Valence in chemistry is the number of chemical bonds an atom can form. Carbon has a valence of four, meaning it forms four bonds. Hydrogen and fluorine have a valence of one. This single concept explains why water is H₂O and not HO or H₃O: oxygen has a valence of two, so it bonds with exactly two hydrogen atoms.
How Valence Works at the Atomic Level
Every atom has electrons arranged in layers, or “shells,” around its nucleus. The electrons in the outermost shell are called valence electrons, and they’re the ones available for bonding. An atom forms bonds by sharing or transferring these outer electrons with other atoms, and it keeps doing so until it reaches a stable arrangement, typically eight electrons in its outer shell. This tendency is known as the octet rule.
The simplest way to calculate an atom’s valence is to look at how many of its outer electrons it uses in bonding. Mathematically, you subtract the number of electrons that sit unused (nonbonding) on the atom from the total number of valence electrons the free atom starts with. For oxygen, that’s 6 valence electrons minus 4 nonbonding electrons (the two lone pairs), giving a valence of 2.
Common Valences for Everyday Elements
The most important valences to know are for the elements that make up most of the molecules you’ll encounter:
- Hydrogen: 1 (forms one bond)
- Oxygen: 2 (forms two bonds)
- Nitrogen: 3 (forms three bonds)
- Carbon: 4 (forms four bonds)
- Fluorine and chlorine: 1 (form one bond each)
- Boron: 3 (forms three bonds)
These numbers directly explain the formulas of familiar compounds. Methane is CH₄ because carbon needs four bonds and each hydrogen provides one. Ammonia is NH₃ because nitrogen needs three. Carbon dioxide is CO₂ because carbon’s four bonds are split between two oxygen atoms, each forming a double bond.
Finding Valence on the Periodic Table
You don’t need to memorize individual valences for every element. The periodic table organizes elements into groups that share the same number of valence electrons:
- Group 1 (lithium, sodium, potassium): 1 valence electron
- Group 2 (beryllium, magnesium, calcium): 2 valence electrons
- Group 13 (boron, aluminum): 3 valence electrons
- Group 14 (carbon, silicon): 4 valence electrons
- Group 15 (nitrogen, phosphorus): 5 valence electrons
- Group 16 (oxygen, sulfur): 6 valence electrons
- Group 17 (fluorine, chlorine): 7 valence electrons
- Group 18 (noble gases): 8 valence electrons (except helium, which has 2)
Note that the number of valence electrons isn’t always the same as the valence itself. Oxygen has 6 valence electrons but a valence of 2, because it only needs two more electrons to complete its octet. Elements in Group 1 have 1 valence electron and a valence of 1. Elements in Group 17 have 7 valence electrons but also a valence of 1, since they need just one more electron to reach eight. The pattern: for groups 1 through 4, the valence equals the group’s electron count. For groups 5 through 7, the valence equals 8 minus that count.
Using Valence to Predict Chemical Formulas
One of the most practical uses of valence is figuring out the formula of a compound before you ever step into a lab. The principle is simple: in a stable compound, the total bonding capacity of all atoms must balance out.
Take magnesium (valence 2) and chlorine (valence 1). Magnesium needs two bonds, but each chlorine atom can only provide one. So you need two chlorine atoms per magnesium atom, giving MgCl₂. For aluminum oxide, aluminum has a valence of 3 and oxygen has a valence of 2. To balance these, you need two aluminum atoms (providing 6 total bonds) and three oxygen atoms (also needing 6 total bonds), yielding Al₂O₃. This “criss-cross” method, where you swap the valences to get subscripts, works reliably for simple compounds.
Variable Valence in Transition Metals
Not every element sticks to a single valence. Transition metals, the elements in the middle block of the periodic table, routinely display multiple valences. Iron, for instance, commonly shows valences of 2 and 3. Chromium can be 2 or 3. Cobalt is usually 2 or 3 but can reach 4 or even 5 in certain compounds.
This happens because transition metals have partially filled inner electron shells (specifically, the d orbitals) that can contribute different numbers of electrons to bonding depending on what they’re reacting with. The type of atom or molecule bonding to the metal influences how many electrons the metal makes available. This is why iron can form both FeO (using a valence of 2) and Fe₂O₃ (using a valence of 3), two compounds with very different properties. Rust is the latter.
Expanded Valence: Breaking the Octet Rule
The octet rule works well for small, light elements, but atoms in the third row of the periodic table and beyond can exceed eight electrons in their outer shell. Phosphorus, for example, typically has a valence of 3 but can expand to 5, as in phosphorus pentachloride (PCl₅), where it bonds with five chlorine atoms. Sulfur can expand to a valence of 6, forming sulfur hexafluoride (SF₆). Xenon, a noble gas that “shouldn’t” bond at all, forms XeF₄ with four bonds to fluorine.
These expanded valences are possible because larger atoms have access to additional orbital space (vacant d orbitals) that smaller atoms like carbon and nitrogen lack. This is why you’ll never see carbon form six bonds, but sulfur can.
Valence vs. Oxidation State
Valence and oxidation state are related but not the same thing, and mixing them up is one of the most common sources of confusion in chemistry. Valence is always a positive number (or zero). It simply counts how many bonds an atom forms. Oxidation state, on the other hand, can be positive or negative because it tracks the hypothetical charge an atom would carry if every bond were completely broken and all shared electrons were handed to the more electronegative atom.
In water, oxygen has a valence of 2 (it forms two bonds) and an oxidation state of −2 (it “claims” both shared electron pairs because it’s more electronegative than hydrogen). Hydrogen has a valence of 1 and an oxidation state of +1. The valence tells you about bonding structure. The oxidation state tells you about how electron density is distributed. You need to draw an actual molecular structure to determine valence, while oxidation state can often be assigned from a chemical formula alone.
For simple compounds of main group elements, the absolute value of the oxidation state often matches the valence, which is why the two get conflated. But for transition metals and more complex molecules, they frequently diverge, and keeping them separate will save you a lot of confusion.