David A. Katz Department of Chemistry

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1 Oxidation Reduction and Electrochemistry David A. Katz Department of Chemistry PimaCommunity College

2 Oxidation Reduction Reactions In an oxidation reduction (Redox) reaction, electrons are transferred from one species to another. For example, in a single replacement reaction Cu (s) + 2 AgNO 3 (aq) 2 Ag (s) + Cu(NO 3 3) 2 (aq) The Cu atoms lose electrons to form Cu 2+ in the Cu(NO 3 ) 2 and the Ag + gains electrons to form metallic Ag

3 Oxidation Reduction Reactions This can be more easily observed by writing the net ionic equation for the reaction: Cu (s) + 2 Ag + (aq) 2 Ag (s) + Cu 2+ (aq) The metallic Cu atoms are uncombined, so they are considered to have an oxidation number of zero. The combined Ag atoms are in a +1 oxidation state. Each Cu atom will lose 2 electrons to 2 Ag + ions The resulting Ag atoms are considered to have an oxidation number of zero

4 Oxidation Reduction Reactions Cu (s) + 2 Ag + (aq) 2 Ag (s) + Cu 2+ (aq)

5 Oxidation Numbers In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

6 Oxidation and Reduction A species is oxidized when it loses electrons. Here zinc loses two electrons to go from neutral zinc Here, zinc loses two electrons to go from neutral zinc metal to the Zn 2+ ion.

7 Oxidation and Reduction A species is reduced when it gains electrons. Here each of the H + gains an electron and they Here, each of the H gains an electron and they combine to form H 2.

8 Oxidation and Reduction The species that contains the element that is reduced is the oxidizing agent. H + oxidizes Zn by taking electrons from it. The species that contains the element that is oxidized is the reducing agent. Zn reduces H + by giving it electrons.

9 Assigning Oxidation Numbers 1. Elements intheir elemental form have an oxidation number of The oxidation number of a monatomic ion is the same as its charge.

10 Assigning Oxidation Numbers 3. The oxidation number of metals depends on their position in the periodic table Group IA elements are +1 Group IIA elements are +2 Group IIIA elements are +3 Group IVA metals are usually +2 or +4 Group VA metals tl are usually +3 or +5

11 Assigning Oxidation Numbers 4. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. Oxygen always has an oxidation number of 2, except in the peroxide ion in which it has an oxidation number of 1. 1 Hydrogen is always 1 when bonded to a metal Hydrogen is +1 when bonded to a nonmetal.

12 Assigning Oxidation Numbers 4. Nonmetals (continued). Fluorine always has an oxidation number of 1. The halogens (Cl, Br, and I)have an oxidation number of 1 when they are negative The halogens (Cl, Br, and I) will have positive oxidation numbers in oxyanions (ClO, ClO 2, ClO 3, etc.)

13 Assigning Oxidation Numbers 5. The sumof the oxidation numbers in a neutral compound is The sumof the oxidation numbers in a polyatomic ion is the charge on the ion.

A History of Electricity/Electrochemistry

acquired the property of attracting light

The word electricity comes from "elektron"

Otto von Guericke (1602 1686) invented the

The ball itself was rubbed by hand and the

14 A History of Electricity/Electrochemistry Thales of Miletus ( 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 amber. Otto von Guericke ( ) invented the first electrostatic generator in 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 transportedt to the place where the electric experiment was carried out. Thales of Miletus Otto von Guericke

15 Eventually, a glass globe replaced the sulfur sphere used by Guericke Later, large disks were used

16 Ewald Jürgen von Kleist ( ), 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 ( ) and/or his lawyer friend Andreas Cunnaeus in Leyden/the Netherlands Leyden jars could be joined together to store large electrical charges

17 In 1752, Benjamin Franklin ( ) demonstrated that lightning gwas electricity in his famous kite experiment In 1780, Italian physician and physicist Luigi Aloisio Galvani ( ) 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.

18 Count Alessandro Giuseppe Antonio Anastasio Volta ( ) developed the first electric cell, called a Voltaic Pile, in A voltaic pile consist of alternating layers of two dissimilar i il metals, tl separated tdby pieces of cardboard soaked in a sodium chloride solution or sulfuric acid. Volta determined that the best combination of metalswas zincand silver

In 1800, English chemist William Nicholson (1753 1815) 1815)

separation of water into Soon thereafter, Ritter discovered

of metal deposited and the amount of oxygen produced during an

electrodes Humphrey Davy (1778 1829) utilized the voltaic

which was soon followed by sodium, barium, calcium, strontium,

19 In 1800, English chemist William Nicholson ( ) 1815) and surgeon Anthony Carlisle ( ) separated water into hydrogen and oxygen by electrolysis. Johann Wilhelm l Ritter ( ) 1810) repeated Nicholson s separation of water into hydrogen and oxygen by electrolysis. Soon thereafter, Ritter discovered the process of 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 ( ) utilized the voltaic pile, in 1807, to isolate elemental potassium by electrolysis which was soon followed by sodium, barium, calcium, strontium, magnesium. William Nicholson Johann Wilhelm Ritter Humphrey Davy

20 Michael Faraday ( ) 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. TheSecond Law of Electrochemistry The quantities of different elements deposited by a given amount of electricity are in the ratio of their chemical equivalent weights.

21 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. th hl id f l di l t l t d h l t l d l th t the chloride of lead is an electrolyte, and when electrolyzed evolves the two ions, chlorine and lead, the former being an anion, and the latter a cation.

John Frederic Daniell (1790 1845), professor of

Daniell's research into development of constant

1830s) that commercial telegraph systems began

Daniell's copper battery (1836) became the

These experiments caused serious injury to

22 John Frederic Daniell ( ), 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 standard dfor British and American tl telegraph 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.

23 Voltaic Cells In spontaneous oxidationreduction (redox) reactions, electrons are transferred and energy is released.

24 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.

Cu(NO 3 ) 2 The two solutions are connected by a salt bridge containing NaNO 3

25 Voltaic Cells This is a typical voltaic cell A strip of zinc metal is immersed in a solution of Zn(NO 3 ) 2 A strip of copper metal is immersed in a solution of Cu(NO 3 ) 2 The two solutions are connected by a salt bridge containing NaNO 3 The oxidation oidationoccurs at the anode (Zn) The reduction occurs at the cathode (Cu)

Voltaic Cells To prevent electron flow directly from the zinc to

U shaped tube that contains a salt solution, sealed with porous

charges balanced and forces the electron to move through the

26 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.

27 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 formed dissolve into the solution in the anode compartment.

28 Voltaic Cells As the electrons reach thecathode 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.

29 Electromotive Force (emf) Thepotentialdifference between theanode and cathode in a cell is called the electromotive force (emf). It is also called the cell potential, and is designated E cell.

30 Cell Potential Cell potential is measured in volts (V). 1 V = 1 Where J = Joules C = Coulombs Recall that 1 electron has a charge of 1.6 x C J C

31 Standard Reduction Potentials The cell potential is the difference between two electrode potentials. By convention, electrode potentials are written as reductions Reduction potentials for most common electrodes are tabulated as standard reduction potentials.

32 Standard Cell Potentials The cell potential at standard conditions is calculated E cell = E red (cathode) E red (anode) Substance reduced Substance oxidized Because cell potential is based on the potential energy per unit of charge, it is an intensive property.

33 Cell Potentials Oxidation: E red = V Reduction: E red = V

34 Cell Potentials E cell = E red (cathode) E red (anode) = V ( 0.76 V) = V As a generalization, for most common voltaic cells, the cell As a generalization, for most common voltaic cells, the cell potential (voltage) will be approximately 1.5 V

35 Applications of Oxidation Reduction Reactions

36 Why Study Electrochemistry? Batteries Corrosion Industrial production of chemicals such as Cl 2, NaOH, F 2 and Al Biological redox reactions The heme group

37 BATTERIES Primary, Secondary, and Fuel Cells

38 Batteries Since most batteries only produce 1.5 V, batteries are combined to produce higher voltages

39 Dy Dry Cell Battery Primary battery uses redox reactions that cannot be restored by recharge. Anode (-) Zn Zn e- Cathode (+) 2 NH e- 2 NH 3 + H 2

40 Alkaline Battery Nearly same reactions as in common dry cell, but under basic conditions. Anode ( ): Zn + 2 OH ZnO + H 2 O + 2e Cathode (+): 2 MnO 2 + H 2 O + 2e Mn 2 O OH

41 Alkaline Batteries

42 Lead Storage Battery Secondary battery Uses redox reactions that can be reversed. Canbe restored by recharging

43 Lead Storage Battery Anode (-) E o = V Pb +HSO 4- +H + PbSO 4 +2e- Cathode (+) E o = V PbO +HSO - +3H e- PbSO 4 +2HO 2

44 Ni-Cad Battery Anode (-) Cd + 2 OH - Cd(OH) 2 + 2e- Cathode (+) NiO(OH) + H 2 O + e- Ni(OH) + OH - 2

45 Fuel Cells: H 2 as a Fuel Fuel cell reactants are supplied continuously from anexternal source. Cars can use electricity generated by H 2 /O 2 fuel cells. H 2 carried intanks or generated from hd hydrocarbons.

46 Hydrogen Air Fuel Cell See Figure 20.12

47 Hydrogen Fuel Cells

48 H 2 as a Fuel Comparison of the volumes of substances required to store 4 kg of hydrogen relative to car size.

49 Storing H 2 as a Fuel One way to store H 2 is to adsorb the gas onto a metal or metal alloy.

50 Electrolysis Using electrical energy to produce chemical change. Sn 2+ (aq) + 2 Cl (aq) Sn () (s) + Cl 2(g)

51 Electrolysis of Aqueous NaOH Electric Energy f Chemical Change Anode (+) 4 OH - O 2(g) + 2 H 2 O + 4e- Cathode (-) 4 H 2 O + 4e- 2 H OH - E o for cell = V Anode Cathode

52 Electrolysis Electric Energy f Chemical Change Electrolysis of molten electrons NaCl. Here a battery pumps electrons from Cl to Na +. NOTE: Polarity of electrodes is reversed from batteries. + Anode BATTERY Cl - Na + Cathode

53 Electrolysis of Molten NaCl See Figure 20.18

54 + Electrolysis of Molten NaCl electrons BATTERY Anode (+) 2 Cl - Cl 2 (g) + 2e- Anode Cathode Cathode (-) Cl - Na + Na + + e- Na E o for cell (in water) = E c E a = 271V 2.71 (+1.36 V) = 4.07 V (in water) Et External energy needed db because E o is ( ).

55 Electrolysis of Aqueous NaCl Cells like these are the source of NaOH and Cl 2. In 1995: 25.1 x 10 9 lb Cl 9 2 and 26.1 x 10 lb NaOH Also the source of NaOCl for use in bleach.

56 Electrolysis of Aqueous NaI Anode (+): 2 I - I 2 (g) + 2e- Cathode (-): 2 H 2 O + 2e- H OH - E o for cell = V

57 Electrolysis of Aqueous CuCl 2 Anode (+) 2 Cl - Cl 2 (g) + 2e- electrons Cathode (-) Cu e- Cu + Anode BATTERY Cathode E o for cell = V Note that Cu is more easily reduced than Cl - Cu 2+ H 2 O either H OorNa + 2.

58 Electrolytic Refining of Copper Impure copper is oxidized to Cu 2+ at the anode. The aqueous Cu 2+ ions are reduced to Cu metal at the cathode. The copper formed at the cathode is over 99% pure

59 Producing Aluminum 2 Al 2 O C 4 Al + 3 CO 2 Charles Hall ( ) 1914) developed electrolysis process. Founded Alcoa.

60 Corrosion and

61 Corrosion Prevention The zinc protects the iron from oxidizing (rusting)

David A. Katz Department of Chemistry

An Introduction to Electrochemistry David A. Katz Department of Chemistry PimaCommunity College A History of Electricity/Electrochemistry Thales of Miletus (640 546 B.C.) is credited with the discovery