Definition and History of Electrochemistry
– Electrochemistry is the branch of physical chemistry that studies the relationship between electrical potential difference and chemical change.
– It involves reactions where electrons move between electrodes through an electrolyte.
– Electrochemical reactions are driven by an electrical potential difference or result in a potential difference.
– Electrical understanding began in the 16th century with William Gilbert’s experiments on magnetism and electricity.
– Otto von Guericke created the first electric generator in 1663.
– Charles François de Cisternay du Fay proposed the two-fluid theory of electricity in the mid-18th century.
– Luigi Galvani established a connection between chemical reactions and electricity in 1791.
– Alessandro Volta developed the first practical battery in response to Galvani’s findings.
– William Nicholson and Johann Wilhelm Ritter decomposed water into hydrogen and oxygen using electrolysis in 1800.
– Thomas Johann Seebeck discovered thermoelectric currents and the Seebeck effect in the 1810s.
– Humphry Davy’s work with electrolysis led to the isolation of metallic sodium and potassium in 1808.
– Hans Christian Ørsted discovered the magnetic effect of electric currents in 1820.
– Georg Ohm formulated Ohm’s Law in 1827.
– John Daniell invented a primary cell in 1836 that solved the problem of polarization.
– William Grove produced the first fuel cell in 1839.
– Georges Leclanché patented the zinc-carbon cell in 1868.
– Svante Arrhenius published his thesis on electrolytic conductivity in 1884.
– Paul Héroult and Charles M. Hall developed the Hall-Héroult process for obtaining aluminum through electrolysis in 1886.
– Walther Hermann Nernst developed the theory of the electromotive force of the voltaic cell in 1888.
– Friedrich Ostwald conducted important studies on the conductivity and electrolytic dissociation of organic acids in 1894.
– Electrochemistry contributed to the understanding of chemical reactions and the principles of electricity.
– The field continues to advance with new discoveries and applications in various industries.
Principles of Electrochemistry
– Oxidation and reduction are key principles in electrochemistry.
– Redox refers to electrochemical processes involving electron transfer.
– Oxidation is the loss of electrons, while reduction is the gain of electrons.
– The oxidizing agent is always being reduced, and the reducing agent is always being oxidized.
– Oxygen is a common oxidizing agent, but other substances can also act as oxidants.
– Redox stands for reduction-oxidation.
– Oxidation state is the hypothetical charge an atom would have if all bonds were 100% ionic.
– Sodium donates an electron to chlorine, resulting in oxidation and reduction.
– Mnemonic devices like ‘OIL RIG’ and ‘LEO the lion says GER’ can help remember oxidation and reduction.
Balancing Redox Reactions
– Electrochemical reactions in water can be analyzed using the ion-electron method.
– In acidic medium, H+ ions and water are added to balance each half-reaction.
– Balancing redox reactions involves adding electrons, H+ ions, OH- ions, and H2O molecules.
– The ion-electron method allows for the balancing of redox reactions in a systematic way.
– Balancing redox reactions is essential for understanding and predicting chemical reactions.
Electrochemical Cells
– Electrochemical cells produce an electric current from a spontaneous redox reaction.
– Anode is where oxidation occurs, and cathode is where reduction takes place.
– Electrodes can be made from various conductive materials.
– The electrolyte contains ions that can freely move.
– Galvanic cells, such as the Daniell cell, use different metal electrodes in electrolytes to generate electric current.
– An electrochemical cell consists of an anode and a cathode.
– The anode is where oxidation occurs, while the cathode is where reduction occurs.
– The electrolyte allows ion flow while minimizing electrolyte mixing.
– A salt bridge can be used to further minimize electrolyte mixing.
– The cell diagram traces the path of electrons in the cell.
Applications of Electrochemistry
– Fuel cells combine gaseous oxygen and hydrogen to produce water and energy.
– Batteries and fuel cells are based on spontaneous electrochemical reactions.
– Non-spontaneous electrochemical reactions can be driven forward by applying a current.
– The equilibrium constant and Gibbs free energy are related to electrochemical cells.
– The Nernst equation accounts for the effect of reactant concentration on cell potential.
– Corrosion is an electrochemical process that causes rust or tarnish on metals.
– Coinage metals like copper and silver slowly corrode over time.
– Copper forms a green-blue patina of copper carbonate with exposure to water. Source: https://en.wikipedia.org/wiki/Electrochemical
Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically-conducting phase (typically an external electrical circuit, but not necessarily, as in electroless plating) between electrodes separated by an ionically conducting and electronically insulating electrolyte (or ionic species in a solution).
When a chemical reaction is driven by an electrical potential difference, as in electrolysis, or if a potential difference results from a chemical reaction as in an electric battery or fuel cell, it is called an electrochemical reaction. Unlike in other chemical reactions, in electrochemical reactions electrons are not transferred directly between atoms, ions, or molecules, but via the aforementioned electronically-conducting circuit. This phenomenon is what distinguishes an electrochemical reaction from a conventional chemical reaction.
