Fundamentals of Oxidation and Reduction

TL;DR

Oxidation is losing electrons, and reduction is gaining electrons; together these are called redox reactions. We can remember which is which using simple mnemonics like OIL RIG. These reactions are fundamental in chemistry, powering everything from batteries to digestion.

1. The Mental Model

Think of chemical reactions as a game of electron 'hot potato'. Sometimes an atom passes an electron to another (oxidation), and sometimes it catches one (reduction). These electron transfers are always linked – one can't happen without the other.

2. The Core Material

Oxidation and reduction are everywhere in chemistry, from how your body processes food to how batteries produce electricity. You'll often hear them referred to jointly as redox reactions.

What are Oxidation and Reduction?

The most fundamental way to think about oxidation and reduction is in terms of electron transfer.

• Oxidation: This is the loss of electrons. When an atom or ion loses electrons, its charge becomes more positive.
• Reduction: This is the gain of electrons. When an atom or ion gains electrons, its charge becomes more negative.

A highly useful mnemonic to remember this is:
OIL RIG
Oxidation Is Losing (electrons)
Reduction Is Gaining (electrons)

Remember, for every electron lost by one species, another species must gain it. Oxidation and reduction always happen simultaneously. You can't have one without the other. The substance that gets oxidized is called the reducing agent because it causes the other substance to be reduced. The substance that gets reduced is called the oxidizing agent because it causes the other substance to be oxidized.

Oxidation Numbers (or Oxidation States)

Sometimes, electron transfer isn't completely obvious, especially when we're dealing with covalent compounds (where electrons are shared, not fully transferred). To help track electron shifts in these situations, we use oxidation numbers (or oxidation states). Think of these as a way to assign imaginary charges based on electronegativity, even if no full electron transfer occurs.

Here are the basic rules for assigning oxidation numbers:

1. Elemets in their elemental form have an oxidation number of 0.
   • Examples: H₂ (0), O₂ (0), Na (0), Fe (0).
2. The oxidation number of a monatomic ion equals its charge.
   • Examples: Cl⁻ (-1), Na⁺ (+1), Ca²⁺ (+2).
3. In most compounds, fluorine (F) is -1. It's the most electronegative element, so it always pulls electrons toward itself.
4. In most compounds, oxygen (O) is -2.
   • Exceptions: Peroxides (like H₂O₂) where oxygen is -1, and when bonded to fluorine (like OF₂) where oxygen is +2.
5. In most compounds, hydrogen (H) is +1.
   • Exception: When bonded to metals (metal hydrides, like NaH) where hydrogen is -1.
6. The sum of oxidation numbers in a neutral compound is 0.
7. The sum of oxidation numbers in a polyatomic ion equals the charge of the ion.

You can use these rules to figure out the oxidation number of a specific atom in a more complex molecule or ion. If an atom's oxidation number increases, it's been oxidized (lost electrons). If its oxidation number decreases, it's been reduced (gained electrons).

3. Worked Example

Let's look at the reaction between magnesium metal and oxygen gas to form magnesium oxide:

2 Mg (s) + O₂ (g) → 2 MgO (s)

1. Assign oxidation numbers to each element in the reactants:

   • Mg (s): This is an element in its elemental form, so its oxidation number is 0.
   • O₂ (g): This is an element in its elemental form, so its oxidation number is 0.

2. Assign oxidation numbers to each element in the product (MgO):

   • MgO is a neutral compound. Oxygen usually has an oxidation number of -2.
   • For the compound to be neutral, Mg must balance the -2 charge from oxygen. So, Mg will have an oxidation number of +2.
   • Therefore, in MgO, Mg is +2 and O is -2.

3. Identify changes in oxidation numbers:

   • Magnesium (Mg): Went from 0 to +2. Since its oxidation number increased, Mg was oxidized.
   • Oxygen (O): Went from 0 to -2. Since its oxidation number decreased, O was reduced.

This confirms it's a redox reaction. Magnesium is the reducing agent (it got oxidized), and oxygen is the oxidizing agent (it got reduced).

4. Key Takeaways

• Oxidation is the loss of electrons, leading to a more positive oxidation state.
• Reduction is the gain of electrons, leading to a more negative oxidation state.
• Redox reactions always involve both oxidation and reduction occurring simultaneously.
• The oxidizing agent is the species that gets reduced (gains electrons).
• The reducing agent is the species that gets oxidized (loses electrons).
• Oxidation numbers help track electron shifts, especially in covalent compounds, even without full electron transfer.
• An increase in oxidation number means oxidation, while a decrease means reduction.

Common Mistakes to Avoid:
- Don't confuse "oxidizing agent" with "being oxidized" – they are opposites.
- Don't forget that oxidation and reduction are always coupled; one can't happen alone.
- Be careful with exceptions to oxidation number rules, especially for oxygen and hydrogen.
- Never assign oxidation numbers to elements you can't balance out (e.g., trying to assign to only half of a polyatomic ion without considering the overall charge).

5. Now Try It

Determine the oxidation numbers for every atom in the following reaction and identify which species is oxidized and which is reduced:

Fe₂O₃ (s) + 3 CO (g) → 2 Fe (s) + 3 CO₂ (g)

What to do: For each compound (Fe₂O₃, CO, Fe, CO₂), assign the oxidation number to each element present. Then, compare the numbers from reactants to products to state which element was oxidized and which was reduced.

What success looks like: You'll have a list of oxidation numbers for each element in each compound, correctly identifying the element that gained electrons and the element that lost electrons. You'll also correctly name the oxidizing and reducing agents.