In chemistry, keeping track of electrons is crucial for understanding how reactions proceed, particularly during oxidation and reduction. Oxidation states, also known as oxidation numbers, provide a book-keeping tool to monitor the electron distribution in atoms, molecules, and ions.
🔑 Key Principle
An oxidation state represents the theoretical charge an atom would carry if all bonds to other atoms were considered completely ionic. It is a vital tool for identifying which species have been oxidised or reduced in a chemical reaction.
What is an Oxidation State?
Unlike actual ionic charges, which describe real physical charges on ions, oxidation states are formal values assigned using a set of international rules. These values tell us whether an atom is in an electron-rich or electron-deficient state relative to its neutral elemental form.
The formal charge that would reside on an atom if all shared covalent bonding electrons were assigned completely to the more electronegative atom in the bond.
A chemical reaction in which both reduction (gain of electrons / decrease in oxidation state) and oxidation (loss of electrons / increase in oxidation state) take place simultaneously.
Rules for Assigning Oxidation States
To determine the oxidation state of any atom in a chemical formula, chemists use a strict hierarchy of rules. These rules must be applied in order, from the highest priority to the lowest.
When stating oxidation states in exams, you must include the sign (positive or negative) before the number, for example: +5 or -2. Writing "5+" or "2-" represents ionic charge, not an oxidation state, and will result in a loss of marks. Additionally, do not forget to include the + sign for positive states.
Using Roman Numerals in Nomenclature
Many elements, especially transition elements and non-metals like sulfur, nitrogen, and chlorine, can exhibit several different oxidation states. In systematic IUPAC naming, Roman numerals are used to specify the oxidation state of these elements in a compound.
Numbers written in parentheses immediately after the name of an element to indicate its oxidation state in that specific compound (e.g., iron(II) indicates Fe with an oxidation state of +2).
Consider the oxides of iron: iron(II) oxide contains \( \text{Fe}^{2+} \) (oxidation state +2), so its formula is \( \text{FeO} \). Iron(III) oxide contains \( \text{Fe}^{3+} \) (oxidation state +3), so its formula is \( \text{Fe}_2\text{O}_3 \).
Similarly, for polyatomic anions, the Roman numeral indicates the oxidation state of the central non-metal:
- Sulfate(IV) ion: \( \text{SO}_3^{2-} \) (sulfur is +4)
- Sulfate(VI) ion: \( \text{SO}_4^{2-} \) (sulfur is +6)
- Nitrate(III) ion: \( \text{NO}_2^{-} \) (nitrogen is +3)
- Nitrate(V) ion: \( \text{NO}_3^{-} \) (nitrogen is +5)
- \( \text{NO}_2 \)
- \( \text{NO}_3^{-} \)
- \( \text{NH}_4^{+} \)
1. \( \text{NO}_2 \): This is a neutral molecule, so the sum of oxidation states is 0. Oxygen has a state of -2. \[ \text{N} + 2(-2) = 0 \implies \text{N} - 4 = 0 \implies \text{N} = +4 \] Systematic name: nitrogen(IV) oxide.
2. \( \text{NO}_3^{-} \): This is a polyatomic ion with a charge of -1, so the sum is -1. \[ \text{N} + 3(-2) = -1 \implies \text{N} - 6 = -1 \implies \text{N} = +5 \] Systematic name: nitrate(V) ion.
3. \( \text{NH}_4^{+} \): This is a polyatomic ion with a charge of +1, so the sum is +1. Hydrogen is +1. \[ \text{N} + 4(+1) = +1 \implies \text{N} + 4 = +1 \implies \text{N} = -3 \] Systematic name: ammonium ion (nitrogen is in the -3 state).
Let the oxidation state of chromium be \( x \). The dichromate ion has a total charge of -2, and oxygen is assigned its standard oxidation state of -2.
Set up the algebraic equation for the sum of states: \[ 2(x) + 7(-2) = -2 \] \[ 2x - 14 = -2 \] \[ 2x = +12 \] \[ x = +6 \] Therefore, the oxidation state of chromium in the dichromate ion is +6 (systematically named as chromate(VI) species).
Be careful when dealing with peroxides and hydrides. In sodium hydride (\( \text{NaH} \)), sodium is a Group 1 metal and must be +1, which forces hydrogen to be -1. In hydrogen peroxide (\( \text{H}_2\text{O}_2 \)), hydrogen is +1, which forces oxygen to be -1. These exceptions are frequently tested to trap unwary students.
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