Transition metals are the primary group of elements that form cations with varying positive charges. Iron, copper, cobalt, chromium, manganese, and many others in the d-block of the periodic table can each lose different numbers of electrons, producing ions with two or more possible charges. A smaller group of post-transition metals, notably tin and lead, also exhibit this behavior.
Why Some Metals Have Variable Charges
Metals on the left side of the periodic table, like sodium (+1) and calcium (+2), almost always form a single type of cation. Their outermost electrons are easy to remove, and the energy cost jumps sharply after those electrons are gone. Transition metals are different. They have electrons in two nearby energy levels (the outermost s-orbital and the d-orbitals just beneath it), and removing varying numbers of those electrons doesn’t require dramatically different amounts of energy. That makes it relatively easy for a single metal to exist as ions with different charges depending on the chemical environment.
For post-transition metals like tin and lead, a different mechanism is at work. These heavier elements can hold onto their outermost pair of s-electrons unusually tightly, a phenomenon chemists call the inert pair effect. This means tin can lose either two or four electrons, giving Sn²⁺ or Sn⁴⁺, and lead can form Pb²⁺ or Pb⁴⁺.
The Most Common Multivalent Metals
Here are the metals you’ll encounter most often with variable charges, along with their typical ionic forms:
- Iron (Fe): Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)). Iron loses its two outermost electrons to form Fe²⁺, or a third electron from its d-orbitals to reach the especially stable half-filled d⁵ configuration of Fe³⁺.
- Copper (Cu): Cu⁺ (copper(I)) and Cu²⁺ (copper(II)). Copper(II) is the more common form in everyday compounds like copper(II) nitrate.
- Manganese (Mn): Ranges from Mn²⁺ all the way up to Mn⁷⁺. The +2 state is the most stable as a free ion, but manganese appears as +4 in manganese dioxide (MnO₂) and +7 in the deep purple permanganate ion (MnO₄⁻). Manganese has the widest range of oxidation states in the first row of transition metals because it has five unpaired d-electrons that can all be removed.
- Chromium (Cr): Cr²⁺, Cr³⁺, and Cr⁶⁺. The +3 state is the most stable. At the +6 level, chromium doesn’t exist as a simple ion in water but instead combines with oxygen to form chromate (CrO₄²⁻) or dichromate (Cr₂O₇²⁻) ions.
- Cobalt (Co): Co²⁺ (cobalt(II)) and Co³⁺ (cobalt(III)).
- Titanium (Ti): Ti²⁺, Ti³⁺, and Ti⁴⁺.
- Vanadium (V): V²⁺, V³⁺, V⁴⁺, and V⁵⁺. Above +3, vanadium bonds with oxygen in solution rather than existing as a bare ion.
- Tin (Sn): Sn²⁺ (tin(II)) and Sn⁴⁺ (tin(IV)).
- Lead (Pb): Pb²⁺ (lead(II)) and Pb⁴⁺ (lead(IV)). The +2 state is far more common and stable.
D-Block Metals That Don’t Vary
Not every transition metal behaves this way. Scandium essentially only forms Sc³⁺, and zinc only forms Zn²⁺. These are the two exceptions in the first row of the d-block. Scandium has just one d-electron beyond its argon core, so removing all three valence electrons to reach a noble-gas configuration is strongly favored. Zinc sits at the opposite end with a completely filled d¹⁰ shell, making it energetically costly to pull electrons from that stable arrangement. Its two s-electrons come off easily, but the d-electrons stay put.
A Trend Across the Periodic Table
The number of possible oxidation states generally increases as you move toward the middle of the transition metal rows, then decreases again. Manganese, sitting near the center of the first transition row, has the most oxidation states (seven). Elements at the edges, like scandium and zinc, have the fewest. This pattern reflects the number of d-electrons available: metals in the middle have enough unpaired d-electrons to lose several but haven’t filled the d-shell to the point where those electrons become difficult to remove.
Moving down a column in the periodic table also changes which charges are most stable. Chromium (+3 is most stable) sits above molybdenum and tungsten, which are both more stable in the +4 and +5 states. Heavier transition metals can access higher oxidation states more easily because their d-electrons are held less tightly by the nucleus.
How Charges Are Shown in Chemical Names
Because these metals can form more than one cation, their names need to specify which one you mean. The modern system (called Stock notation) places a Roman numeral in parentheses after the metal’s name to indicate the charge. FeCl₂ is iron(II) chloride, meaning iron carries a +2 charge. FeCl₃ is iron(III) chloride. SnF₄ is tin(IV) fluoride.
An older naming system used Latin-root suffixes instead: “-ous” for the lower charge and “-ic” for the higher one. Under that system, Fe²⁺ was ferrous and Fe³⁺ was ferric; Sn²⁺ was stannous and Sn⁴⁺ was stannic. You’ll still see these names on older reagent bottles and in some industrial contexts, but Roman numerals are the accepted standard.
Why Variable Charges Matter in Biology
The ability to switch between charges is not just a chemistry-class curiosity. It’s the reason iron and copper are essential to life. Iron is a critical component of the enzymes and protein complexes that carry out cellular respiration, the process your cells use to convert food into energy. It works by cycling between Fe²⁺ and Fe³⁺, picking up and releasing electrons along the way.
Copper plays a similar role. It sits at the heart of several proteins that shuttle electrons during essential reactions, including the final step of the electron transport chain (cytochrome c oxidase). Copper toggles between its +1 and +2 states throughout each catalytic cycle. Without metals capable of changing charge, the electron-transfer chemistry that powers nearly all living organisms would not be possible.