 \rightarrow Cu_2S + 2FeO + 3SO_2$
3. Smelt the roasted ore in a furnace with silica ($SiO_2$) to remove iron as slag
4. Purify the crude copper through electrolytic refining, where impure copper serves as the anode and pure copper deposits on the cathode

Silver is often extracted as a byproduct of lead, zinc, or copper mining from the ore argentite ($Ag_2S$):

1. Crush the ore and treat it with a sodium cyanide ($NaCN$) solution to form a soluble silver-cyanide complex: $Ag_2S + 4NaCN \rightarrow 2Na[Ag(CN)_2] + Na_2S$
2. Recover silver from the solution by displacement with zinc or through electrolysis

Oxidation states in transition metals

The oxidation state of a transition metal directly affects the character of the compounds it forms. This shows up clearly in halides, oxides, and salts.

Halides shift from ionic to covalent as oxidation state increases:

• Lower oxidation states tend to form ionic halides ($FeCl_2$, iron(II) chloride)
• Higher oxidation states tend to form covalent halides ($FeCl_3$, iron(III) chloride, which actually exists as a dimer in the gas phase)

Oxides shift from basic to acidic as oxidation state increases:

• Lower oxidation states form basic oxides ($Cr_2O_3$, chromium(III) oxide)
• Higher oxidation states form acidic oxides ($CrO_3$, chromium(VI) oxide)

Salts vary in stability depending on the metal's oxidation state and the nature of the anion:

• $MnSO_4$ (manganese(II) sulfate) is stable
• $Mn_2(SO_4)_7$ (manganese(VII) sulfate) would be highly unstable because Mn(VII) is such a strong oxidizing agent that it tends to oxidize the sulfate anion rather than form a stable salt

The general trend: as oxidation state increases, compounds become more covalent, oxides become more acidic, and the metal becomes a stronger oxidizing agent.

Applications of transition metals

Transition metals are heavily used as catalysts in industrial chemistry. Their variable oxidation states let them cycle between different electron configurations during a reaction, lowering activation energy without being consumed. Iron catalyzes the Haber process for ammonia synthesis, and vanadium(V) oxide catalyzes the Contact process for sulfuric acid production.

In biological systems, transition metal complexes play essential roles. The iron-containing heme group in hemoglobin binds and releases oxygen for transport through the bloodstream. Cobalt sits at the center of vitamin B12, and zinc is critical for the function of many enzymes. These biological roles depend on the same coordination chemistry that defines transition metal behavior in the lab.