Electrochemistry

Electrochemistry is the branch of chemistry concerned with conversion of chemical energy to electrical energy OR electrical energy to chemical energy.

What we will study in this unit:

Oxidation and Reduction
Redox Reactions
Determining Spontaneity of REDOX reaction
Balancing half-reactions
Balancing REDOX reactions − Half-reaction Method
Determining Oxidation Numbers
Balancing REDOX reactions − Oxidation Number Method
REDOX Titrations
Electrochemical Cells
Electrolysis

Oxidation and Reduction

Oxidation occurs when a species LOSES electrons. When iron rusts, it is oxidizing:

2Fe

O₂

→

2FeO

What is happening to the IRON in this equation:

2Fe

→

Fe²⁺

2e⁻

OXIDATION half-reaction

Fe charge becomes MORE positive and 2e⁻ are released.

Where do the electrons go?

½O₂

2e⁻

→

O²⁻

REDUCTION half-reaction

Reduction occurs when a species GAINS electrons. This is what occurs when a metal is produced during a single replacement reaction.

Cu²⁺

2e⁻

→

REDUCTION half-reaction

Oxidizing Agent − is a species in a reaction that causes another species to oxidize. An oxidizing agent will be reduced during the reaction.

Reducing Agent − is a species in a reaction that causes another species to reduce. A reducing agent will be oxidized during the reaction.

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REDOX Reactions

In a REDOX reaction one species is REDuced and the other species is OXidized. Most SINGLE REPLACEMENT reactions are REDOX reactions.

In a REDOX reaction you will be expected to identify:

species oxidized
species reduced
oxidizing agent (O.A.)
reducing agent (R.A.)

Example:

Cu²⁺	+	Zn	→	Cu	+	Zn²⁺
reduced		oxidized
O.A.		R.A.

Memory aids:

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Determining the Spontaneity of a REDOX reaction

1) Find the two half reactions on the data table.
2) Strongest oxidizing agents are highest up on the left. Strongest reducing agents are lowest down on the right. If a pair of reactions is written so you have strongest oxidizing agent with strongest reducing agent then the reaction is spontaneous.

Things to be careful about:

some species such as H₂O₂ appear on both sides
some species will ONLY react in presence of H⁺ or OH⁻, if H⁺ or OH⁻ is part of the equation then the reaction will only go if ALL the conditions are met.
the exception to the point above is if the H⁺ or OH⁻ have a 10⁻⁷M written in brackets since this is the amount of H⁺ and OH⁻ that naturally appear in water.

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Balancing half-reactions
The steps you should follow are:
1. Balance MAJOR species using coefficients
2. Balance O by adding H₂O
3. Balance H by adding H⁺
4. Balance charge by adding e⁻ to the side of the equation that has too much + charge
Example
1) Ru²⁺ → Ru

STEP 1: major species balanced
STEP 2: no O in equation
STEP 3: no H in equation
STEP 4: add 2e⁻ to LEFT side

Ru²⁺ + 2e⁻ → Ru

2) Al₂O₃ → 2Al

Al₂O₃ + 6H⁺ + 6e⁻ → 2Al + 3H₂O

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Balancing REDOX Equations − Half-Reaction Method
Steps:
1. Identify and write each half-reaction
2. Balance each half-reaction
3. Multiply each half-reaction equation so the number of electrons is equal
4. Cancel electrons and other species that appear on both sides
Examples
1) Mn²⁺ + IO₃⁻ → MnO₄⁻ + I₂ (acidic)

Mn²⁺ + 4H₂O → MnO₄⁻ + 8H⁺ + 5e⁻

2IO₃⁻ + 12H⁺ + 10e⁻ → I₂ + 6H₂O

_________________________________________________________

2×( Mn²⁺ + 4H₂O → MnO₄⁻ + 8H⁺ + 5e⁻ )

2IO₃⁻ + 12H⁺ + 10e⁻ → I₂ + 6H₂O

_________________________________________________________

2Mn²⁺ + ²8H₂O → 2MnO₄⁻ + ⁴16H⁺ + ~~10e~~⁻

2IO₃⁻ + ~~12H~~⁺ + ~~10e~~⁻ → I₂ + 6H₂O

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2Mn²⁺ + 2IO₃⁻ + 2H₂O → 2MnO₄⁻ + 4H⁺ + I₂

2) C₂O₄²⁻ + IO₃⁻ → CO₂ + I₂ (acidic)

C₂O₄²⁻ → 2CO₂ + 2e⁻

2IO₃⁻ + 12H⁺ + 10e⁻ → I₂ + 6H₂O

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5×( C₂O₄²⁻ → 2CO₂ + 2e⁻ )

2IO₃⁻ + 12H⁺ + 10e⁻ → I₂ + 6H₂O

_________________________________________________________

5C₂O₄²⁻ → 10CO₂ + ~~10e~~⁻

2IO₃⁻ + 12H⁺ + ~~10e~~⁻ → I₂ + 6H₂O

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2IO₃⁻ + 5C₂O₄²⁻ + 12H⁺ → I₂ + 10CO₂ + 6H₂O

3) Balancing in basic solution: O₂ + Cr(OH)₃ → HO₂⁻ + CrO₄²⁻ (basic)

O₂ + H⁺ + 2e⁻ → HO₂⁻

Cr(OH)₃⁻ + H₂O → CrO₄²⁻ + 5H⁺ + 3e⁻

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3×( O₂ + H⁺ + 2e⁻ → HO₂⁻ )

2×( Cr(OH)₃⁻ + H₂O → CrO₄²⁻ + 5H⁺ + 3e⁻ )

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3O₂ + 3H⁺ + 6e⁻ → 3HO₂⁻

2Cr(OH)₃⁻ + 2H₂O → 2CrO₄²⁻ + 10H⁺ + 6e⁻

_________________________________________________________

3O₂ + 2Cr(OH)₃⁻ + 3H⁺ + 2H₂O → 3HO₂⁻ + 2CrO₄²⁻ + ⁷10H⁺

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3O₂ + 2Cr(OH)₃⁻ + 2H₂O → 3HO₂⁻ + 2CrO₄²⁻ + 7H⁺
+ 7OH⁻ + ~~7OH~~⁻
^{‾‾‾‾‾‾‾‾}
⁵7H₂O

_________________________________________________________

3O₂ + 2Cr(OH)₃⁻ + 7OH⁻ → 3HO₂⁻ + 2CrO₄²⁻ + 5H₂O

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Determining Oxidation Numbers
Finding the oxidation number of a substance is like finding the "charge" or combining capacity of an atom in a molecule.

Rules
- Elements ALWAYS have oxidation number ZERO
- Hydrogen is ALWAYS +1
- Alkali ions are ALWAYS +1
- Alkaline Earth ions are ALWAYS +2
- Oxygen is USUALLY −2, but is −1 in peroxides
- ALL OTHERS MUST BE CALCULATED! Use the sum of charges of species
Example
HPO₄²⁻
H = +1 P = x O = −2
⇒ 1 + x + 4×(−2) = −2
⇒ x = +5

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Balancing REDOX Equations − Oxidation Number Method
See Hebden!

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REDOX Titrations
Uses stoichiometry of REDOX reactions to find the concentration of an unknown species.

Most Common Oxidizing Agents
MnO₄⁻ and Cr₂O₇²⁻ because they both change colour when they reduce ⇒ these substances act as an indicator as well as a reactant.

Example
A 25.00mL sample of an aqueous Fe²⁺ salt solution is titrated using acidified KMnO₄. 24.83mL of 0.7864M KMnO₄ was required to react all of the Fe²⁺. What is the [Fe²⁺]?

5×( Fe²⁺ → Fe³⁺ + e⁻ )

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

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5Fe²⁺ + MnO₄⁻ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O

Most Common Reducing Agent
I⁻ is used as a reducing agent then the I₂ can be titrated since starch will indicate for the presence of I₂. This takes place in two reactions:

REACTION 1: Producing I₂ Using Excess I⁻ (note: substance that is our unknown is the LIMITING REAGENT!)

2I⁻ → I₂ + 2e⁻

H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O

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2I⁻ + H₂O₂ + 2H⁺ → I₂ + 2H₂O

REACTION 2: Titrating the I₂

I₂ + 2e⁻ → 2I⁻

2S₂O₃²⁻ → S₄O₆²⁻ + 2e⁻

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2S₂O₃²⁻ + I₂ → S₄O₆²⁻ + 2I⁻

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Electrochemical Cells
Electrochemical cells are SPONTANEOUS redox reactions, they produce a flow of electrons. This is what happens when you use a battery in a cellphone or camera.

The parts of an electrochemical cell are:
- Beaker with one half reaction that will be the reduction reaction.
- Beaker with one half reaction that will be the oxidation reaction.
- Voltmeter and wires to allow electron flow
- Salt Bridge containing KNO₃ or NaNO₃ to allow ion flow
Steps:
1. Determine which half-reaction is the oxidation and which one is the reduction.
2. The OXIDATION reaction occurs at the ANODE. The REDUCTION reaction occurs at the CATHODE.
3. Label the anode and cathode.
4. Label the direction of ion flow
In the diagram electrons flow FROM the anode TO the cathode through the wire.

Cations flow TOWARDS the cathode. Anions flow TOWARDS the anode.

Memory Aid
An Ox chases a Red Cat

Cell Potential
- Since we have a POTENTIAL DIFFERENCE there needs to be a point of comparison. The standard chosen is the HYDROGEN half-reaction. This is assigned a value of ZERO, and all other values are determined based on this standard.
- Use table of reduction potentials to determine the potential of a cell.
- Calculate the cell potential using: E^o_cell = E^o_red − E^o_ox
- Cell potential is POSITIVE for a SPONTANEOUS REDOX reaction.
The cell potential for the diagram is calculated using: E^o_cell = E^o_red − E^o_ox = +0.34 − (−0.76) = 1.10V

Corrosion
Corrosion is a spontaneous redox reaction which causes a metal to oxidize. When iron corrodes this process is called rusting

The corrosion of iron is shown in the diagram below:

You should note that corrosion is really an application of an electrochemical cell that forms in the presence of oxygen and moisture.

Corrosion Prevention
- Using a barrier ⇒ prevents moisture and oxygen from reaching the metal by adding a coating
  - paint
  - plastic coatings
  - grease
- Cathodic Protection = Sacrificial Protection = Anodizing
  - attach a more active metal which will corrode preferentially
  - used for pipelines and large structures like oil rigs as well as household objects
  - galvanizing is a specific example of cathodic protection where iron is coated with Zn.
- Change the chemical environment
  - remove moisture
  - remove oxygen
  - increase the pH
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Electrolysis
In an electrolysis the reaction is NOT spontaneous ⇒ electricity is required to make the reaction occur.

There are THREE different types of electrolytic cell that you will be expected to know about
1. Electrolysis of a Molten (Melted=Liquid) Binary Salt
2. Electrolysis of an Aqueous Solution with INERT electrodes (Graphite or Pt)
3. Electrolysis of an Aqueous Solution with REACTIVE electrodes
Electrolysis of a Molten Salt

ANODE reaction: 2Cl⁻ → Cl₂ + 2e⁻ Cathode reaction: Na⁺ + e⁻ → Na

Electrolysis of an Aqueous Salt With Inert Electrodes

ANODE reaction: 2Cl⁻ → Cl₂ + 2e⁻ Cathode reaction: Na⁺ + e⁻ → Na
OR:
H₂O → ½O₂ + 2H⁺ + 2e⁻ 2H₂O + 2e⁻ → H₂ + 2OH⁻

Look at the table and find the CLOSEST together REDOX pair. In this case the closest pair are:

ANODE reaction: 2Cl⁻ → Cl₂ + 2e⁻ Cathode reaction: 2H₂O + 2e⁻ → H₂ + 2OH⁻

Electrolysis of an Aqueous Salt With Reactive Electrodes

1) Electroplating:

2) Electrorefining

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3O₂	+	2Cr(OH)₃⁻	+	2H₂O	→	3HO₂⁻	+	2CrO₄²⁻	+	7H⁺
	+	7OH⁻							+	~~7OH~~⁻
										^{‾‾‾‾‾‾‾‾}
										⁵7H₂O

Electrochemistry

Oxidation and Reduction

REDOX Reactions

Memory aids:

Determining the Spontaneity of a REDOX reaction

Balancing half-reactions

Example

Balancing REDOX Equations − Half-Reaction Method

Examples

Determining Oxidation Numbers

Rules

Example

Balancing REDOX Equations − Oxidation Number Method

REDOX Titrations

Most Common Oxidizing Agents

Example

Most Common Reducing Agent

Electrochemical Cells

The parts of an electrochemical cell are:

Steps:

Memory Aid

Cell Potential

Corrosion

Corrosion Prevention

Electrolysis

Electrolysis of a Molten Salt

Electrolysis of an Aqueous Salt With Inert Electrodes

Electrolysis of an Aqueous Salt With Reactive Electrodes

1) Electroplating:

2) Electrorefining