Oxidation-Reduction Reactions

Oxidation-Reduction Reactions

Meaning of Oxidation and Reduction

• Oxidation-Reduction reactions are often known as RedOx reactions.
• Oxidation was initially seen as the combination of an element with oxygen to produce oxides.
• Combustion is one type of oxidation.

Oxidation and Reduction

• Oxidation:
  • A species is oxidized when it loses electrons.
  • Example: Zinc loses two electrons to become $Zn^{2+}$.
• Reduction:
  • A species is reduced when it gains electrons.
  • Example: Hydrogen ions gain electrons to form $H_2$.

Oxidation

• Not all oxidation reactions involve oxygen, nor are they all combustion.
• Corrosion of a metal (e.g., iron rusting) is a slow oxidation process.
• $Fe + O<em>2 \rightarrow Fe</em>2O_3$
• Removal of fabric stains with bleach is another example of oxidation where the $OCl^-$ ion releases oxygen, which oxidizes stains.

Reduction

• Reduction is the opposite of oxidation and originally meant the removal of oxygen from a compound.
• Example: Iron ore is heated and gives up oxygen.
  • $2Fe<em>2O</em>3 + 3C \rightarrow 4Fe + 3CO_2$
• Reduction of iron also involves an oxidation process: the carbon is oxidized.

RedOx

• Oxidation and reduction reactions occur simultaneously; one cannot occur without the other.
• Reagents are known as oxidizers/oxidizing agents or reducers/reducing agents.
• A reducing agent is oxidized, and an oxidizing reagent is reduced.
• The overall reaction is often called a redox reaction.
• Many types of reactions are RedOx reactions:
  • Single replacement reactions
  • Combination reactions that involve two elements
  • Decomposition reactions that yield pure elements
  • Combustion reactions
• If no electron shifts occur, a reaction is not RedOx.
  • Double replacement reactions
  • Acid-base reactions

Electron Shift

• Many oxidation reactions do not involve oxygen in any way.
• Redox reactions are currently defined as the shift of electrons between reactants, although they can still be characterized as the gain or loss of oxygen.

Mnemonic

• LEO the lion says GER
  • Lose Electrons Oxidation
  • Gain Electrons Reduction

RedOx Reactions with Ions

• As metals react with nonmetals, electrons are transferred.
• Example: $Mg + S \rightarrow MgS$ $(Mg^{+2} + S^{-2})$
  • Oxidation: $Mg \rightarrow Mg^{+2} + 2e^-$
  • Reduction: $S + 2e^- \rightarrow S^{-2}$
• The two half-reactions occur simultaneously.
• The oxidizing reagent is reduced, and the reducing agent is oxidized.
• In the example above:
  • Magnesium loses electrons and is oxidized, but it reduces the sulfur.
  • Likewise, sulfur gains electrons and is reduced, but it oxidizes the magnesium.

RedOx with Covalent Compounds

• Electrons are not completely transferred in covalent compounds, unlike ionic compounds.
• Example: The reaction to form water.
• Bonding electrons are pulled towards oxygen in the H-O bond.

Oxidation and Reduction

• Table summarizing processes leading to oxidation and reduction:
  • Oxidation:
    • Complete loss of electrons (ionic reactions)
    • Shift of electrons away from an atom in a covalent bond
    • Gain of oxygen
    • Loss of hydrogen by a covalent compound
    • Increase in oxidation number
  • Reduction:
    • Complete gain of electrons (ionic reactions)
    • Shift of electrons toward an atom in a covalent bond
    • Loss of oxygen
    • Gain of hydrogen by a covalent compound
    • Decrease in oxidation number

Corrosion

• Corrosion is a widespread problem that costs billions every year.
• Occurs more rapidly in the presence of salts and acids.
• Not all metals corrode at equal rates (e.g., Au, Pt).
• Some metals corrode easily but form a protective oxide coating (e.g., Al).

Corrosion( continued )

• Sometimes corrosion is actually desired because it can enhance the surface appearance.

Corrosion Control

• Methods for controlling corrosion:
  • Coating with paint, oil/grease, plastic, another metal (e.g., chrome plating).
  • Sacrifice of a metal to protect another metal (e.g., galvanized iron).

Types of RedOx Reactions

• Combination: Oxidizing agent of one element will react with the reducing agent of the same element to produce the free element.
  • $I^- + IO<em>3^- + H^+ \rightarrow I</em>2 + H_2O$
• Decomposition:
  • Peroxides to oxides
  • Chlorates to chlorides
  • Electrolysis of compounds into elements
  • Carbonates to oxides

Oxidation Numbers

• Positive or negative number assigned to an atom to show the degree of oxidation or reduction.
• Assigned by a set of rules.
• As elements are oxidized, the oxidation number moves in a positive direction; as elements are reduced, the oxidation number moves in a negative direction.

Oxidation Numbers (continued)

• Example:
  • $Fe + Cl<em>2 \rightarrow FeCl</em>2$
    • The oxidation number of Fe goes from 0 to +2 because it loses two electrons.
  • $FeCl<em>2 + Cl</em>2 \rightarrow FeCl_3$
    • The oxidation number of Fe goes from +2 to +3 because it loses an electron.

Rules to Assign Oxidation Numbers

1. The oxidation number of a monatomic ion is equal in magnitude and sign to its ionic charge.
   • Example: The oxidation number of $Br^-$ is -1; $Fe^{3+}$ is +3.
2. The oxidation number of hydrogen in a compound is +1, except in metal hydrides (e.g., NaH), where it is -1.
3. The oxidation number of oxygen in a compound is -2, except in peroxides (e.g., $H<em>2O</em>2$), where it is -1, and in compounds with fluorine, where it is positive.
4. The oxidation number of an atom in uncombined (elemental) form is 0.
   • Example: The oxidation number of K in potassium metal (K) or N in nitrogen gas ($N_2$) is 0.
5. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0.
6. For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.

Oxidation Numbers (continued)

• Many elements have multiple oxidation numbers.
  • Carbon can be +4, -4, and +2.
• Transition metals often have multiple oxidation numbers.
  • Chromium can have 6, 5, 4, 3, 2, 1, -1, -2, although 2, 3, and 6 are the most common.

Problems – Oxidation Numbers

• Determine the oxidation number of each element in the following:
  • $SO_3$
  • $P<em>2O</em>5$
  • $NO_3^-$
• Determine the oxidation number of chlorine in each of the substances:
  • $KClO_3$
  • $Cl_2$
  • $Ca(ClO<em>4)</em>2$
  • $Cl_2O$

Changes in Oxidation Number

• In RedOx reactions, oxidation numbers of atoms change.
• $2AgNO<em>3 + Cu \rightarrow Cu(NO</em>3)_2 + 2Ag$
• $+1 \ \ 0 \ \ +2 \ \ 0$
• Which element is oxidized, and which is reduced in this example?

Balancing RedOx Reactions

• Color changes can indicate RedOx reactions, especially RedOx reactions that involve transition metals (e.g., $Cr^{+6}$ is orange, $Cr^{+3}$ is green).
• Often RedOx reactions are too complex to be balanced by trial and error.
• Two systematic methods that can be used:
  1. Use Oxidation Number Changes
  2. Use half-reactions
• Both use the fact that the total electrons in oxidation equals the total electrons in reduction.

Oxidation Number Changes

• Start with the skeleton equation for the redox equation.
• Example: $Fe<em>2O</em>3 + CO \rightarrow Fe + CO_2$
• Step 1 – Assign oxidation numbers to all atoms and write numbers above atoms on a per-atom basis.
• $+3 \ -2 \ \ +2 \ -2 \ \ 0 \ \ +4 \ -2$
• $Fe<em>2O</em>3 + CO \rightarrow Fe + CO_2$

Oxidation Number Changes (continued)

• Step 2: Identify which atoms are oxidized and which are reduced.
• $Fe<em>2O</em>3 + CO \rightarrow Fe + CO_2$
• Fe oxidation number decreases from +3 to 0, therefore Fe is reduced; C oxidation number increases from +2 to +4, therefore C is oxidized.

Oxidation Number Changes (continued)

• Step 3: Use a bracket to connect atoms that undergo oxidation and another bracket to connect atoms that undergo reduction; write oxidation number change beside each bracket.
• $Fe<em>2O</em>3 + CO \rightarrow Fe + CO_2$
• Change is the number of $e^-$ transferred.

Oxidation Number Changes (continued)

• Step 4: Make the total increase in oxidation number equal to the total decrease in oxidation number with coefficients.
• $Fe<em>2O</em>3 + 3CO \rightarrow 2Fe + 3CO_2$

Oxidation Number Changes (continued)

• Step 5: Finally, make certain that the equation is balanced for both atoms and charge.

Examples

• $K<em>2Cr</em>2O<em>7 + H</em>2O + S \rightarrow KOH + Cr<em>2O</em>3 + SO_2$
• $HNO<em>2 + HI \rightarrow NO + H</em>2O + I_2$

Half-Reactions

• Write a separate half-reaction each for oxidation and reduction. Balance the atoms in each half-reaction and then balance electrons in each half-reaction.
• Example: $HNO<em>3 + S \rightarrow SO</em>2 + NO + H_2O$
• Step 1: Write the unbalanced equation in ionic form.
• $H^+ + NO<em>3^- + S \rightarrow SO</em>2 + NO + H_2O$

Half Reactions (continued)

• Step 2: Write separate half-reactions for oxidation and reduction processes and put the oxidation numbers into place in each reaction.
  • Oxidation: $S \rightarrow SO_2$ (0 to 4)
  • Reduction: $NO_3^- \rightarrow NO$ (+5 to +2)
• NOTE – $H^+$ and $H_2O$ are not in the half-reactions initially because they are neither oxidized nor reduced; however, they will be needed to balance the half-reactions.

Half-Reaction

• Step 3: Balance the atoms in the half-reactions.
  • Oxidation - S is balanced but O is not; since the reaction occurs in acid, water and $H^+$ can be used to balance O and H as needed.
  • Add 2 water to balance O and 4 $H^+$ to balance H.
  • $2H<em>2O + S \rightarrow SO</em>2 + 4H^+$
  • The half-reaction is balanced in terms of atoms.

Half-Reactions (continued)

• Step 3, cont.:
  • Reduction – balance the half-reaction with the addition of two molecules of water to balance O and $H^+$ to balance H.
  • NOTE: add $H^+$ and $H_2O$ such that all atoms are on both sides of the reaction
  • $4H^+ + NO<em>3^- \rightarrow NO + 2H</em>2O$
  • The half-reaction is balanced in terms of atoms.
  • NOTE – a reaction that is run in base uses $OH^-$ and water to balance, rather than $H^+$ and water.

Half-Reactions (continued)

• Step 4: Add electrons to one side of each half-reaction to balance the charges (don’t add electrons to both sides).
  • Oxidation (electrons are on the product side)
    • $2H<em>2O + S \rightarrow SO</em>2 + 4H^+ + 4e^-$
  • Reduction (electrons are on the reactant side)
    • $4H^+ + NO<em>3^- + 3e^- \rightarrow NO + 2H</em>2O$
  • Half-reactions are balanced in terms of both atoms and charge.

Half-Reactions (continued)

• Step 5: Multiply each half-reaction by an appropriate number (usually the product is the lowest common multiple) to make the numbers of electrons equal in both.
  • Oxidation
    • $6H<em>2O + 3S \rightarrow 3SO</em>2 + 12H^+ + 12e^-$
  • Reduction
    • $16H^+ + 4NO<em>3^- + 12e^- \rightarrow 4NO + 8H</em>2O$

Half-Reactions (continued)

• Step 6: Add the two balanced half-reactions to give an overall equation.
• $6H<em>2O + 3S + 16H^+ + 4NO</em>3^- + 12e^- \rightarrow 3SO<em>2 + 12H^+ + 12e^- + 4NO + 8H</em>2O$
• Subtract out all terms that appear on both sides of the equation (in this example, subtract out $e^-$, $H^+$, and $H_2O$).
• $3S + 4H^+ + 4NO<em>3^- \rightarrow 3SO</em>2 + 4NO + 2H_2O$

Choice of Method to Balance

• Some RedOx reactions are simple enough to balance by inspection.
• Balance by oxidation number works best if oxidized and reduced species only occur once on each side of the equation and no acids or bases are involved.

Choice of Method to Balance (continued)

• Reactions that take place in acids or bases or where the same element is oxidized and reduced are best done as half-reactions.

Examples

• Basic solution
  • $Zn + NO<em>3^- \rightarrow Zn(OH)</em>4^{-2} + NH_3$
  • $MnO<em>4^- + C</em>2O<em>4^{-2} \rightarrow MnO</em>2 + CO_2$
• Acidic solution
  • $ClO<em>3^- + Cl^- \rightarrow Cl</em>2 + ClO_2$
  • $PH<em>3 + I</em>2 \rightarrow H<em>2PO</em>3^- + I^-$

Examples of reactions that are redox reactions (write equations)

• A piece of solid bismuth is heated strongly in oxygen to yield bismuth (V) oxide.
• A strip of copper metal is added to a concentrated solution of sulfuric acid.
• Magnesium turnings are added to a solution of iron (III) chloride.

Examples of redox reactions, cont.

• A stream of chlorine gas is passed through a solution of cold, dilute sodium hydroxide to produce sodium hypochlorite and sodium chloride.
• A solution of tin (II) chloride is added to an acidified solution of potassium permanganate; products are $MnO_2$ and $Sn^{+4}$.
• A solution of potassium iodide is added to an acidified solution of potassium dichromate; products are $Cr<em>2O</em>3$ and $I_2$