Redox Reactions and Oxidation Numbers

4.3 Identifying Redox Reactions

Skill #1 – Oxidation Numbers for Elements & Monatomic Ions

• Simple redox reactions are easy to identify by observing which substance becomes more positive (oxidized) and which becomes less positive (reduced).
• More complex reactions require assigning "Oxidation Numbers" (States) to identify the two half-reactions.

Oxidation Number (O.N.)

• Definition: A number indicating how many electrons a substance has lost or gained compared to its isolated state.

1st Set of Rules: Isolated Elements & Monatomic Ions

A. Isolated Elements

• O.N. = 0
• Examples: Ca(s), Hg(l), F2(g), H2(g), O2(g) all have O.N. = 0
• Note: Oxygen (O) and Hydrogen (H) are special cases (explained later).

B. Monatomic Ions (and Multivalent Ions) in Solution, Isolated

• O.N. = ionic charge as given in the question
• Examples:
  • Cu$^{2+}$ (aq) = +2 (positive sign goes first)
  • S$^{2-}$ (aq) = -2 (negative sign goes first)
  • Mn$^{2+}$ (aq) = +2

C. Monatomic Ions (not Multivalent) as Part of a Compound

• O.N. = ionic charge as given on the Periodic Table
• Examples:
  • NaCl: Na = +1, Cl = -1
  • ZnI$_2$: Zn = +2, I = -1
  • Ag$<em>2$CO$</em>3$: Ag = +1, CO$_3$??? (explained later)

Examples: Writing Oxidation Numbers

• Write the oxidation number (state) under each term in the chemical reaction.
• Ex 1: 3Mg(s) + 2Cr$^{3+}$ (aq) → 3Mg$^{2+}$ (aq) + 2Cr (s)
• Ex 2: 8Cl$<em>2$ (g) + 8K$</em>2$S (aq) → S$_8$ (s) + 16 KCl (aq)

Answers to Examples

• Ex 1: 3Mg(s) + 2Cr$^{3+}$ (aq) → 3Mg$^{2+}$ (aq) + 2Cr (s)
  • 0 +3 +2 0
• Ex 2: 8Cl$<em>2$ (g) + 8K$</em>2$S (aq) → S$_8$ (s) + 16 KCl (aq)
  • 0 +1,-2 0 +1,-1

Skill #2 – 2nd Set of Rules: Complete Compounds, Polyatomic Ions

D. Complete Compounds (not Ions)

• O.N. = 0
• Examples:
  • NaCl = 0
  • Li$<em>2$CO$</em>3$ = 0
  • C$<em>6$H$</em>{12}$O$_6$ = 0

E. Polyatomic Ions

• O.N. = the regular ionic charge (either isolated or as part of an ionic compound)
• Examples:
  • SO$_4$ $^{2-}$ = -2 (overall, for the whole ion)
  • NH$_4$ $^{1+}$ = +1
  • NO$<em>3$ in Cu(NO$</em>3$)$_{2}$ = -1

F. Special Case: Oxygen Atoms within Compounds

• Each atom of oxygen, O.N. = -2
• Examples:
  • Within CrO$_4$ $^{2-}$, each O = -2
  • Within H$_2$O, O = -2
• Exception: O within O$_2$ (g) has O.N. = 0 (Rule A: Isolated elements)

G. Special Case: Hydrogen Atoms within Compounds

• Each atom of hydrogen, O.N. = +1
• Examples:
  • Within KOH, H = +1
  • Within H$<em>2$PO$</em>3$$^- , each H = +1
• Exception: H_2 (g) has O.N. = 0 (Rule A: Isolated elements)

H. The Tricky One: Elements (other than O & H) in Polyatomic Ions

• No direct rule! You have to solve for it.
• It varies depending on the compound.
• Procedure:
  1. Take the known O.N. for the compound (or ion).
  2. Assign O.N. for O, H, and monatomic ions using previous rules.
  3. Solve for the O.N. of the remaining element.
• Example: AgNO_3 … solve for N
  • Mark in O.N. for each element: +1 x -2
  • Set up equation & solve: 1(+1) + 1(x) + 3(-2) = 0
  • 1 + x - 6 = 0
  • x = +5 for N (in THIS compound!)

Examples: Calculating O.N. in Polyatomic Ions

• Ex2: Calculate O.N. for N in NH_3
  • 1(x) + 3(1) = 0
  • X = -3 in THIS compound
• Ex3: S in the sulfate ion SO_4$$^{2-}
  • 1(x) + 4(-2) = -2
  • X = +6 for S in THIS ion

Skill #3: Using O.N. to Determine if a Reaction is Redox or Not

• Redox reactions involve the transfer of electrons from one element to another.
• If two elements undergo changes in O.N., then the reaction is a redox reaction.
• If none of the elements experience a change in O.N., then the reaction is not a redox reaction.
• Procedure: Mark the O.N. for all elements before & after the reaction and compare.

The "Usual Four"

• The element whose O.N. increases is being oxidized (and is the reducing agent).
• The element whose O.N. decreases is being reduced (and is the oxidizing agent).

Examples: Identifying Redox Reactions and the "Usual Four"

• Ex1: Cu (s) + 2 AgNO3$(aq) → 2 Ag (s) + Cu(NO$3$)$_{2} (aq)
  • 0 +1+5-2 0 +2 +5 -2
  • Cu goes from O.N. = 0 to +2… Increase = oxidized
  • Ag goes from O.N. = +1 to 0… Decrease = reduced.
  • Changes in O.N. indicate that the reaction IS redox.
  • Substance oxidized = Cu(s)
  • Substance reduced = Ag^+ (aq)
  • Oxidizing agent = Ag^+$(aq) [or AgNO$_3]
  • Reducing agent = Cu (s)
  • NO_3$$^{1-}(aq) is a spectator ion (does nothing).
• Ex2: NaCl(aq) + AgNO3$(aq) → NaNO$3(aq) + AgCl(s)
  • +1-1 +1+5-2 +1+5-2 +1-1
  • Is it redox? No.
  • Why not? Because none of the oxidation numbers change. Nothing is oxidized (and so nothing is reduced).
• Ex3: C4$H$8$(g) + 5O$2$(aq) → 3CO$2$(g) + 4H$_2O(g)
  • -2+1 0 +4-2 +1-2
  • Is it redox? Yes
  • Why? Because the O.N. of C & O have changed!
  • Substance oxidized = C
  • Substance reduced = O
  • Oxidizing agent = O (or O_2, I suppose..)
  • Reducing agent = C (or C4$H$8, I suppose…)

Skill #4: Spontaneity of Redox Reactions

Example 1

• If a strip of Zn (s) is placed in a 0.1 mol/L solution of Cu(NO3$)${2}, after 5 minutes orange-brown deposits have formed on the surface of the Zn (s) and the blue-colour intensity of the solution has faded.
  • A) Was the redox reaction spontaneous or non-spontaneous?
  • B) What was the oxidation half-reaction?
  • C) What was the reduction half-reaction?
  • D) What was the oxidizing agent?

Answers

• A) Spontaneous
• B) Zn (s) → Zn^{2+}$(aq) + 2e$^-
• C) Cu^{2+}$(aq) + 2e$^- → Cu (s)
• D) Cu^{2+}$(aq) because it's the substance being reduced. The reaction was spontaneous, meaning Cu$^{2+} (aq) is a relatively strong oxidizing agent.

Example 2

• Another strip of Zn (s) is now placed in a 0.1 mol/L solution of Al(NO3$)${3}. After 5 minutes nothing has happened!
  • A) Was the redox reaction spontaneous or non-spontaneous?
  • B) What would have been the oxidizing agent?
  • C) Between Ex1 & Ex2, which was the stronger oxidizing agent?

Answers

• A) Non-spontaneous
• B) Al^{3+} (aq)
• C) Cu^{2+}$(aq) is a stronger oxidizing agent (was able to oxidize Zn) than Al$^{3+}$$ (aq) (was not able to oxidize Zn)
  • This leads us to the S.R.P. data sheet, which lists oxidizing agents in order of increasing oxidizing strength.