EC Chapter
Introduction to Electrochemistry
Definition
Electrochemistry is the study of the relationship between electricity and chemical reactions. It focuses on two main processes: the generation of electrical energy from spontaneous chemical reactions and the use of electrical energy to drive non-spontaneous chemical reactions. Understanding electrochemistry is crucial for various applications, including energy production, battery technologies, and environmental management.
Applications
Chemical Production: Electrochemical processes are used in the industrial production of essential chemicals like sodium hydroxide, chlorine, fluorine, and other compounds.
Energy Storage: Batteries and fuel cells store chemical energy and convert it into
Learning Objectives
Describe Electrochemical Cells: Understand the two main types: galvanic (voltaic) cells that produce energy from spontaneous reactions, and electrolytic cells that require energy input.
Apply Nernst Equation: Use the Nernst equation to calculate the electromotive force (emf) of galvanic cells under varying conditions.
Relate Thermodynamic Concepts: Connect Gibbs free energy, standard electrode potentials, and equilibrium constants in electrochemical reactions.
Differentiate Conductive Properties: Define and compare resistivity, conductivity, and molar conductivity, and understand their significance in electrochemistry.
Measure Conductivity: Learn to measure the conductivity of electrolytic solutions and observe how molar conductivity changes with concentration.
Explore Kohlrausch's Law: Understand Kohlrausch's Law and its implications for conductivity and electrolysis applications.
Quantitative Electrochemistry: Grasp the quantitative aspects of electrolysis, battery construction, and fuel cell operation.
Electrochemical Corrosion: Explain the electrochemical process of corrosion and explore methods for prevention.
Electrochemical Cells
Galvanic Cells
Function: Convert chemical energy from spontaneous redox reactions into electrical energy that can be harnessed for work, such as powering electronic devices.
For this, the Gibbs Energy of the spontaneous reaction is usually converted into electrical energy
Example: Daniell Cell, where zinc and copper electrodes are used.
Reactions: In a typical galvanic cell:
Oxidation: Zn(s) ➔ Zn²⁺(aq) + 2e⁻ (Zinc loses electrons)
Reduction: Cu²⁺(aq) + 2e⁻ ➔ Cu(s) (Copper ions gain electrons)
Each of these are known as half reactons, with the copper being the reduction half reaction and the sinz being the oxidation half reaction. these two together form the redox couple.
Potential: The cell's standard potential is 1.1 V under standard conditions (1 mol/dm³ concentration).
Standard Electrode Potential -
Making the Electrodes: We dip each metallic electrode into the electrolyte and connect them with a metallic wire for the transfer of electrons. This is external.We also connect them with a salt bridge internally so that the flow of ions can be maintained. this allows it to be elevtrically neutral so thats good. Salt bridge is made of inert electrolytes
Understanding Current Flow: Current flows from the cathode (positive electrode) to anode (negative electrode), while electrons flow in the opposite direction from anode to cathode.
Effects of External Voltage:
Eext < 1.1 V: Electron flow occurs, driving the reaction forward (discharging).
Eext = 1.1 V: The system reaches equilibrium, and there is no net charge flow.
Eext > 1.1 V: The reaction reverses direction and behaves like an electrolytic cell, requiring external energy.
Lay mans explanation:
Why does it even do anything?
We get the cell like that. With Zinc Sulphate as the Anode and Copper Sulphate as the cathode. Zinc undergoes oxidation which is the process of losing electrons, while copper undergoes reduction, gaining those electrons to form metallic copper. We can think of connecting the two as opening a door between the two. When this door is opened, zinc gives away electrons because it isnt as electronegative meaning it doesnt hold onto its electrons well wherereas copper is, meaning itll pull those electrons towards itself when the door is opened. Zinc is also kinda wanting to lose those electrons to become stable. So the electrons are wanting to move from where they arent wanted, so at zinc, to a place thats more chill with them being there, like copper.
Explanation of electrode potential:
electrode potenial is essentially an indicator of what is happening at the electrode and thus what its charge is and also how likely it is to do what. Positive EP means it is likely to take electrons and deposit the metal and thus cathodes have positive ep, whereas anodes are more likely to disolve as ions into the soluion and leave behind electrons and thus their EP is negative. For exaple, for something to have a negative electron potential, it is more negative than positive. This happens because more of the electrode is forming metal ions in the solution and leaving eletcrons behind, thus making it a negatively charged region. At the cathode, the reverse is ocuring, i.e more of the metal ions are being deposited at the metal electrode by taking electrons than the reverse, making it positively charged due to lack of electrons in the region.
Electrolytic Cells
Function: Use external voltage to drive non-spontaneous chemical reactions, which would not occur without a power source.
Combination of Reactions: In this cell, reactions are forced to occur, such as the oxidation of Cu(s) while Cu²⁺ ions discharge at the cathode to form copper metal.
Electrode Potentials and Standard Reactions
Cell Potential: A measure of the electromotive force (emf) of an electrochemical cell, measured in volts. It indicates the maximum potential difference available.
Standard Electrode Potential: Defined under standard conditions (1 M concentration, 1 atm pressure). It reflects the tendency of a chemical species to be reduced.
Nernst Equation: Relates emf to concentration, providing a way to calculate cell potentials under non-standard conditions. It shows how potential changes as concentrations change.
Electrode Potential Measurement: The standard hydrogen electrode (SHE) is used as a reference point with a potential of E° = 0 V at all temperatures.
Example Cell Potentials:
Cu²⁺ + 2e⁻ ➔ Cu(s): E° = 0.34 V
Zn²⁺ + 2e⁻ ➔ Zn(s): E° = -0.76 V
Conductance in Electrolytic Solutions
Definitions
Conductivity (κ): A measure of a solution's ability to conduct electricity, typically expressed in siemens per meter (S/m).
Resistivity (ρ): The opposition a material offers to the flow of electric current, inversely related to conductivity. Lower resistivity means higher conductivity.
Molar Conductivity (Λm): Defined as κ/c (where c is the concentration), showing how conductivity varies with concentration. It provides insight into how ions behave in solution.
Measurement Techniques
Resistance Measurement: Resistance is measured using a Wheatstone bridge, which accounts for the geometry of the cell and leads to accurate conductivity values.
Conductivity Cells: Platinum electrodes are commonly used to measure resistance in a solution to calculate conductivity (κ).
Variation with Concentration
Strong Electrolytes: Typically show a slow increase in molar conductivity with dilution, indicating partial dissociation.
Weak Electrolytes: Generally demonstrate a steep increase in molar conductivity with dilution as they dissociate more effectively when concentrated.
Faraday's Laws of Electrolysis
First Law: The mass of a substance deposited or liberated in an electrochemical reaction is directly proportional to the charge passed through the cell. This means that more current yields more product.
Second Law: The masses of different substances deposited by the same quantity of electricity are proportional to their equivalent weights. This helps us understand the relationships between different ions.
Battery Technology
Primary Batteries
Example: Dry cell, which is a type of primary battery that can only be used until the chemicals are exhausted. They cannot be recharged and are often used in everyday electronic devices.
Secondary Batteries
Example: Lead-acid battery, which is rechargeable. This technology allows the chemical reactions that occur during discharging to be reversed during charging, allowing the battery to be reused multiple times.
Fuel Cells
Definition: Fuel cells convert chemical energy directly into electrical energy, typically using hydrogen and oxygen as reactants. This process is efficient and clean, producing water as a byproduct.
Efficiency: Fuel cells have a higher efficiency (about 70%) compared to traditional thermal power plants, which are less efficient due to energy losses.
Components: Fuel cells rely on catalysts to facilitate reactions at the electrodes, improving overall efficiency and performance.
Overall Reaction: 2H₂ + O₂ ➔ 2H₂O, illustrating how the fuel cell processes hydrogen and oxygen to generate electricity and produce water.
Corrosion Processes
Electrochemical Nature: Corrosion is an electrochemical process where metals oxidize due to reactions with their environment, leading to degradation and loss of material. It represents a significant challenge for many metal structures and equipment.
Prevention: Methods to prevent corrosion include physical barriers (like paints), sacrificial anodes that corrode instead of the main structure, and protective coatings that shield metals from exposure to corrosive elements.
Conclusion
Electrochemistry is a fascinating field that combines extensive knowledge of chemistry and electricity with practical applications in energy, technology, and environmental science. Through the study of electrochemical cells, battery technology, and corrosion prevention, we can better understand how to harness and utilize energy efficiently while minimizing environmental impacts.