4.3 Oxidation Reduction Reactions
PRE-HEALTH PREP
Oxidation-Reduction Reactions
An Introduction to Redox
General Chemistry 4.3
Lesson Overview
Oxidation-reduction (redox) reactions involve the transfer of electrons between chemical species, with oxidation defined as electron loss and reduction as electron gain. A systematic set of rules governs the assignment of oxidation states to atoms in elements, ions, and compounds, enabling identification of which species are oxidized and reduced in a given reaction. The activity series ranks metals by their tendency to lose electrons, providing a straightforward method for predicting whether a single replacement reaction will occur spontaneously.
Table of Contents
What Are Oxidation and Reduction?
Oxidation-reduction reactions — commonly called redox reactions — are fundamentally electron transfer reactions. In every redox reaction, at least one species loses electrons and at least one species gains them. The names "oxidation" and "reduction" each describe one side of this exchange:
- Oxidation is the loss of electrons.
- Reduction is the gain of electrons.
Classic example from the slides:
2Na(s) + Cl2(g) ➔ 2NaCl(s)
Mnemonics: OIL RIG and LEO the Lion Says GER
Two classic mnemonics make these definitions stick:
OIL RIG: Oxidation Is Loss, Reduction Is Gain.
LEO the lion says GER: Lose Electrons Oxidation; Gain Electrons Reduction. This one is worth memorizing — the letter E explicitly reminds you that electrons are what's being lost or gained, which is easy to forget under exam pressure.
In every redox reaction you will encounter at this level, oxidation and reduction always occur together — you cannot have one without the other. One species loses electrons only because another is there to accept them.
How Do Oxidation States Track Electron Transfer?
To track who loses electrons and who gains them, chemists use oxidation states (also called oxidation numbers). Rather than counting every electron directly, you assign each atom a number that represents its effective charge — and then watch for changes. A change in oxidation state is your evidence that electron transfer has occurred.
Example: Zinc and Copper(II) Ion
Consider the reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
- Zinc starts as a neutral element (oxidation state = 0) and becomes Zn²⁺ (oxidation state = +2). It lost two electrons — zinc was oxidized.
- Copper starts as Cu²⁺ (oxidation state = +2) and becomes neutral copper metal (oxidation state = 0). It gained two electrons — copper was reduced.
Why Losing Electrons Raises the Oxidation State
At first glance it seems backwards: losing something should make a number go down, not up. The key is that electrons carry a negative charge. When you lose negative charges, your overall charge becomes more positive — so your oxidation state increases. When you gain negative charges, your overall charge becomes more negative — so your oxidation state decreases. That is precisely why the process of gaining electrons is called reduction: the oxidation state gets reduced (lowered).
Memory check: Zinc goes from 0 to +2 (more positive) — it lost electrons. Copper goes from +2 to 0 (more negative, i.e., reduced) — it gained electrons. If this feels backwards, re-read the logic above until it clicks; this is a concept that trips up nearly every student the first time.
4 Rules for Assigning Oxidation States
Apply these rules in order. As soon as a rule assigns an oxidation state to an atom, move on; the later rules handle whatever is left.
| DETERMINING OXIDATION STATES | |
|---|---|
| 1 | Atoms in their elemental form are in the zero oxidation state. |
| 2 | Monatomic ions in ionic compounds get their typical ionic charge. |
| 3 | Atoms in Molecular Compounds and Polyatomic Ions Oxygen −2 except in peroxide O22− Hydrogen +1 (when bonded with any nonmetals) Fluorine −1 always Halogens −1 except when bonded to Oxygen |
| 4 | The oxidation numbers sum to zero for a compound or to the overall charge of a polyatomic ion. |
Rule 1: Elements in Elemental Form — Oxidation State = 0
Any element that exists in its pure, uncharged elemental form has an oxidation state of zero. This applies whether the element is monatomic (e.g., Na, Fe) or diatomic (e.g., O₂, H₂). The absence of a charge is the telltale sign: no charge means elemental form, which means oxidation state of zero.
Rule 2: Monatomic Ions — Oxidation State Equals the Ionic Charge
For simple, single-atom ions (not polyatomic ions), the oxidation state is identical to the ion's charge. Na⁺ is +1; Cl⁻ is −1; Mg²⁺ is +2.
Transition metals and variable oxidation states: Main-group metals like sodium (Group 1) and calcium (Group 2) always carry the same ionic charge in compounds. Transition metals, however, can adopt multiple oxidation states. For example, copper can be +1 or +2. When you encounter a transition metal in a compound, determine its oxidation state from the charges of the other ions present. In CuCl₂, chloride is always −1 and there are two of them (total −2), so copper must be +2 to give a neutral compound.
Rule 3: Common Non-Metals with Predictable Oxidation States
After handling elemental forms and monatomic ions, assign oxidation states to the non-metals that almost always appear with the same value:
- Oxygen is almost always −2. The main exception is in peroxides, where oxygen is −1.
- Hydrogen is +1 when bonded to non-metals, and −1 when bonded exclusively to metals (metal hydrides).
- Fluorine is always −1 in compounds (it is the most electronegative element).
- Other halogens (Cl, Br, I) are typically −1 as monatomic ions, but can have other oxidation states in oxoanions (e.g., chlorate, perchlorate).
Rule 4: Oxidation States Must Sum to the Overall Charge
For a neutral compound, all oxidation states (multiplied by the number of each atom) must sum to zero. For a polyatomic ion, they must sum to the ion's overall charge. Use this rule to solve for the oxidation state of the one element you have not yet assigned — treat it algebraically if needed.
Example: Sodium Bicarbonate (NaHCO₃)
Work through the rules in order:
- No elements are in elemental form (it is a compound).
- Na⁺ is a monatomic ion from Group 1: oxidation state = +1.
- Hydrogen is bonded to non-metals here (not a metal hydride): oxidation state = +1. Oxygen is not part of a peroxide: oxidation state = −2 each; three oxygens total = −6.
- Running tally: (+1) + (+1) + (−6) = −4. The compound is neutral, so carbon must supply +4 to bring the total to zero. Carbon's oxidation state = +4.
Worked Examples: Assigning Oxidation States
Practice: Determine oxidation states for all elements in the following:
Na O2 NaCl NaHCO3
CO2 H2O2 SO42−
Carbon Dioxide (CO₂)
CO₂ is a molecular compound, so there are no monatomic ions. Oxygen is not in a peroxide, so it is −2; two oxygens contribute −4 total. Carbon must be +4 to bring the sum to zero.
Hydrogen Peroxide (H₂O₂) — The Peroxide Exception
Hydrogen peroxide is the most important exception to the "oxygen = −2" rule. Here oxygen is −1, not −2.
How to recognize a peroxide: Look for two oxygen atoms bonded together in the presence of hydrogen or a Group 1 metal. In H₂O₂, the two oxygens are each −1 and the two hydrogens are each +1: (+1)(2) + (−1)(2) = 0. Everything checks out. If you forget the peroxide rule, you can also derive it: hydrogen is always +1 with non-metals, so to make the sum zero with two hydrogens (+2 total), the two oxygens must together be −2 — meaning each oxygen is −1.
Sulfate Ion (SO₄²⁻) — Working with Polyatomic Ions
The key difference here is that the oxidation states must sum to −2 (the ion's charge), not zero.
Example: Sulfate Ion SO₄²⁻
- This is a polyatomic ion, so no elemental forms and no monatomic ions.
- Oxygen is not in a peroxide: oxidation state = −2 each; four oxygens = −8 total.
- The ion's overall charge is −2, so sulfur must satisfy: S + (−8) = −2, giving S = +6.
Important distinction: The oxidation state of oxygen in sulfate is −2 per atom. If an exam asks for the oxidation state of oxygen, the answer is −2, not −8. Oxidation state is always defined on a per-atom basis, even though you multiply by atom count when summing.
3 Shortcuts for Identifying Redox Reactions
Given a set of reactions, you need to determine which ones involve electron transfer. The defining characteristic of a redox reaction is a change in oxidation state for at least one element. Rather than assigning every oxidation state in every reaction from scratch, use these pattern-recognition shortcuts to work efficiently — especially under exam conditions.
Practice: Which of the following are oxidation-reduction reactions?
2Mg(s) + O2(g) ➔ 2MgO(s)
CaCO3(s) ➔ CaO(s) + CO2(g)
AgNO3(aq) + HCl(aq) ➔ AgCl(s) + HNO3(aq)
2AgNO3(aq) + Cu(s) ➔ Cu(NO3)2(aq) + 2Ag(s)
Shortcut 1: Look for Elements in Elemental Form
If any element appears in its elemental form (oxidation state = 0) on one side of the equation but not on the other, a change in oxidation state has definitely occurred. That is a redox reaction — no further calculation needed.
Example: Magnesium Combustion
2 Mg(s) + O₂(g) → 2 MgO(s)
Both Mg and O₂ begin in the elemental form (oxidation state = 0). Neither appears in elemental form on the right side. Both changed — this is a redox reaction. Specifically, magnesium is oxidized (0 → +2) and oxygen is reduced (0 → −2).
Shortcut 2: Single Replacement Reactions Are Always Redox
A single replacement (single displacement) reaction always involves one element displacing another from a compound. The displacing element starts in the elemental form (oxidation state = 0) and ends up in a compound, while the displaced element goes from a compound to the elemental form. A change in oxidation state is guaranteed.
Key rule: All single replacement reactions are redox reactions. Not all redox reactions are single replacement reactions, but the reverse is always true.
Shortcut 3: Double Replacement Reactions Are Never Redox
In a double replacement reaction, cations and anions simply trade partners. No element enters or leaves a compound in its elemental form; atoms just switch ionic partners. As a result, oxidation states remain unchanged throughout, and no electron transfer occurs.
Key rule: Double replacement reactions are never redox reactions. Once you recognize that reaction type, you can immediately rule out electron transfer without doing any calculations.
When in Doubt: Assign Oxidation States Directly
For reaction types that do not fit the shortcuts above — such as decomposition reactions — you need to assign oxidation states to all elements on both sides and check for any changes. Many decomposition reactions turn out not to be redox (no oxidation state changes), but this cannot be assumed without checking. Becoming fast and accurate at assigning oxidation states is essential, both for identifying redox reactions and for exam questions that ask for the oxidation state of a specific element.
Single Replacement Reactions and the Activity Series
A single replacement reaction occurs when a free element displaces another element from a compound. For example, a metal might displace another metal from an aqueous salt solution, or a metal might displace hydrogen from an acid. The key question is always: will this reaction actually occur? The answer depends on the activity series.
The Activity Series
The activity series ranks metals (and hydrogen) from most reactive at the top to least reactive at the bottom. "Most reactive" means most easily oxidized — most willing to give up electrons. To predict whether a single replacement reaction is spontaneous, apply this rule:
The activity series rule: A single replacement reaction is spontaneous only if the element being oxidized (the one that starts in elemental form) is higher on the activity series than the element being reduced. If it is lower, the reaction does not occur as written.
The activity series is almost always provided on exams — you are not expected to memorize it. If an exam asks which metal is "most active" or "most easily oxidized," simply choose the one highest on the list.
Example 1: Magnesium and Hydrochloric Acid (Spontaneous)
Reaction: Mg(s) + 2 HCl(aq) → MgCl₂(aq) + H₂(g)
Magnesium begins in elemental form (oxidation state = 0) and ends as Mg²⁺ (+2) — it is being oxidized. Hydrogen begins as H⁺ (+1) in the acid and ends as H₂ gas (0) — it is being reduced.
Consulting the activity series, magnesium sits well above hydrogen. The species being oxidized (Mg) is more easily oxidized than the species being reduced (H), so this reaction is spontaneous. Drop a piece of solid magnesium into hydrochloric acid and you will see the metal dissolve while hydrogen gas bubbles vigorously out of solution.
Example 2: Copper and Zinc Nitrate (Non-Spontaneous)
Reaction (as written): Cu(s) + Zn(NO₃)₂(aq) → Cu(NO₃)₂(aq) + Zn(s)
Copper begins in elemental form (0) and would become Cu²⁺ (+2) — it would be oxidized. Zinc would go from +2 to 0 — it would be reduced.
On the activity series, zinc is above copper, meaning zinc is more easily oxidized. Because the species being oxidized (Cu) is less easily oxidized than the species being reduced (Zn), this reaction does not occur. Drop a piece of solid copper into an aqueous zinc nitrate solution and nothing will happen — the copper simply sinks to the bottom.
Example 3: Manganese and Nickel — Net Ionic Form
Net ionic reaction: Mn(s) + Ni²⁺(aq) → Mn²⁺(aq) + Ni(s)
When spectator ions have been removed and only the net ionic equation is shown, you cannot rely on the visual shortcut of "who is alone vs. who is in a couple." Instead, reason directly from oxidation states: the element that starts in elemental form (manganese, oxidation state = 0 → +2) is being oxidized. Nickel is being reduced (+2 → 0).
For the reaction to be spontaneous, manganese must be higher on the activity series than nickel. Checking the list confirms that manganese is indeed above nickel, so this reaction is spontaneous.
Applying the Activity Series: The Full Logic
Whether you are looking at a complete molecular equation or a stripped-down net ionic equation, the underlying logic is the same. Identify which element is being oxidized (it starts in elemental form or its oxidation state increases). That element must be higher on the activity series than the element being reduced. If it is, the reaction proceeds. If it is not, the reaction does not occur as written.
Quick-check shortcut (complete molecular equations only): In a complete molecular equation, the element that starts all alone (in elemental form) is the one doing the displacing. If that element is higher on the activity series than the element it is displacing, the reaction is spontaneous. This shortcut breaks down for net ionic equations — there you must think in terms of oxidation states.
7 Key Takeaways
- Redox reactions are electron transfer reactions. Oxidation is the loss of electrons (oxidation state increases); reduction is the gain of electrons (oxidation state decreases). Remember LEO the lion says GER.
- Oxidation states track electron transfer. Assign them in order: (1) elemental form = 0, (2) monatomic ions = ionic charge, (3) predictable non-metals (O usually −2, H usually +1, F always −1, halogens usually −1), (4) remaining atoms must balance the total charge.
- The peroxide exception: Oxygen is −1, not −2, in peroxides. Recognize peroxides by the O–O bond, typically in the presence of hydrogen or a Group 1 metal.
- For polyatomic ions, oxidation states sum to the ion's overall charge, not zero. Oxidation state is always defined per atom, not per formula unit.
- Identifying redox reactions: Any change in oxidation state signals a redox reaction. Single replacement reactions are always redox; double replacement reactions are never redox.
- The activity series ranks metals by ease of oxidation. In a single replacement reaction, the reaction is spontaneous only if the element being oxidized is higher on the activity series than the element being reduced.
- This lesson is an introduction. A full treatment of redox reactions — including electrochemical cells, cell potentials, and electrolysis — appears later in the General Chemistry 2 electrochemistry chapter.