Oxidation Numbers and Redox Systems

TL;DR

Oxidation numbers help us track electron movement in chemical reactions. We use a set of rules to assign these numbers to elements in compounds and ions. Changes in oxidation numbers tell us if a reaction is an oxidation (losing electrons) or a reduction (gaining electrons).

1. The Mental Model

Think of oxidation numbers as a way to "score" how many electrons an atom appears to have gained or lost compared to its neutral state. This score isn't always real, but it's a helpful bookkeeping tool to see which atoms are changing electron partners.

2. The Core Material

Redox reactions (short for reduction-oxidation) are fundamental chemical processes involving the transfer of electrons. Understanding them hinges on correctly assigning oxidation numbers (also called oxidation states).

Rules for Assigning Oxidation Numbers

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Here's a step-by-step guide to assigning oxidation numbers. Always apply them in order:

1. Elements in their elemental form: The oxidation number of an uncombined element (like O₂, Na, H₂) is 0.
2. Monatomic ions: The oxidation number of a monatomic ion (like Na⁺, Cl⁻) is equal to its charge.
3. Group 1 metals: Always +1 in compounds.
4. Group 2 metals: Always +2 in compounds.
5. Fluorine: Always -1 in compounds.
6. Hydrogen: Usually +1 in compounds, except when bonded to a metal (e.g., NaH), where it's -1.
7. Oxygen: Usually -2 in compounds, except in peroxides (e.g., H₂O₂, where it's -1) or when bonded to fluorine (e.g., OF₂, where it's +2).
8. Sum of Oxidation Numbers:
   • For a neutral compound, the sum of oxidation numbers of all atoms is 0.
   • For a polyatomic ion, the sum of oxidation numbers of all atoms equals the charge of the ion.

Identifying Oxidation and Reduction

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Once you've assigned oxidation numbers to all relevant atoms in a reaction:

• Oxidation: An atom's oxidation number increases. This means it loses electrons.
• Reduction: An atom's oxidation number decreases. This means it gains electrons.

Remember the mnemonic "OIL RIG": Oxidation Is Loss (of electrons); Reduction Is Gain (of electrons).

Here's how you'd typically apply the rules:

graph TD
    A["Start: Assigning Oxidation Numbers"] --> B{"Is it an uncombined element?"}
    B -- Yes --> C["Oxidation Number = 0"]
    B -- No --> D{"Is it a monatomic ion?"}
    D -- Yes --> E["Oxidation Number = Ion Charge"]
    D -- No --> F{"Is it a Group 1/2 metal or Fluorine?"}
    F -- Yes --> G["Assign +1/+2/-1 respectively"]
    F -- No --> H{"Is it Hydrogen?"}
    H -- Yes --> I{"Is it bonded to a metal?"}
    I -- Yes --> J["Oxidation Number = -1"]
    I -- No --> K["Oxidation Number = +1"]
    H -- No --> L{"Is it Oxygen?"}
    L -- Yes --> M{"Is it a peroxide or bonded to F?"}
    M -- Yes --> N["Assign -1 (peroxide) or +2 (OF₂)"]
    M -- No --> O["Oxidation Number = -2"]
    L -- No --> P["Assign other known elements (if any)"]
    C --> Q["Calculate sum for neutral compound (=0) or polyatomic ion (=charge)"]
    E --> Q
    G --> Q
    J --> Q
    K --> Q
    N --> Q
    O --> Q
    P --> Q
    Q --> R["Deduce unknown oxidation number from the sum"]
    R --> S["End: All oxidation numbers assigned"]

3. Worked Example

Let's look at the reaction between copper(II) oxide and hydrogen gas:

CuO(s) + H₂(g) → Cu(s) + H₂O(l)

1. Assign oxidation numbers to reactants:

   • In CuO: Oxygen is -2 (Rule 7). Since the compound is neutral, Copper must be +2 (Rule 8).
   • In H₂: This is an uncombined element (Rule 1), so Hydrogen is 0.

2. Assign oxidation numbers to products:

   • In Cu: This is an uncombined element (Rule 1), so Copper is 0.
   • In H₂O: Oxygen is -2 (Rule 7). There are two Hydrogens, so each Hydrogen must be +1 for the compound to be neutral (2 * +1 + -2 = 0) (Rule 8).

3. Identify changes:

   • Copper (Cu): Goes from +2 in CuO to 0 in Cu. Its oxidation number decreased.
   • Hydrogen (H): Goes from 0 in H₂ to +1 in H₂O. Its oxidation number increased.
   • Oxygen's oxidation number remained -2 throughout.

4. Conclusion:

   • Since Copper's oxidation number decreased (from +2 to 0), Cu in CuO was reduced (gained electrons).
   • Since Hydrogen's oxidation number increased (from 0 to +1), H in H₂ was oxidized (lost electrons).

4. Key Takeaways

• Oxidation numbers are a tool to track electron distribution, not always actual charges.
• The sum of oxidation numbers in a neutral compound is zero; in an ion, it equals the ion's charge.
• An increase in oxidation number means oxidation (loss of electrons).
• A decrease in oxidation number means reduction (gain of electrons).
• Every oxidation must be accompanied by a reduction; they always occur together.
• Learn the hierarchy of rules for assigning oxidation numbers.

Common Mistakes to Avoid:

• Forgetting that elemental forms have an oxidation number of 0.
• Confusing the charge of an ion with an oxidation number for a single atom in a polyatomic ion.
• Applying the oxygen rule before checking if it's a peroxide or bonded to fluorine.
• Not considering the total number of atoms for polyatomic ions or compounds when summing oxidation numbers.

5. Now Try It

For the reaction: 2Na(s) + Cl₂(g) → 2NaCl(s)

1. Assign oxidation numbers to each atom in the reactants (Na and Cl₂).
2. Assign oxidation numbers to each atom in the product (NaCl).
3. Identify which element was oxidized and which was reduced, explaining why based on the change in oxidation number.

Success looks like correctly assigning all oxidation numbers and clearly stating which species was oxidized and which was reduced, along with the change in their respective oxidation numbers.