Electrochemistry notes

MDCAT Chemistry

Electrochemistry deals with the relationship between chemical reactions and electrical energy. This chapter covers oxidation and reduction, oxidation numbers, balancing redox equations, electrochemical cells, electrode potentials, the Standard Hydrogen Electrode, salt bridges, and electrolysis.

Redox Reactions

A redox reaction is a chemical reaction in which oxidation and reduction occur at the same time. Oxidation means loss of electrons, while reduction means gain of electrons. Since electrons lost by one species are gained by another species, oxidation and reduction cannot occur separately in an ordinary redox reaction.

The substance that loses electrons is the reducing agent. The substance that gains electrons is the oxidising agent. A reducing agent causes reduction of another substance and is itself oxidised. An oxidising agent causes oxidation of another substance and is itself reduced.

Oxidation: loss of electrons, increase in oxidation number, addition of oxygen, or removal of hydrogen.
Reduction: gain of electrons, decrease in oxidation number, removal of oxygen, or addition of hydrogen.
Reducing agent: donates electrons and is itself oxidised.
Oxidising agent: accepts electrons and is itself reduced.
Example: Zn + Cu2+ → Zn2+ + Cu. Zinc is oxidised and acts as the reducing agent. Cu2+ is reduced and acts as the oxidising agent.
Redox reaction: oxidation and reduction take place simultaneously.
Electron flow in an electrochemical cell is from the anode to the cathode through the external circuit.

Oxidation Number

Oxidation number is the apparent charge assigned to an atom in a compound or ion according to certain rules. It helps identify which atoms are oxidised or reduced. An increase in oxidation number represents oxidation, while a decrease represents reduction.

The oxidation number of a free element is zero. In a neutral compound, the sum of all oxidation numbers is zero. In a polyatomic ion, their sum is equal to the charge on the ion.

The oxidation number of a free element, such as Mg, O2, H2, or Cl2, is 0.
The oxidation number of a monoatomic ion is equal to its charge.
The usual oxidation number of oxygen is -2, but it is -1 in peroxides such as H2O2 and +2 in OF2.
The usual oxidation number of hydrogen is +1, but it is -1 in metal hydrides such as NaH.
Fluorine has oxidation number -1 in its compounds.
The oxidation number of potassium is +1, and that of alkaline earth metals is usually +2.
In K2MnO4: 2(+1) + Mn + 4(-2) = 0, so Mn has oxidation number +6.
In NH4NO3, nitrogen in NH4+ has oxidation number -3, while nitrogen in NO3- has oxidation number +5.
Beryllium generally has the same oxidation number, +2, in its compounds.
In OF2, oxygen is +2 because fluorine is -1. In H2O2, oxygen is -1.

Oxidation and Reduction in Terms of Oxidation Number

Oxidation and reduction can be recognised by comparing oxidation numbers before and after a reaction. This method is especially useful when the reaction does not clearly show transfer of electrons.

The total increase in oxidation number must equal the total decrease in oxidation number. This equality is the basis of the oxidation number method for balancing redox equations.

Increase in oxidation number means oxidation.
Decrease in oxidation number means reduction.
An element whose oxidation number increases is the reducing agent because it loses electrons.
An element whose oxidation number decreases is the oxidising agent because it gains electrons.
In Fe2+ → Fe3+ + e-, iron is oxidised because its oxidation number increases from +2 to +3.
In Cu2+ + 2e- → Cu, copper is reduced because its oxidation number decreases from +2 to 0.
Oxidation number is not always the actual charge on an atom. It is an assigned or apparent charge based on bonding rules.

Balancing Redox Equations by the Ion-Electron Method

The ion-electron method, also called the half-reaction method, balances oxidation and reduction separately. The two half-reactions are then combined after making the number of electrons equal.

The method depends on the medium. In acidic medium, H2O balances oxygen atoms and H+ balances hydrogen atoms. In basic medium, H2O and OH- are used. Charge must be balanced by adding electrons.

Step 1: Write separate oxidation and reduction half-reactions.
Step 2: Balance all atoms except oxygen and hydrogen.
Step 3: In acidic medium, use H2O to balance oxygen atoms.
Step 4: In acidic medium, use H+ to balance hydrogen atoms.
Step 5: Balance charge by adding electrons to the more positive side.
Step 6: Multiply half-reactions so that electrons lost equal electrons gained.
Step 7: Add the half-reactions, cancel common species, and check atoms and charge.
In basic medium, first balance as an acidic equation, then add OH- to both sides to neutralise each H+ and simplify water.
Example half-reaction: MnO4- → Mn2+ in acidic medium becomes MnO4- + 8H+ + 5e- → Mn2+ + 4H2O.

Electrochemical Cells

An electrochemical cell converts chemical energy into electrical energy or electrical energy into chemical energy. In a voltaic cell, also called a galvanic cell, a spontaneous redox reaction produces electricity. In an electrolytic cell, electrical energy is supplied to carry out a non-spontaneous reaction.

Oxidation always occurs at the anode and reduction always occurs at the cathode. In a voltaic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive and the cathode is negative. The names anode and cathode are based on the reactions, not on their signs.

Voltaic cell: converts chemical energy into electrical energy through a spontaneous reaction.
Electrolytic cell: uses electric current to carry out a non-spontaneous redox reaction.
Oxidation occurs at the anode in both types of cells.
Reduction occurs at the cathode in both types of cells.
Electrons move through the external wire from anode to cathode.
A voltaic cell can be converted into a reverse galvanic or electrolytic cell by applying an external source of electrical energy.
Electricity in a voltaic cell is produced by a spontaneous redox reaction and the directed flow of electrons.
In a voltaic cell, the cathode has a greater reduction potential than the anode.

Electrode Potential and Standard Reduction Potential

When a metal electrode is dipped into a solution containing its ions, a potential difference develops between the electrode and the solution. This is called electrode potential. It results from the tendency of the metal atoms to lose electrons or the ions to gain electrons.

Electrode potentials are usually written as reduction potentials. A more positive reduction potential means a greater tendency to gain electrons and undergo reduction. A more negative reduction potential means a greater tendency to lose electrons and form positive ions.

Standard reduction potential is measured under standard conditions, including 1 mol dm-3 ion concentration, 1 atm gas pressure, and a temperature of 298 K.
The standard reduction potential of the Standard Hydrogen Electrode is assigned a value of 0.00 V.
Greater reduction potential means greater tendency for reduction and smaller tendency to form positive ions.
Smaller or more negative reduction potential means greater tendency of the element to lose electrons and form positive ions.
Coinage metals such as Cu, Ag, and Au are comparatively less reactive because they have positive reduction potentials.
The standard reduction potential of Zn2+/Zn is -0.76 V.
For a spontaneous cell reaction, E°cell = E°cathode - E°anode.
A positive E°cell indicates that the cell reaction is spontaneous under standard conditions.

Standard Hydrogen Electrode and Salt Bridge

The Standard Hydrogen Electrode, or SHE, is the primary reference electrode used to measure standard electrode potentials. It consists of a platinum electrode in contact with hydrogen gas and an aqueous solution containing H+ ions of unit concentration.

The potential of another electrode is measured by connecting it to the SHE. A voltmeter measures the potential difference between the two electrodes. The SHE is assigned zero potential by convention, so the measured cell voltage gives the standard reduction potential of the other electrode when the other electrode acts as the cathode.

SHE contains a platinum electrode, hydrogen gas, and an acidic solution containing H+ ions.
Hydrogen gas is passed at 1 atm pressure over platinum at 298 K.
The standard concentration of H+ ions is 1 mol dm-3.
The value of SHE is taken as 0.00 V.
The electrode reaction is 2H+ + 2e- ⇌ H2.
A voltmeter measures the potential difference when the standard electrode potential is determined.
A salt bridge commonly contains KCl in a gel.
The salt bridge completes the electrical circuit, maintains electrical neutrality, and reduces direct mixing of the two solutions.
K+ and Cl- ions migrate to the compartments where their charges need to be balanced.

Electrolysis and Industrial Applications

Electrolysis is the chemical decomposition of an electrolyte by passing electric current through its molten or aqueous solution. At the cathode, positive ions gain electrons and are reduced. At the anode, negative ions or suitable molecules lose electrons and are oxidised.

The products of aqueous electrolysis depend on the ions present, their electrode potentials, and whether the electrodes are active or inert. Inert electrodes, such as platinum or graphite, do not normally take part in the reaction.

In dilute H2SO4 solution with inert electrodes, H2 is released at the cathode and O2 gas is released at the anode.
At the anode in acidified water: 2H2O → O2 + 4H+ + 4e-.
At the cathode: 2H+ + 2e- → H2.
Electrolysis requires an external source of electrical energy.
In the Castner-Kellner process for the manufacture of sodium hydroxide, titanium anodes are used.
The Castner-Kellner process uses electrolysis of brine and separates products formed at the electrodes.
At the cathode, water is reduced to produce hydrogen gas and hydroxide ions.
At the anode, chloride ions are oxidised to chlorine gas when brine is electrolysed.

Redox Titration and Oxidising Strength

Some redox reactions are used in volumetric analysis. A solution of known concentration is added to another solution until the reacting species are present in exact stoichiometric amounts. The end point may be detected by a colour change.

Potassium permanganate is a strong oxidising agent in acidic medium. Its intense purple colour allows it to act as its own indicator, so a separate indicator is usually not required.

KMnO4 is called a self-indicator because its purple colour shows the end point.
In acidic medium, MnO4- is reduced to Mn2+ and accepts 5 electrons per permanganate ion.
Dichromate ion, Cr2O7^2-, is reduced to Cr3+ in acidic medium and accepts 6 electrons per ion.
One mole of KMnO4 oxidises 5 moles of Fe2+ in acidic medium.
One mole of K2Cr2O7 oxidises 6 moles of Fe2+ in acidic medium.
Equal volumes of 1 M KMnO4 and 1 M K2Cr2O7 can oxidise different amounts of Fe2+ because their electron-accepting capacities differ.
For equal volume and concentration, K2Cr2O7 can oxidise more Fe2+ than KMnO4 under acidic conditions.

Key terms

Redox reaction: A reaction in which oxidation and reduction occur simultaneously.
Oxidation: Loss of electrons or increase in oxidation number.
Reduction: Gain of electrons or decrease in oxidation number.
Oxidising agent: A substance that accepts electrons and is itself reduced.
Reducing agent: A substance that donates electrons and is itself oxidised.
Oxidation number: The apparent charge assigned to an atom according to oxidation number rules.
Electrode potential: The potential difference developed between an electrode and its ionic solution.
Standard reduction potential: The reduction potential measured under standard conditions relative to the SHE.
Voltaic cell: A cell that produces electrical energy from a spontaneous redox reaction.
Electrolytic cell: A cell in which electrical energy drives a non-spontaneous redox reaction.
Anode: The electrode at which oxidation occurs.
Cathode: The electrode at which reduction occurs.
Standard Hydrogen Electrode: A reference electrode assigned a standard potential of 0.00 V.
Salt bridge: A device containing an electrolyte that completes the circuit and maintains electrical neutrality.
Electrolysis: The use of electric current to bring about a chemical decomposition or non-spontaneous redox reaction.
Self-indicator: A titrant whose own colour change indicates the end point, as KMnO4 in acidic medium.

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Chemistry shortcuts

Finding the limiting reactant and percentage composition

Convert every given mass or volume into moles first. The reactant that produces the least amount of the required product is the limiting reactant.

Write the balanced equation and calculate moles using n = mass/Mr.
Use the mole ratio to calculate the product. For percentage composition, use percentage = mass of element in one mole of compound divided by molar mass, multiplied by 100.
Example: Percentage of nitrogen in KNO3 = 14/101 × 100 = 13.86%.
Answer: 13.86% nitrogen.

Use gas volume at molar volume only when the gas conditions are stated or are standard conditions.

Using gas volume, pressure and temperature relations

At the same temperature and pressure, gas volume is directly proportional to the number of molecules. For changing conditions, use P1V1/T1 = P2V2/T2.

At constant temperature and pressure, divide or multiply the volume in the same ratio as the number of molecules.
Example: 10 mL H2 contains 2 × 10^3 molecules. Oxygen in 200 mL contains 20 × 2 × 10^3 = 4 × 10^4 molecules.
Answer: 4 × 10^4 molecules.
For a rigid container, increasing temperature increases molecular speed and mean free path if the gas remains in the same phase.

The direct volume to molecule ratio does not apply when temperature or pressure changes.

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