Cell potential is a fundamental concept in electrochemistry, representing the driving force behind an electrochemical reaction. It quantifies the voltage an electrochemical cell can produce, which is the electrical energy available from its chemical processes. This potential difference drives the flow of electrons through an external circuit, enabling devices like batteries to generate electricity.
Understanding the Electrochemical Cell
Understanding the basic structure of an electrochemical cell is helpful. A cell consists of two half-cells, each containing an electrode submerged in an electrolyte solution. These half-cells are connected externally by a wire and internally by a salt bridge.
In an electrochemical cell, chemical reactions involve electron transfer, known as redox reactions. Oxidation, the loss of electrons, occurs at the anode, while reduction, the gain of electrons, takes place at the cathode. The salt bridge maintains electrical neutrality within the half-cells by allowing ions to migrate, ensuring a continuous flow of charge.
Calculating Potential Under Standard Conditions
The standard cell potential ($E^\circ_{cell}$) refers to the voltage generated by an electrochemical cell under specific standard conditions. These conditions are defined as 25°C (298 K), 1 M concentration for all dissolved species, and 1 atmosphere partial pressure for any gases. Standard electrode potentials are listed in tables, measured against a standard hydrogen electrode (SHE), defined as 0 V.
To calculate the standard cell potential, use the standard reduction potentials of the two half-reactions. The formula subtracts the standard reduction potential of the anode from that of the cathode: $E^\circ_{cell} = E^\circ_{cathode} – E^\circ_{anode}$. The half-reaction with the more positive standard reduction potential undergoes reduction at the cathode, while the other is oxidized at the anode. A positive $E^\circ_{cell}$ indicates a spontaneous reaction under standard conditions.
Calculating Potential Under Non-Standard Conditions
Electrochemical reactions rarely occur under standard conditions, as concentrations and temperatures often vary. To determine cell potential under non-standard conditions, the Nernst equation is employed. This equation accounts for the influence of reactant and product concentrations, as well as temperature, on the cell’s potential.
The Nernst equation is given as $E_{cell} = E^\circ_{cell} – (RT/nF) \ln Q$. Here, $E_{cell}$ is the cell potential under non-standard conditions, and $E^\circ_{cell}$ is the standard cell potential. Variables include the universal gas constant (R, 8.314 J/(mol·K)), temperature in Kelvin (T), moles of electrons transferred (n), and Faraday’s constant (F, approximately 96,485 C/mol e⁻). The term Q represents the reaction quotient, considering concentrations or partial pressures of reactants and products.
Practical Measurement of Cell Potential
Beyond theoretical calculations, cell potential can be directly measured in a laboratory setting. This is achieved using a voltmeter, which measures the potential difference between two points. In an electrochemical cell setup, the voltmeter’s leads are connected to the two electrodes.
When connected, the voltmeter measures the voltage generated by electron flow from the anode to the cathode. For an accurate reading, ensure a positive voltage is displayed, which indicates the spontaneous direction of the reaction. This practical measurement provides the cell’s electrical output under the specific experimental conditions.