Redox Reaction Calculator

Calculate oxidation states, determine electron transfer, and classify redox reactions. Use the half-reaction method to balance equations in acidic or basic solutions. Identify oxidizing and reducing agents.

Oxidation State Rules: 1. Free elements = 0 (e.g., O₂, Fe, Cl₂) 2. Monatomic ions = charge (e.g., Na⁺ = +1, Cl⁻ = -1) 3. Oxygen usually = -2 (except peroxides -1, OF₂ +2) 4. Hydrogen usually = +1 (except metal hydrides -1) 5. Sum of oxidation states = total charge Electron Transfer: Electrons transferred = |Change in ox. state| × Number of atoms Oxidation: Ox. state increases (loses e⁻) Reduction: Ox. state decreases (gains e⁻) Half-Reaction Method (Acidic): 1. Write oxidation and reduction half-reactions 2. Balance atoms (except O and H) 3. Balance O by adding H₂O 4. Balance H by adding H⁺ 5. Balance charge by adding e⁻ 6. Multiply to equalize electrons 7. Add half-reactions and cancel Half-Reaction Method (Basic): 1. Balance as if acidic solution 2. Add OH⁻ equal to H⁺ to both sides 3. Combine H⁺ + OH⁻ → H₂O 4. Cancel H₂O appearing on both sides 5. Simplify Key Definitions: • Oxidation: Loss of electrons (OIL) • Reduction: Gain of electrons (RIG) • Oxidizing agent: Gets reduced (accepts e⁻) • Reducing agent: Gets oxidized (donates e⁻) • Disproportionation: Same element both oxidized and reduced Cell Potential: E°cell = E°(cathode) - E°(anode) Positive E°cell = spontaneous reaction Nernst Equation: E = E° - (RT/nF) × ln(Q) E = E° - (0.0592/n) × log₁₀(Q) at 25°C
Example 1 (Iron Oxidation): Fe²⁺ → Fe³⁺ Initial oxidation state: +2 Final oxidation state: +3 Change: +3 - (+2) = +1 Electrons: 1 e⁻ lost Process: Oxidation Half-reaction: Fe²⁺ → Fe³⁺ + e⁻ Fe²⁺ is the reducing agent Example 2 (Permanganate Reduction in Acid): MnO₄⁻ → Mn²⁺ Mn oxidation state: +7 → +2 Change: -5 (reduction, gains 5e⁻) Balance in acidic solution: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O MnO₄⁻ is the oxidizing agent Example 3 (Complete Redox in Acid): MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ (in H⁺) Oxidation: Fe²⁺ → Fe³⁺ + e⁻ (×5) Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O Balanced: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O 5 electrons transferred total Example 4 (Dichromate Reduction): Cr₂O₇²⁻ → 2Cr³⁺ Cr: +6 → +3 (2 atoms) Change: -3 per Cr × 2 = 6e⁻ total Half-reaction in acid: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O Example 5 (Chlorine Disproportionation): Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O (basic) Cl₂: 0 → -1 (reduction) and 0 → +1 (oxidation) One Cl atom reduced, one oxidized 1 electron transferred Example 6 (Zinc and Copper): Zn + Cu²⁺ → Zn²⁺ + Cu Zn: 0 → +2 (oxidation, loses 2e⁻) Cu: +2 → 0 (reduction, gains 2e⁻) 2 electrons transferred Zn is reducing agent, Cu²⁺ is oxidizing agent Example 7 (Hydrogen Peroxide): H₂O₂ → H₂O + O₂ O in H₂O₂: -1 O in H₂O: -2 (reduction) O in O₂: 0 (oxidation) Disproportionation reaction

What is a redox reaction?

A redox (reduction-oxidation) reaction involves the transfer of electrons between species. Oxidation is the loss of electrons (increase in oxidation state), and reduction is the gain of electrons (decrease in oxidation state). Both always occur together: one species is oxidized while another is reduced. Example: Zn + Cu²⁺ → Zn²⁺ + Cu (Zn loses electrons, Cu²⁺ gains electrons).

What is oxidation state (oxidation number)?

Oxidation state is the charge an atom would have if all bonds were 100% ionic. Rules: 1) Free elements = 0, 2) Monatomic ions = their charge, 3) Oxygen usually -2, 4) Hydrogen usually +1, 5) Sum equals total charge. Example: In SO₄²⁻, S is +6 because 4(O at -2) + S = -2, so S = +6.

How do you identify what is oxidized and reduced?

Compare oxidation states before and after reaction. Oxidation = oxidation state increases (loses electrons). Reduction = oxidation state decreases (gains electrons). Example: Fe²⁺ → Fe³⁺ is oxidation (+2 to +3, loses 1e⁻). Cl₂ → 2Cl⁻ is reduction (0 to -1, gains electrons). Use mnemonic: OIL RIG = Oxidation Is Loss, Reduction Is Gain.

What is the half-reaction method?

The half-reaction method separates redox reactions into oxidation and reduction half-reactions, balances each separately, then combines them. Steps: 1) Identify oxidation states, 2) Write half-reactions, 3) Balance atoms (except O, H), 4) Balance O with H₂O, 5) Balance H with H⁺ (acidic) or OH⁻ (basic), 6) Balance charge with e⁻, 7) Equalize electrons, 8) Add half-reactions.

How do you balance redox reactions in acidic solution?

Use half-reaction method in acidic medium: 1) Split into half-reactions, 2) Balance atoms except O and H, 3) Balance O by adding H₂O, 4) Balance H by adding H⁺, 5) Balance charge by adding e⁻, 6) Multiply to equalize electrons, 7) Add and cancel common terms. Example: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ in H⁺.

How do you balance redox reactions in basic solution?

Balance in acidic solution first, then convert to basic: After balancing in acid (with H⁺), add OH⁻ equal to H⁺ to both sides. H⁺ + OH⁻ combines to form H₂O. Cancel water on both sides if needed. Example: If you have 8H⁺ on left, add 8OH⁻ to both sides, forming 8H₂O on left.

What is a disproportionation reaction?

Disproportionation is when the same element is simultaneously oxidized and reduced. One atom increases oxidation state while another of the same element decreases. Example: 2H₂O₂ → 2H₂O + O₂. Oxygen in H₂O₂ is -1; in H₂O it's -2 (reduced), in O₂ it's 0 (oxidized). Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O is another example.

How many electrons are transferred in a redox reaction?

Count the change in oxidation state multiplied by the number of atoms. Example: 2Fe²⁺ → 2Fe³⁺ transfers 2 electrons total (2 atoms × 1 e⁻ each). For Cr₂O₇²⁻ → 2Cr³⁺, chromium goes from +6 to +3, change = -3 per Cr, total = 2 × 3 = 6 electrons gained. The electrons lost by one species equals electrons gained by the other.

What are common oxidizing and reducing agents?

Common oxidizing agents (get reduced, accept e⁻): MnO₄⁻ (permanganate), Cr₂O₇²⁻ (dichromate), H₂O₂, Cl₂, O₂, HNO₃. Common reducing agents (get oxidized, donate e⁻): Metals (Zn, Fe, Al), H₂, C, H₂S, SO₃²⁻, Fe²⁺, Sn²⁺. Strength varies; consult reduction potential table.

What is the relationship between redox and electrochemistry?

Redox reactions are the basis of electrochemistry. Standard reduction potentials (E°) predict spontaneity: E°cell = E°cathode - E°anode. Positive E°cell means spontaneous reaction. The Nernst equation relates E to concentrations: E = E° - (RT/nF)lnQ. Batteries, fuel cells, and electrolysis all involve controlled redox reactions.

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