IB Chemistry SLTopic 3 — Classifying the ElementsPaper 1 & 2Core skill~11 min read
Oxidation States
An oxidation state is a piece of accounting. It is the charge an atom would have if every bond in the compound were completely ionic — a pretence, but an extremely useful one. It lets you see at a glance which atoms have gained electrons and which have lost them, even in reactions where nothing looks like an ion at all.
📚 What you need to know
An oxidation state is the charge an atom would carry if all bonding were fully ionic.
An uncombined element is always 0; a simple ion equals its charge.
In a compound the oxidation states sum to zero; in a polyatomic ion they sum to the ion’s charge.
Fixed values worth learning: Group 1 = +1, Group 2 = +2, F = –1, H = +1 (except –1 in metal hydrides), O = –2 (except in peroxides).
Oxidation = oxidation state increases (electrons lost). Reduction = it decreases (electrons gained).
Stock notation puts the oxidation state in Roman numerals after the metal, e.g. iron(III) oxide.
What it actually means
Take water. The O–H bonds are covalent — the electrons are shared, not transferred. But oxygen is much more electronegative than hydrogen, so it takes the larger share. Now pretend the sharing was total: oxygen would have grabbed both bonding pairs, giving it a charge of –2, and each hydrogen would be left at +1.
Those are the oxidation states: O = –2, H = +1. Nobody is claiming water contains ions. It is a deliberate simplification that makes electron movement visible.
The rule underneath every fixed value is the same: in any bond, the more electronegative atom is assigned the negative oxidation state. Everything else follows from that.
The rules
Rule
Example
An uncombined element is 0
Na, O2, P4, Cl2 → all 0
A monatomic ion equals its charge
Mg2+ = +2, S2– = –2
Group 1 = +1, Group 2 = +2, F = –1
In KBr, K = +1
Hydrogen = +1, but –1 in metal hydrides
HCl: H = +1; NaH: H = –1
Oxygen = –2, but –1 in peroxides and +2 in OF2
H2O: O = –2; H2O2: O = –1
A neutral compound sums to 0
NaCl: (+1) + (–1) = 0
A polyatomic ion sums to its charge
NO3–: sums to –1
The more electronegative atom takes the negative value
OF2: F = –1, so O = +2
Notice the pattern in the exceptions. Hydrogen goes negative only when bonded to something less electronegative (a metal). Oxygen goes positive only when bonded to fluorine, the one element that beats it.
Working out an unknown
You are almost never asked for a value you have memorised. You are asked for the one that is left over — and there is a reliable method.
🧩 Finding an unknown oxidation state
Write down the fixed values first — usually oxygen at –2 and hydrogen at +1.
Multiply each by the number of those atoms in the formula.
Set the total equal to 0 for a neutral compound, or to the charge for an ion.
Solve for the unknown. Divide by the number of those atoms if there is more than one.
Write it with the sign in front: +6, not 6.
WORKED EXAMPLE
Deduce the oxidation state of nitrogen in the nitrate ion, NO3–.
Fixed value: oxygen is −2, and there are 33 × (−2) = −6The total must equal the ion’s charge, −1N + (−6) = −1N = +5
WORKED EXAMPLE
Deduce the oxidation state of chromium in the dichromate ion, Cr2O72–.
Oxygen is −2, and there are 77 × (−2) = −14Total must equal −22Cr + (−14) = −2, so 2Cr = +12Cr = +6Two chromium atoms share the +12, so divide — a very common slip.
WORKED EXAMPLE
Deduce the oxidation state of oxygen in hydrogen peroxide, H2O2.
Hydrogen is +1, and there are 22 × (+1) = +2Neutral compound, so the total is 02O + (+2) = 0, so 2O = −2O = −1This is the peroxide exception — and here the maths proves it rather than you having to remember it.
Spotting oxidation and reduction
Once you can assign oxidation states, redox becomes obvious. Compare each element before and after:
Direction along the line is the whole test: right is oxidation, left is reduction, whatever the sign of the numbers.
Oxidation state goes up → electrons lost → that element has been oxidised.
Oxidation state goes down → electrons gained → that element has been reduced.
WORKED EXAMPLE
In Cl2 + 2KBr → 2KCl + Br2, identify what has been oxidised and what has been reduced.
Chlorine: element → chloride0 → −1, a decreasechlorine is reducedBromine: bromide → element−1 → 0, an increasebromide is oxidisedPotassium stays at +1 throughout — a spectator ion.
Naming with oxidation states
Transition metals can have more than one oxidation state, so the name has to say which one. That is what the Roman numerals in Stock notation are for.
Formula
Metal’s oxidation state
Name
FeO
+2
iron(II) oxide
Fe2O3
+3
iron(III) oxide
Cu2O
+1
copper(I) oxide
KMnO4
+7
potassium manganate(VII)
K2Cr2O7
+6
potassium dichromate(VI)
The Roman numeral is always positive and has no sign, and it refers to that one element, not to the whole compound. Non-metals are usually named with prefixes instead — SO2 is sulfur dioxide, not sulfur(IV) oxide.
Can oxidation states be fractions?
Occasionally the maths gives you something like +2.5. That does not mean half an electron has moved. It means the atoms of that element are in different environments within the ion, and the value you calculated is the average across them. A single atom always has a whole-number oxidation state.
💡 Exam tip
Always write the sign first: +2, –1. A bare “2” can be read as a charge and lose the mark.
Divide when there is more than one atom of the unknown element. This is the single most common arithmetic error here.
To prove redox, quote the oxidation state before and after for each element you mention.
Oxidation state and ionic charge are written differently: +2 for the oxidation state, 2+ for the charge on an ion.
⚠️ Common mix-up
An oxidation state is not a real charge. The chlorine in HCl is not an ion, but its oxidation state is still –1.
Oxidation is an increase, even when the number stays negative — –2 to –1 is oxidation.
Hydrogen is not always +1 (metal hydrides) and oxygen is not always –2 (peroxides, OF2).
The Roman numeral in a name is never negative and never carries a sign.
Don’t forget to multiply by the subscript before balancing the total.
That completes Classifying the Elements: The Periodic Table. You started with a grid of 118 boxes and finished able to predict an element’s size, reactivity, oxide and electron bookkeeping from its position alone — which is exactly what the table was built to let you do.
Want this explained one-to-one?
Book a free session with an experienced IB Chemistry tutor and get your trickiest topics made simple.
