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Introduction to Electrochemistry

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Introduction to Electrochemistry An Introduction to Electroch emi stry David A. Katz Department of Chemistry Pima Community College A History of Electricity/Electrochemistry • Thales of Miletus (640‐546 B.C.) is credited with the discovery that amber when rubbed with cloth or fur acquired the property of attracting light objects. • The word electricity comes from "elektron" the Greek word for Thales of Miletus Otto von Guericke amber. • Otto von Guericke (1602‐1686) invented the first electrostatic generator in 1675. It was made of a sulphur ball which rotated in a wooden cradle. The ball itself was rubbed by hand and the charged sulphur ball had to be ttdtransported to theplace where the electric experiment was carried out. • Eventually, a glass globe replaced the sulfur sphere used by GikGuericke • Later, large disks were used • Ewald Jürgen von Kleist (1700‐1748), invented the Leyden Jar in 1745 to store electric energy. The Leyden Jar contained water or mercury and was placed onto a metal surface with ground connection. • In 1746, the Leyden jar was independently invented by physicist Pieter van Musschenbroek (1692‐ 1761) and/or his lawyer friend Andreas Cunnaeus in Leyden/the Netherlands • Leyden jars could be joined together to store large electrical charges • In 1752, Benjamin Franklin (1706‐1790) demonstrated that liggghtning was electricity in his famous kite experiment • In 1780, Italian physician and physicist Luigi Aloisio Galvani (1737‐1798) discovered that muscle and nerve cells produce electricity. Whilst dissecting a frog on a table where he had been conducting experiments with static electricity, Galvani touched the exposed sciatic nerve with his scalpel, which had picked up an electric charge. He noticed that the frog’s leg jumped. Count Alessandro Giuseppe Antonio Anastasio Volta (1745 – 1827) developed the first electric cell, called a Voltaic Pile, in 1800. A voltaic pile consist of alternating layers of two dissi mil ar metltals, separatdted by pieces of cardboard soaked in a sodium chloride solution or sulfuric acid. Volta determined that the best combination of metals was zinc and silver • In 1800, English chemist William Nicholson (1753–1815) and surgeon Anthony Carlisle (1768‐1840) separated water into hydrogen and oxygen by electrolysis. • JhJohann Wilhel m Ritter (1776‐1810) repeatdted Nicholson’s separation of water into hydrogen and oxygen by electrolysis. Soon thereafter, Ritter discovered the process of William Nic ho lson electroplating He also observed that the amount of metal deposited and the amount of oxygen produced during an electrolytic process depended on the distance between the electrodes • Humphrey Davy (1778‐1829) utilized the voltaic pile, in 1807, to isolate elemental potassium by electrolysis which was soon followed by sodium, barium, calcium, Johann Wilhelm Humphrey Davy strontium, magnesium. Ritter • Michael Faraday (1791‐1867) began his career in 1813 as Davy's Laboratory Assistant. • In 1834, Faraday developed the two laws of electrochemistry: • The First Law of Electrochemistry The amount of a substance deposited on each electrode of an electrolytic cell is directly proportional to the amount of electricity passing through the cell. • The Second Law of Electrochemistry The quantities of different elements deposited by a given amount of electricity are in the ratio of their chemical equivalent weights. • Faraday also defined a number of terms: The anode is therefore that surface at which the electric current, according to our present expression, enters: it is the negative extremity of the decomposing body; is where oxygen, chlorine, acids, etc., are evolved; and is against or opposite the positive electrode. The cathode is that surface at which the current leaves the decomposing body, and is its positive extremity; the combustible bodies, metals, alkalies, and bases are evolved there, and it is in contact with the negative electrode. Many bodies are decomposed directly by the electric current, their elements being set free; these I propose to call electrolytes.... Finally, I require a term to express those bodies which can pass to the electrodes, or, as they are usually called, the poles. Substances are frequently spoken of as being electro‐negative or electro‐positive, according as they go under the supposed influence of a direct attraction to the positive or negative pole...I propose to distinguish such bodies by calling those anions which go to the anode of the decomposing body; and those passing to the cathode, cations; and when I have occasion to speak of these together, I shall call them ions. …the chlor ide of ldlead is an eltltlectrolyte, and when eltllectrolyzed evolves the two ions, chlorine and lead, the former being an anion, and the latter a cation. • John Frederic Daniell (1790‐1845), professor of chemistry at King's College, London. • Daniell's research into development of constant current cells took place at the same time (late 1830s) that commercial telegraph systems began to appear. Daniell's copper battery (1836) became the stddtandard for Briti sh and AiAmerican tltelegraph systems. • In 1839, Daniell experimented on the fusion of metals with a 70‐ cell battery. He produced an electric arc so rich in ultraviolet rays that it resulted in an instant, artificial sunburn. These experiments caused serious injury to Daniell's eyes as well as the eyes of spectators. • Ultimately, Daniell showed that the ion of the metal, rather than its oxide, carries an electric charge when a metal‐salt solution is electrolyzed. Left: An early Daniell Cell Right:Daniell cells used by Sir William Robert Grove, 1839. Voltaic Cells In spontaneous oxidation‐ reduction (redox) reactions, electrons are transferred and energy is released. Voltaic Cells • If the reaction is separated into two parts, we can use that energy to do work if we make the electrons flow through an external device. • This type of setup is called a voltaic cell. Voltaic Cells • This is a typical voltaic cell • A strip of zinc metal is immersed in a solution of Zn(NO3)2 • A strip of copper metal is immersed in a solution of Cu(NO3)2 • The two solutions are connected by a salt bridge containing NaNO3 • The oidationoxidation occu rs at the anode (Zn) • The reduction occurs at the cathode (Cu) Voltaic Cells • To prevent electron flow directly from the zinc to the copper, a salt bridge is used • The salt bridge consists of a U‐shaped tube that contains a salt solution, sealed with porous plugs, or an agar solution of the salt • The salt bridge keeps the charges balanced and forces the electron to move through the wire – Cations move toward the cathode. – Anions move toward the anode. Voltaic Cells • In the cell, then, electrons leave the anode and flow through the wire to the cathode. • As the electrons leave the anode, the cations fdformed dissolve into the solution in the anode compartment. Voltaic Cells • As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode. • The electrons are taken by the cation, and the neutral metal is deposited on the cathode. Electromotive Force (emf) • The potential difference between the anode and cathode in a cell is called the electromotive force (emf). • It is also called the cell potential, and is designated Ecell. • Cell potential is measured in volts (V). Standard Reduction Potentials The cell potential is the difference between two electrode potentials. By convention, electrode potentials are written as rediductions Reduction potentials for most common electrodes are tabulated as standard reduction potentials. Standard Cell Potentials The cell potential at standard conditions is calculated Ecell = Ered(cathode) − Ered (anode) Substance reduced Substance oxidized Because cell potential is based on the potential energy per unit of charge, it is an intensive property. Cell Potentials Oxidation: E°red = -0.76 V Reduction: E°red = +0.34 V Cell Potentials Ecell = Ered (cathode) − Ered (anode) = +0.34 V − (−0.76 V) = +1.10 V As a generalization, for most common voltaic cells, the cell potential (voltage) will be approximately 1.5 V Applications of Oxidation‐ Reduction Reactions Why Study Electrochemistry? • Batteries • Corrosion • Industrial production of chemicals such as Cl2, NaOH,NaOH, F2 and Al • Biological redox reactions The heme group BATTERIES Primary, Secondary, and Fuel Cells Batteries Since most batteries only produce 1.5 V, batteries are combined to produce higher voltages DDyry Cell Battery Primary battery ——usesuses redox reactions that cannot be restored by recharge. Anode (-) Zn Zn2+ + 2e- Cathode (+) + 2 NH4 + 2e2e-- 2 NH3 + H2 Alkaline Battery Nearly same reactions as in common dry cell, bu t un der bas ic con ditions. ‐ Anode (‐): Zn + 2 OH ZnO + H2O + 2e‐ ‐ Cathode (+): 2 MnO2 + H2O + 2e‐ Mn2O3 + 2 OH Alkaline Batteries Lead Storage Battery • Secondary battery • Uses redox reactions that can be reversed. • Can be restored by recharging Lead Storage Battery Anode ((--)) Eo = +0.36 V - + Pb +HSO+ HSO4 PbSO4 +H+ H +2e+ 2e- Cathode (+) Eo = +1.68 V - + PbO2 +HSO+ HSO4 +3H+ 3 H +2e+ 2e- PbSO4 +2H+ 2 H2O NiNi--CadCad Battery Anode (-(-)) - Cd +2OH+ 2 OH Cd(OH)2 +2e+ 2e- Cathode (+) - NiO(()OH) + H2O + ee-- Ni(()OH)2 + OH Fuel Cells: H2 as a Fuel •Fuel cell ‐ reactants are supplied continuously from an external source. •Cars can use electricity generated by H2/O2 fuel cells. •H2 carried in tanks or generated from hdhydrocarb ons. HydrogenHydrogen——AirAir Fuel Cell See Figure 20.12 Hydrogen Fuel Cells H2 as a Fuel Comparison of the volumes of substances required to store 4 kg of hydrogen relative to car size. Storing H2 as a Fuel One way to store H2 is to adsorb the gas onto a metal or metal alloy. Electrolysis Using electrical energy to produce chemical change. 2+ ‐ Sn (aq) + 2 Cl (aq) Sn()(s) + Cl2(g ) Electrolysis of Aqueous NaOH Electric Energy f Chemical Change Anode (+) Anode Cathode - 4 OH O2(g) + 2 H2O + 4e4e-- Cathode ((--)) - 4 H2O + 4e-4e- 2 H2 + 4 OH Eo for cell = --1.231.23 V Electrolysis Electric Energy f Chemical Change • Electrolysis of molten electrons NaClNaCl.
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