ELECTROCHEMISTRYSCI3023

http://www.public.coe.edu/departments/ElectrochemEd/

Electrochemistry study of how electricity produces chemical reactions and chemical reactions produces electricity

Involves redox reactionsElectrochemical cell: any device which converts chemical energy into electrical energy or vs.

IntroductionElectrochemistry is a branch of chemistry that studies chemical reactions

which take place in a solution at the interface of an electron conductor (a metal or a semiconductors) and an ionic conductor (the electolyte), and which involve electron transfer between the electrode and the electrolyte or species in solution.

• he field of electrochemistry encompasses a huge array of different phenomena (e.g., electrophoresis and corrosion), devices (electrochromic displays, electro analytical sensors, batteries, and fuel cells), and technologies (the electroplating of metals and the large-scale production of aluminum and• chlorine).

• Alessandro Volta's discovery, in 1793, that electricity could be produced by placing two dissimilar metals on opposite sides of a moistened paper

In 1800, Nicholson and Carlisle, using Volta’s primitive battery as a source, showed that an electric current could decompose water into oxygen and hydrogen.

By 1812, the Swedish chemist Berzelius could propose that all atoms are electrified, hydrogen and the metals being positive, the nonmetals negative.

Humphry Davy prepared the first elemental sodium by electrolysis of a sodium hydroxide melt.

Michael Faraday, to show that there is a direct relation

between the amount of electric charge passed through the solution and the quantity of electrolysis products

Chemical reactions where electrons are transferred between molecules are called oxidation/reduction (redox) reactions. In general, electrochemistry deals with situations where oxidation and reduction reactions are separated in space or time, connected by an external electric circuit to understand each process.

Electron transfer reactions are oxidation-reduction or redox reactions.

Results in the generation of an electric current (electricity) or be

caused by imposing an electric current.

Therefore, this field of chemistry is often called

ELECTROCHEMISTRY.

Electron Transfer Reactions

Terminology for Redox ReactionsOXIDATION :loss of electron(s) by a species; increase in oxidation

number; increase in oxygen.

REDUCTION: Gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.

OXIDIZING AGENT: Electron acceptor; species is reduced.

REDUCING AGENT: Eelectron donor; species is oxidized.

OXIDATION-REDUCTION REACTIONS

Direct Redox ReactionOxidizing and reducing agents in direct contact.

Cu(s) + 2 Ag+(aq) ---> Cu2+(aq) + 2 Ag(s)

• Batteries• Corrosion• Industrial production of

chemicals such as Cl2, NaOH, F2 and Al• Biological redox reactions

The heme group

Why Study Electrochemistry?

Why electrochemistry? According to Sawyer et al. (Electrochemistry for Chemists, 2nd ed.), "..chemical questions amenable to treatment by electochemistry include...." - standard potentials of oxidation-reduction reactions - evaluation of the solution thermodynamics - determination of the electron stoichiometry of oxidation-reduction reactions - evaluation of the heterogeneous electron-transfer kinetics and mechanism - determination of the effect of solvent and electrode material on electron-transfer kinetics - study of reaction and product absorption processes in relation to heterogeneous catalysis - study of pre- and post- chemical reactions associated with the electron-transfer reactions - preparation and study of unstable intermediates - evaluation of the valence of the metal in new compounds - determination of the formulas and stability constants of metal complexes - evaluation of M-X, H-X, and O-Y covalent bond formation energies - studies on the effects of solvent, supporting electrolyte, and solution acidity on oxidation-reduction reactions

Electrochemical Cells and Reactions• In electrochemical systems, we are concerned with the processes and

factors that affect the transport of charge across the interface between chemical phases, for example, between an electronic conductor (an electrode) and an ionic conductor (an electrolyte).• Our concerned with the electrode/electrolyte interface and the

events that occur there when an electric potential is applied and current passes. • Charge is transported through the electrode by the movement of

electrons (and holes).

• Typical electrode materials include solid metals (e.g., Pt, Au), liquid metals (Hg, amalgams), carbon (graphite), and semiconductors (indium-tin oxide, Si). In the electrolyte phase,• charge is carried by the movement of ions. The most frequently

used electrolytes are liquid solutions containing ionic species, such as, H+, Na+, Cl~, in either water or a nonaqueous solvent.

To be useful in an electrochemical cell, the solvent/electrolyte systemmust be of sufficiently low resistance (i.e., sufficiently conductive) for the electrochemical experiment envisioned. Less conventional electrolytes include fused salts (e.g., molten NaCl-KCl eutectic) and ionically conductive polymers (e.g., Nafion, polyethylene oxide-LiClO4). Solid electrolytes also exist (e.g., sodium β-alumina, where charge is carried by mobile sodium ions that move between the aluminum oxide sheets)

Fig. 1. Illustration of electrochemical terms

electrochemical cells. These systems are definedmost generally as two electrodes separated by at least one electrolyte phase.

A difference in electric potential can be measured between the electrodes in an electrochemical cell. Typically this is done with a high impedance voltmeter.

This cell potential, measured in volts (V), where 1 V = 1 joule/coulomb (J/C), is a measure of the energy available to drive charge externally between the electrodes.

The overall chemical reaction taking place in a cell is made up of two independent half-reactions, which describe the real chemical changes at the two electrodes

Each half reaction (and, consequently, the chemical composition of the system near the electrodes) responds to the interfacial potential difference at the corresponding electrode. the electrode at which it occurs is called the working (or indicator) electrode.

To focus on it, one standardizes the other half of the cell by using an electrode (called a reference electrode) made up of phases having essentially constant composition.

The internationally accepted primary reference is the standard hydrogen electrode (SHE), or normal hydrogen electrode (NHE), which has all components at unit activity

Potentials are often measured and quoted with respect to reference electrodes other than the NHE, which is not very convenient from an experimental standpoint. A common reference is the saturated calomel electrode (SCE), which is

Fig.1.1.2 a

Figure 1.1.2; Representation of (a) reduction and (b) oxidation process of a species, A, in solution.

The molecular orbitals (MO) of species A shown are the highest occupied MO and the lowest vacant MO.

These correspond in an approximate way to the E°s of the A/A- and A+/Acouples, respectively.

The illustrated system could represent an aromatic hydrocarbon (e.g., 9,10-diphenylanthracene) in an aprotic solvent (e.g., acetonitrile) at a platinum electrode.

Reduction

Oxidation

Consider a typical electrochemical experiment where a working electrode and a reference electrode are immersed in a solution, and the potential difference between the electrodes is varied by means of an external power supply (Figure 1.1.3).

This variation in potential, E, can produce a current flow in the external circuit, because electrons cross the electrode/solution interfaces as reactions occur.

The number of electrons that cross an interface is related stoichiometrically to the extent of the chemical reaction (i.e., to the amounts of reactant consumed and product generated).

The number of electrons is measured in terms of the total charge, Q, passed in the circuit. Charge is expressed in units of coulombs (C), where 1 С is equivalent to 6.24 X 1018 electrons.

The relationship between charge and amount of product formed is given by Faraday's law; that is, the passage of 96,485.4 С causes 1 equivalent of reaction (e.g., consumption of 1 mole of reactant or production of 1 mole of product in a one-electron reaction).

The current, i, is the rate of flow of coulombs (or electrons), where a current of 1 ampere (A) is equivalent to 1 C/s.

Consider a typical electrochemical experiment where a working electrode and a reference electrode are immersed in a solution, and the potential difference between the electrodes is varied by means of an external power supply (Figure 1.1.3).

Compare this with Fig 1.1.1

The overall chemical reaction taking place in a cell is made up of two independent half-reactions, which describe the real chemical changes at the two electrodes

Equilibrium potential does not exist

When a power supply (e.g., a battery) and a microammeter are connected across the cell, and the potential of the Pt electrode is made more negative with respect to the Ag/AgBr reference electrode. The first electrode reaction that occurs at the Pt is the reduction of protons,