Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry ↓


Electrochemistry


Electrochemistry is a branch of chemistry that studies the relationship between electricity and chemical reactions. It includes the processes by which chemical energy is converted into electrical energy and vice versa. This field is fundamental to many applications, including batteries, electroplating, and electrolysis.

Basic concepts

To understand electrochemistry we need to define some basic concepts:

  • Oxidation: This is the process in which a chemical species loses electrons. An increase in oxidation state is usually associated with this.
  • Reduction: This is the process in which a chemical species gains electrons, usually with a decrease in oxidation state.
  • Redox reaction: Short form of reduction-oxidation reaction, where both processes occur simultaneously.
  • Oxidizing agent: A substance that gains electrons and is reduced in a chemical reaction.
  • Reducing agent: A substance that loses electrons in a chemical reaction and becomes oxidized.

Electrochemical cells

An electrochemical cell is a system that uses chemical reactions to generate electrical energy, or uses electrical energy to drive chemical reactions. There are two types of electrochemical cells:

  1. Galvanic (or voltaic) cells
  2. Electrolytic cell

Galvanic cells

Galvanic cells derive energy from spontaneous redox reactions that occur within the cell. Here is an example setup of a galvanic cell:

Zn Anode Cu Cathode Zinc Cube salt bridge

For example, in a zinc-copper galvanic cell, zinc loses electrons at the anode:

Zn → Zn²⁺ + 2e⁻

The electrons flow through an external circuit to the copper cathode, where they react with copper ions in the solution:

Cu²⁺ + 2e⁻ → Cu

As these reactions continue, electrons flow through the circuit, providing electrical power. The salt bridge works to maintain electrical neutrality by allowing the exchange of ions.

Electrolytic cell

Unlike a galvanic cell, an electrolytic cell requires an external source of energy to drive chemical reactions. These cells are used to induce non-spontaneous reactions. A typical application of an electrolytic cell is the electrolysis of water.

Negative Electrode Positive Electrode H₂ O₂

In the electrolysis of water, an external voltage is applied and water is decomposed into hydrogen and oxygen gases:

2H₂O(l) → 2H₂(g) + O₂(g)

The reactions taking place at the electrodes are as follows:

  • Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻
  • Anode: 2H₂O → O₂ + 4H⁺ + 4e⁻

Applications of electrochemistry

Understanding electrochemistry is important for many practical applications, including:

  • Batteries: Electrochemical cells used to store electricity. Different types of batteries have different chemical combinations, such as lead-acid, lithium-ion, and nickel-cadmium.
  • Corrosion prevention: Metals can corrode when exposed to certain environments. Electrochemical methods can prevent or slow this process, such as in galvanization where a protective zinc layer is applied.
  • Electroplating: A thin layer of metal is deposited on the surface of a substrate. This is typically used for decorative purposes, protection against corrosion, or to improve electrical conductivity.

Nernst equation

The Nernst equation describes how the concentration of ions affects the potential of an electrochemical cell. It is expressed as:

E = E⁰ - (RT/nF) * ln(Q)

Where:

  • E is the cell potential.
  • E⁰ is the standard cell potential.
  • R is the universal gas constant.
  • T is the temperature in Kelvin.
  • n is the number of moles of electrons.
  • F is the Faraday constant.
  • Q is the reaction quotient.

The Nernst equation allows chemists to calculate cell potentials under non-standard conditions, taking into account the different concentrations and pressures involved in electrochemical reactions.

Conclusion

Electrochemistry is an important field of chemistry that has wide applications in the modern world. From the batteries that power our electronic devices to the industrial processes that produce the metals we use every day, electrochemistry plays a vital role. Understanding the basic principles of redox reactions and the operation of electrochemical cells helps us harness the power of chemical reactions to efficiently generate and use electrical energy.


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Electrochemistry

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