Electrochemical cells are devices that convert chemical energy into electrical energy. They consist of two half-cells connected by a salt bridge, with electrodes immersed in electrolyte solutions. Understanding their components and structure is crucial for grasping how they generate electricity.

Electrode potentials determine the direction of electron flow in electrochemical cells. By calculating cell potentials and relating them to Gibbs free energy, we can predict the spontaneity of redox reactions. This knowledge is essential for designing efficient batteries and fuel cells.

Electrochemical Cell Components and Structure

Half-cells and Electrodes

• An electrochemical cell consists of two half-cells, each containing an electrode immersed in an electrolyte solution
• The anode is the electrode where oxidation occurs, releasing electrons into the external circuit
• The cathode is the electrode where reduction takes place, accepting electrons from the external circuit
• Examples of common electrodes include zinc (Zn) for the anode and copper (Cu) for the cathode

Electrolyte Solutions and Ion Transfer

• The electrolyte solutions in each half-cell contain the ionic species involved in the redox reactions
• These solutions facilitate the transfer of ions between the half-cells
• The salt bridge or porous membrane connects the half-cells, allowing ion flow without mixing the electrolyte solutions
• The salt bridge maintains electrical neutrality by balancing the charges in the half-cells (KCl or NaNO3)

External Circuit and Current Flow

• The external circuit connects the anode and cathode, allowing electrons to flow from the anode to the cathode
• This flow of electrons generates an electric current that can be used to power devices
• The direction of the current is always from the anode to the cathode in the external circuit
• The magnitude of the current depends on factors such as the electrode surface area and the concentration of the electrolyte solutions

Electrode Potential and Electron Flow

Defining Electrode Potential

• Electrode potential is a measure of the tendency of an electrode to gain or lose electrons in a redox reaction
• It is measured relative to a standard hydrogen electrode (SHE) which has an assigned potential of 0 volts
• The electrode potential depends on the nature of the electrode material and the concentration of the ionic species in the electrolyte solution
• Examples of electrode materials with different potentials include zinc (-0.76 V) and copper (+0.34 V) relative to SHE

Determining the Direction of Electron Flow

• The direction of electron flow in an electrochemical cell is determined by the relative electrode potentials of the anode and cathode
• Electrons flow from the electrode with a more negative potential (anode) to the electrode with a more positive potential (cathode) through the external circuit
• The electrode with a higher tendency to lose electrons (more negative potential) will undergo oxidation
• The electrode with a higher tendency to gain electrons (more positive potential) will undergo reduction
• For example, in a Zn-Cu cell, electrons flow from the zinc anode (-0.76 V) to the copper cathode (+0.34 V)

Cell Potential Calculation

Standard Reduction Potentials

• Standard reduction potentials are tabulated values that indicate the tendency of a species to be reduced under standard conditions (25°C, 1 M concentrations, and 1 atm pressure)
• These potentials are measured relative to the standard hydrogen electrode (SHE)
• To calculate the cell potential, identify the half-reactions occurring at the anode and cathode
• Find their corresponding standard reduction potentials in a table
• Examples of standard reduction potentials include Zn2+ + 2e- → Zn (-0.76 V) and Cu2+ + 2e- → Cu (+0.34 V)

Calculating Standard Cell Potential

• The standard cell potential (E°cell) is the difference between the standard reduction potentials of the cathode (E°cathode) and anode (E°anode) under standard conditions
• The standard cell potential can be calculated using the equation: $E°cell = E°cathode - E°anode$
• Subtract the standard reduction potential of the anode from that of the cathode to obtain the standard cell potential
• A positive E°cell indicates a spontaneous redox reaction, while a negative E°cell indicates a non-spontaneous reaction under standard conditions
• For example, in a Zn-Cu cell: $E°cell = E°cathode(Cu) - E°anode(Zn) = 0.34 V - (-0.76 V) = 1.10 V$

Gibbs Free Energy vs Cell Potential

Gibbs Free Energy and Spontaneity

• Gibbs free energy (ΔG) is a thermodynamic quantity that determines the spontaneity of a chemical reaction, including redox reactions in electrochemical cells
• A negative ΔG indicates a spontaneous reaction, while a positive ΔG indicates a non-spontaneous reaction
• The change in Gibbs free energy is related to the cell potential (Ecell) by the equation: $ΔG = -nFEcell$
• In this equation, n is the number of electrons transferred in the redox reaction, and F is Faraday's constant (96,485 C/mol)

Relationship between Cell Potential and Spontaneity

• When the cell potential is positive (Ecell > 0), the Gibbs free energy change is negative (ΔG < 0), indicating a spontaneous redox reaction
• When the cell potential is negative (Ecell < 0), the Gibbs free energy change is positive (ΔG > 0), indicating a non-spontaneous redox reaction under the given conditions
• The relationship between Gibbs free energy and cell potential allows for the prediction of the spontaneity and direction of redox reactions in electrochemical cells
• For example, a Zn-Cu cell with a cell potential of 1.10 V has a negative ΔG, indicating a spontaneous reaction