Electrochemistry Notes

Reduction and Oxidation Chemistry

• Redox reactions involve both reduction and oxidation processes.
• Electrochemistry is the study of these redox reactions.

Oxidation & Reduction

• LEO says GER:
  • Loss of electrons is oxidation.
  • Gain of electrons is reduction.
• Electrochemistry deals with these processes.

Electrochemistry: Oxidation Numbers

• Oxidation numbers are assigned to atoms in a compound to track electron movement.
• Rules for Assigning Oxidation Numbers:
  1. The oxidation number of an uncombined element is 0.
  2. The oxidation number of a monatomic ion is the same as its charge.
  3. The sum of the oxidation numbers of the atoms in a polyatomic ion equals its charge.
• Using the Periodic Table: For instance, chlorine (Cl) as an ion is Cl- (charge -1), so its oxidation number is -1.
• Example: Phosphate Ion (PO43-)
  • Oxygen's oxidation number is -2.
  • The sum of oxidation numbers must equal the ion's charge (-3).
  • $P + 4(O) = -3$
  • $P + 4(-2) = -3$
  • $P = +5$

More Rules for Oxidation Numbers

• 4) Hydrogen (H) in a compound is +1, except in metal hydrides (where it is -1).
• 5) Oxygen (O) in a compound is -2, except in peroxides (where it is -1).
• 6) The sum of oxidation numbers in a neutral compound is zero.
• Metal Hydrides: Metals bonded with Hydrogen (NaH).

Practice: Oxidation Numbers

• Example 1: KNO3
  • O: -2
  • K: +1
  • N: +5 (calculated to balance the charge)
• Example 2: Na2SO3
  • O: -2
  • Na: +1
  • S: +4 (calculated to balance the charge)
• Strategy: Apply known rules first, then deduce the rest.

More Practice: Oxidation Numbers

• Example 1: MgH2
  • H: -1 (metal hydride)
  • Mg: +2
• Example 2: HIO
  • O: -2
  • H: +1
  • I: +1

Oxidation and Reduction Defined

• Oxidation: A species loses electrons.
  • Example: $Zn → Zn^{2+} + 2e^-$
  • Zinc metal becomes a zinc ion by losing two electrons.
• LEO says GER!

Reduction Defined

• Reduction: A species gains electrons.
  • Example: $2H^+ + 2e^- → H_2$
  • Hydrogen ions gain electrons to form hydrogen gas.
• LEO says GER!

Oxidizing and Reducing Agents

• Oxidizing Agent: The species that is reduced.
  • Example: H+ oxidizes Zn by taking electrons from it.
• Reducing Agent: The species that is oxidized.
  • Example: Zn reduces H+ by giving it electrons.

Balancing Redox Equations: Half-Reactions Method

• Consider the reaction: $Zn(s) + Cu(NO<em>3)</em>2(aq) → Zn(NO<em>3)</em>2(aq) + Cu(s)$
  1. Determine Oxidation Numbers:
     • Zn: 0 to +2
     • Cu: +2 to 0
     • $NO_3$ remains unchanged (spectator ion)
  2. Separate Half-Reactions:
     • Oxidation: $Zn → Zn^{2+}$
     • Reduction: $Cu^{2+} → Cu$
     • Spectator ions ($NO_3^-$) are omitted.

Balancing Redox Equations (cont.)

• 3. Balance Atoms:
  • Oxidation: $Zn → Zn^{2+}$
  • Reduction: $Cu^{2+} → Cu$
• 4. Balance Charge with Electrons:
  • Oxidation: $Zn → Zn^{2+} + 2e^-$
  • Reduction: $Cu^{2+} + 2e^- → Cu$
• 5. Equalize Electrons by Multiplying Half-Reactions (if necessary):
  • In this case, both half-reactions already have 2 electrons, so no multiplication is needed.

Balancing Redox Equations (cont.)

• 6. Add Half-Reactions and Cancel Like Terms:
  • $Zn + Cu^{2+} + 2e^- → Zn^{2+} + Cu + 2e^-$
  • Net Ionic Equation: $Zn + Cu^{2+} → Zn^{2+} + Cu$
• 7. Add Spectator Ions and Balance (Optional):
  • $Zn + Cu(NO<em>3)</em>2 → Zn(NO<em>3)</em>2 + Cu$
  • Check that elements and charge are balanced.

Balancing Redox Equations: Another Example

• Consider: $CuS + HNO<em>3 → Cu(NO</em>3)<em>2 + NO</em>2 + SO_2$
  1. Determine Oxidation Numbers:
     • Cu: +2, S: -2, N (in $NO<em>3^-$) : +5, N (in $NO</em>2$) : +4, S (in $SO_2$) : +4
  2. Separate Half-Reactions (leaving out spectators):
     • Oxidation: $S^{-2} → SO_2$
     • Reduction: $NO<em>3^- → NO</em>2$

Balancing Redox Equations (cont.)

• 3. Balance Atoms (add $H_2O$ and $H^+$):
  • Oxidation: $S^{-2} + 2H<em>2O → SO</em>2 + 4H^+$
  • Reduction: $NO<em>3^- + 2H^+ → NO</em>2 + H_2O$
• 4. Balance Charge with Electrons:
  • Oxidation: $S^{-2} + 2H<em>2O → SO</em>2 + 4H^+ + 6e^-$
  • Reduction: $NO<em>3^- + 2H^+ + 1e^- → NO</em>2 + H_2O$
• 5. Equalize Electrons:
  • Multiply the reduction half-reaction by 6.
  • $S^{-2} + 2H<em>2O → SO</em>2 + 4H^+ + 6e^-$
  • $6(NO<em>3^- + 2H^+ + 1e^- → NO</em>2 + H_2O)$

Balancing Redox Equations (cont.)

• 6. Add Half-Reactions and Cancel Like Terms:
  • $S^{-2} + 6NO<em>3^- + 12H^+ + 2H</em>2O + 6e^- → SO<em>2 + 6NO</em>2 + 6H_2O + 4H^+ + 6e^-$
  • $S^{-2} + 6NO<em>3^- + 8H^+ → SO</em>2 + 6NO<em>2 + 4H</em>2O$
• 7. Add Spectator Ions and Balance (Optional):
  • $CuS + 8HNO<em>3 → Cu(NO</em>3)<em>2 + 6NO</em>2 + SO<em>2 + 4H</em>2O$

Voltaic Cells (Galvanic Cells)

• Spontaneous redox reactions release energy via electron transfer.
• Example: $Zn + Cu^{2+} → Zn^{2+} + Cu$

Voltaic Cells: Harnessing Energy

• Energy from electron flow can do work if electrons pass through an external device.
• A voltaic cell is a setup that enables this.

Voltaic Cells: Components

• Anode: Where oxidation occurs.
• Cathode: Where reduction occurs.
• RED CAT AN OX (Reduction occurs at the Cathode and Oxidation occurs at the Anode).

Voltaic Cells: Charge Balance

• Electron flow from anode to cathode creates charge imbalance, stopping the flow.

Voltaic Cells: Salt Bridge

• Salt Bridge: A U-shaped tube containing a salt solution to maintain charge balance.
  • Cations move toward the cathode.
  • Anions move toward the anode.

Voltaic Cells: Electron Flow

• Electrons leave the anode, flow through a wire to the cathode (A before C).
• Cations formed at the anode dissolve into the solution.

Voltaic Cells: Cathode Activity

• Electrons reaching the cathode attract cations in the solution.
• Cations take electrons and deposit as neutral metal on the cathode.

Electromotive Force (emf)

• Electrons flow spontaneously from higher to lower potential energy (like water in a waterfall).

Electromotive Force (emf) Defined

• Electromotive Force (emf): The potential difference (energy) between the anode and cathode.
• Also called cell potential or voltage (
• $E_{cell}$), measured in volts (V).

Standard Reduction Potentials

• Similar to electronegativity and reactivity concept.
• Reactive Nonmetals: Tend to gain electrons (reduction).
• Reactive Metals: Tend to lose electrons (oxidation).
• Halogens: Gain electrons (reduction).
• Alkali Metals: Lose electrons (oxidation).

Standard Reduction Potentials Table

• A table of reduction half-reactions and their corresponding potentials (V) is used for calculations.

Standard Cell Potentials

• The cell potential at standard conditions is calculated as follows:
• $E<em>{cell}^° = E</em>r^°(cathode) - E_r^°(anode)$
• If $E_{cell}^°$ is positive, the reaction is spontaneous.
• If $E_{cell}^°$ is negative, the reaction is non-spontaneous.

Cell Potentials: Example Calculation

• Given:
  • Oxidation: $E_r^° = -0.76 V$
  • Reduction: $E_r^° = +0.34 V$
• Calculation:
  • $E_{cell}^° = (+0.34 V) - (-0.76 V) = +1.10 V$

Standard Hydrogen Electrode (SHE)

• Reduction potentials are referenced to SHE.
• By definition, the reduction potential for hydrogen is 0 V:
  • $2H^+(aq, 1M) + 2e^- → H_2(g, 1 atm)$

Voltaic Cells: Cell Notation (aka Galvanic Cells)

• Cell notation represents voltaic cells.
• Single line (|) represents a phase boundary (electrode to electrolyte).
• Double line (||) represents a physical boundary (porous boundary).

Cell Notation: Examples

• Examples of cell notation matching cell descriptions.
  • A copper-magnesium cell: Mg(s) | Mg2+(aq) || Cu2+(aq) | Cu(s)
  • Two tin electrodes in solution of tin(II) chloride and tin (IV) chloride respectively: Sn(s) | SnCl2(aq) || SnCl4(aq) | Sn(s)

Oxidizing and Reducing Agents: Strength

• Strongest Oxidizers: Have the most positive reduction potentials.
• Strongest Reducers: Have the most negative reduction potentials.

Voltaic Cells: Example - Silver-Copper Cell

• Cell Notation: $Cu(s) | Cu^{2+}(aq) || Ag^+(aq) | Ag(s)$
• An ox ate a red cat:
• Anode: Where strongest reducing agent reacts, Oxidation
• Cathode: Where strongest oxidizing agent reacts, Reduction
• Reduction Half-Reaction (Cathode):
  • $Ag^+(aq) + e^- → Ag(s)$
• Oxidation Half-Reaction (Anode):
  • $Cu(s) → Cu^{2+}(aq) + 2e^-$
• Overall (net) equation.
• Balance the half-reactions and add together to create the net equation.

Voltaic Cells: Silver-Copper Cell Details

• Silver ions ($Ag^+$) are the strongest oxidizing agents, undergoing reduction at the cathode to form Ag(s).
• Copper atoms (Cu) are the strongest reducing agents, undergoing oxidation at the anode, releasing electrons and forming $Cu^{2+}$ ions.
• Electrons flow from the copper anode to the silver cathode.
• Cations from the salt bridge move into the cathode compartment, and anions move into the anode compartment to maintain electrical neutrality.

Standard Cells: Analysis Rules

1. Determine the cathode: The electrode where the strongest oxidizing agent present in the cell reacts.
   • SOA: The OA that is closet to the top on the left side of the redox table.
   • Copy the reduction half-reaction.
2. Determine the anode: The electrode where the strongest reducing agent present in the cell reacts.
   • SRA: The RA that is closet to the bottom on the right side of the redox table
   • Copy the oxidation half-reaction (reverse the half-reaction)
3. Determine the overall cell reaction: Balance the electrons for the two half reactions (but DO NOT change the $E^0r$
4. Determine the standard cell potential: $E^0 cell = E^0r (cathode) – E^0r (anode)$

Standard Cells and Cell Potentials: Example

• #1 Example: What is the standard potential of the cell represented below:
• Determine the cathode and anode
• Determine the overall cell reaction
• Determine the standard cell potential

Oxidizing and Reducing Agents: Potential Difference

• More positive the Ered(V) the greater the difference between the two (cathode and anode), the greater the voltage of the cell.
• Example: Cu2+ + 2e- → Cu $Ered(V)$ +0.34 (Cathode).
• $Zn > Zn^{2+} + 2e^−, Ered(V) -0.76$(Anode)
• $E_{cell} = (+0.34) - (-0.76) = +1.10 V$

Applications of Oxidation-Reduction Reactions

• Batteries: Convert chemical energy into electrical energy.
  • Lead-acid batteries.
  • Alkaline batteries.

Batteries: History and Function

• Alessandro Volta invented the first electric cell, inspired by Luigi Galvani's observation of frog leg muscle twitching with different metals.
  • Galvani thought this was due to “animal electricity”.
  • Volta recognized this experiment’s electric potential.
• Alessandro Volta’s’s battery produced steady electric current

Batteries: Volta's Design

• Volta's first battery consisted of several bowls of brine (NaCl(aq)) connected by metals that dipped from one bowl to another.
• His revised design consisted of a sandwich of two metals separated by paper soaked in salt water.

Batteries: Technological Impact

• Alessandro Volta’s invention was an immediate technological success because it produced electric current more simply and reliably than methods that depended on static electricity.
• It also produced a steady electric current –something no other device could do.

Hydrogen Fuel Cells

• Fuel Cell: A voltaic cell where the oxidation of a fuel produces electric energy.
• Operation:
• Hydrogen is oxidized at the anode, and oxygen is reduced at the cathode, producing water as a byproduct and generating electricity.

Corrosion

• Corrosion: Loss of metal due to oxidation-reduction reactions with the environment.

Corrosion Prevention

• Galvanization: Coating iron with zinc.
• Mechanism:
  • Zinc acts as the anode and is oxidized instead of iron (cathode).
  • Reactions:
    • $Zn → Zn^{2+} + 2e^-$
    • $O<em>2 + 4H+ + 4e^- → 2H</em>2O$

Electrolysis

• Using electrical energy to drive a reaction in a non-spontaneous direction.
• Using electrical energy to bring about a chemical reaction

Electrolysis Applications

• Using electrical energy to drive a reaction in a nonspontaneous direction
• Used for electroplating, electrolysis of water, separation of a mixture of ions, etc.
  • (Most negative reduction potential is easiest to plate out of solution.)