S3.1.6 Oxidation states

Oxidation States: Decoding Electron Behavior in Compounds

What is an Oxidation State?

• In simpler terms, it tells us how many electrons an atom has gained, lost, or shared during bonding.

• Oxidation states are vital for understanding redox reactions, where electrons are transferred between substances.
• They help us identify which species is oxidized (loses electrons) and which is reduced (gains electrons).

Conventions for Assigning Oxidation States

1. To assign oxidation states, chemists follow a systematic set of rules.
2. These rules ensure consistency and accuracy across different compounds.

Rule 1: Free Elements

The oxidation state of an atom in its elemental form is always 0. This applies to:

• Single atoms like Na, O₂, or Cl₂.
• Diatomic molecules like H₂ or N₂.

Rule 2: Sum of Oxidation States

The sum of the oxidation states of all atoms in:

• A neutral compound is 0.
• A polyatomic ion equals the charge of the ion.

Rule 3: Group-Specific Rules

Certain elements have predictable oxidation states in most of their compounds:

• Fluorine is always −1(it’s the most electronegative element).
• Group 1 metals (e.g., Na, K) are always +1.
• Group 2 metals (e.g., Mg, Ca) are always +2.
• Hydrogen is usually +1, except in metal hydrides (e.g., NaH), where it’s −1.
• Oxygen is usually −2, except in peroxides (e.g., H₂O₂, where it’s −1) or when bonded to fluorine (e.g., OF₂, where it’s +2).

Examples of Oxidation States in Common Compounds

Let’s apply these rules to deduce the oxidation states of atoms in some common compounds.

Example 1: Water (H₂O)

1. Hydrogen is +1 (Rule 3).
2. Oxygen is −2 (Rule 3).
3. Total: $2(+1)+(-2)=0$.

Example 2: Hydrogen Peroxide (H₂O₂)

1. Hydrogen is+1 (Rule 3).
2. Oxygen is −1 (except for peroxides).
3. Total: $2(+1)+2(-1)=0$.

Example 3: Nitrate Ion (NO₃⁻)

1. Oxygen is−2 (Rule 3).
2. Let the oxidation state of nitrogen be $x$.
3. Total: $x+3(-2)=-1$.
4. Solving for $x$: $x=+5$.

Example 4: Sulfate Ion (SO₄²⁻)

1. Oxygen is −2 (Rule 3).
2. Let the oxidation state of sulfur be $x$.
3. Total: $x+4(-2)=-2$.
4. Solving for $x$: $x=+6$.

Special Cases: Hydrogen in Metal Hydrides and Oxygen in Peroxides

Sometimes, the oxidation states of hydrogen and oxygen deviate from their usual values. These exceptions are worth noting:

1. Hydrogen in Metal Hydrides: When hydrogen is bonded to a metal (e.g., NaH), it acts as the more electronegative element and has an oxidation state of −1.
2. Oxygen in Peroxides: In peroxides (e.g., H₂O₂), oxygen has an oxidation state of −1 due to the oxygen-oxygen single bond.

Why Do Oxidation States Matter?

Understanding oxidation states allows us to:

1. Predict Redox Reactions: Oxidation states reveal which atoms are oxidized (increase in oxidation state) and which are reduced (decrease in oxidation state).
2. Name Compounds: Oxidation states are used in naming compounds, particularly for transition metals and polyatomic ions. For example:
   • FeCl₂ is iron(II) chloride (Fe has an oxidation state of+2).
   • FeCl₃ is iron(III) chloride (Fe has an oxidation state of+3).
3. Understand Variable Oxidation States: Transition metals often exhibit multiple oxidation states, which influence their chemical reactivity and the colors of their compounds.

Reflection and Practice

Now that you’ve explored the rules and examples, try assigning oxidation states to the following compounds and ions:

1. $K_{2}C{r}_7{O}_4$
2. $C{H}_4$
3. $C{O}_2$
4. $HCl{O}_4$
5. $O{F}_2$