oxidation reduction and el h ielectrochemistry to redox 130.pdf · 2008. 12. 1. ·...

OxidationOxidation‐‐ReductionReductionandand

El h iEl h iElectrochemistryElectrochemistry

David A. KatzDepartment of ChemistryPima Community CollegePima Community College

Oxidation‐Reduction ReactionsOxidation Reduction Reactions

In an oxidation‐reduction (Redox) reaction,In an oxidation reduction (Redox) reaction, electrons are transferred from one species to another.

For example, in a single replacement reaction

Cu (s) + 2 AgNO3 (aq) 2 Ag (s) + Cu(NO3)2 (aq)(s) g 3 (aq) g (s) ( 3)2 (aq)

The Cu atoms lose electrons to form Cu2+ in theThe Cu atoms lose electrons to form Cu in theCu(NO3)2 and the Ag+ gains electrons to form metallic Ag

Oxidation‐Reduction ReactionsOxidation Reduction Reactions

• This can be more easily observed by writing the net y y gionic equation for the reaction:

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

• The metallic Cu atoms are uncombined, so they are , yconsidered 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• Each Cu atom will lose 2 electrons to 2 Ag+ ions• The resulting Ag atoms are considered to have an oxidation number of zero

Oxidation‐Reduction ReactionsCu (s) + 2 Ag+ (aq) 2 Ag (s) + Cu2+ (aq)

Oxidation NumbersOxidation Numbers

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

Oxidation and ReductionOxidation and Reduction

• A species is oxidized when it loses electrons.– Here zinc loses two electrons to go from neutral zincHere, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.


• A species is reduced when it gains electrons.– Here each of the H+ gains an electron and theyHere, each of the H gains an electron and they combine to form H2.


• The species that contains the element that is reduced is the oxidizing agent.

H+ idi Z b ki l f i– 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.

Assigning Oxidation NumbersAssigning Oxidation Numbers

1 Elements in their elemental form have an1. Elements in their elemental form have an oxidation number of 0.

2 The oxidation number of a monatomic ion2. The oxidation number of a monatomic ion is the same as its charge.


3 The oxidation number of metals depends3. The oxidation number of metals depends on their position in the periodic table

• Group IA elements are +1• Group IA elements are +1

• Group IIA elements are +2

• Group IIIA elements are +3• Group IIIA elements are +3

• Group IVA metals are usually +2 or +4

G VA t l ll 3 5• Group VA metals are usually +3 or +5


4. Nonmetals tend to have negative4. 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

id ti b f 1an oxidation number of −1.– Hydrogen is always −1 when bonded to a

metalmetal– Hydrogen is +1 when bonded to a

nonmetal.


4. Nonmetals (continued).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‐, ClO2

‐, ClO3‐, etc.)(ClO , ClO2 , ClO3 , etc.)


5. The sum of the oxidation numbers in a5. The sum of the oxidation numbers in a neutral compound is 0.

6 The sum of the oxidation numbers in a6. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

A History of Electricity/Electrochemistry• Thales of Miletus (640‐546 B.C.) is

credited with the discovery thatamber when rubbed with cloth oramber when rubbed with cloth orfur acquired the property ofattracting light objects.

• The word electricity comes fromThe word electricity comes from"elektron" the Greek word foramber.

• Otto von Guericke (1602‐1686)

Thales of Miletus Otto von Guericke

invented the first electrostaticgenerator in 1675. It was made ofa sulphur ball which rotated in awooden cradle. The ball itself wasrubbed by hand and the chargedsulphur ball had to bet t d t th l htransported to the place wherethe electric experiment wascarried out.

• Eventually, a glass globe replaced the lf h d b G i ksulfur sphere used by Guericke

• 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 wascontained water or mercury and was placed onto a metal surface with ground connection.

6 h d j• 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 NetherlandsNetherlands

• Leyden jars could be joined together to store large electrical charges

• In 1752, Benjamin Franklin (1706‐1790) demonstrated that lightning was g gelectricity in his famous kite experiment

• In 1780, Italian physician and physicist Luigi Aloisio Galvani (1737‐1798) g ( )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 noticedpicked up an electric charge. He noticed that the frog’s leg jumped.

Count Alessandro Giuseppe Antonio Anastasio Volta (1745 – 1827) developed the first electricVolta (1745 1827) developed the first electric cell, called a Voltaic Pile, in 1800.

A voltaic pile consist of alternating layers of two di i il t l t d b i fdissimilar metals, separated by pieces of cardboard soaked in a sodium chloride solution or sulfuric acid.

Volta determined that the best combination of metals was zinc andmetals was zinc and silver

• In 1800, English chemist William Nicholson (1753–1815) and surgeon Anthony Carlisle(1753 1815) and surgeon Anthony Carlisle (1768‐1840) separated water into hydrogen and oxygen by electrolysis.

J h Wilh l Ritt (1776 1810) t d• Johann Wilhelm Ritter (1776‐1810) repeated Nicholson’s separation of water into hydrogen and oxygen by electrolysis. Soon thereafter, Ritter discovered the process of

Willi Ni h lp

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

William Nicholson

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, strontium magnesium

Johann Wilhelm Ritter

Humphrey Davystrontium, magnesium. Ritter

• Michael Faraday (1791‐1867) began his career in 1813 as Davy's Laboratory Assistant.y

• In 1834, Faraday developed the two laws of electrochemistry:

• The First Law of Electrochemistry

The amount of a substance deposited on each electrode of anThe 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 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 gour 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 they, q p pelectrodes, 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 cathodethe 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 d i 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 chemistry at King's College, London.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 th t d d f B iti h d A i t l h tthe standard for British and American 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 , pcaused 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 isits oxide, carries an electric charge when a metal‐salt solution is electrolyzed.

Left: An early Daniell Celly

Right:Daniell cells used by Sir William Robert Grove, 1839.,

Voltaic Cells

In spontaneous oxidation‐reduction (redox) reactions, electrons areelectrons are transferred and energy isenergy is released.

Voltaic Cells

• If the reaction is separated into two parts, we can use that energy to dothat energy to do work if we make the electrons flow through an external device.Thi t f t• This type of setup is called a voltaic cell.

Voltaic Cells• This is a typical voltaic

cellf l• A strip of zinc metal is

immersed in a solution of Zn(NO3)2

• A strip of copper metal• A strip of copper metal is immersed in a solution of Cu(NO3)2

• The two solutions are connected by a salt bridge containing NaNO3The o idation occ rs at• The oxidation occurs at the anode (Zn)

• The reduction occurs at the cathode (Cu)the cathode (Cu)

Voltaic Cells

• To prevent electron flow directly from the zinc to thedirectly from the zinc to the copper, a salt bridge is used

• The salt bridge consists of a U‐shaped tube that pcontains a salt solution, sealed with porous plugs, or an agar solution of the saltThe salt brid e keeps the• The salt bridge keeps the charges balanced and forces the electron to move through the wire g– Cations move toward the

cathode.– Anions move toward the

anode.anode.

Voltaic Cells

• In the cell, then, electrons leave theelectrons leave the anode and flow through the wire to the cathodethe cathode.

• As the electrons leave the anode, the

f dcations formed dissolve into the solution in the anode compartment.

Voltaic Cells

• As the electrons reach the cathode cations inthe cathode, cations in the cathode are attracted to the now negative cathodenegative cathode.

• The electrons are taken by the cation, and the

l lneutral metal is deposited on the cathode.

Electromotive Force (emf)Electromotive Force (emf)

• The potential difference between the anodeThe potential difference between the anode and cathode in a cell is called the electromotive force (emf)electromotive force (emf).

• It is also called the cell potential, and is designated Edesignated Ecell.

Cell PotentialCell Potential

Cell potential is measured in volts (V).

J1 V = 1

JC

Where J = JoulesC = Coulombs

Recall that 1 electron has a charge of 1.6 x 10‐19 C

Standard Reduction Potentials

The cell potential is the differenceis the difference between two electrode potentials.By convention, electrode potentials are written as d ireductions

Reduction potentials for most common electrodes are tabulated as standard reduction potentials.p

Standard Cell PotentialsStandard Cell Potentials

The cell potential at standard conditions isThe 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 PotentialsCell Potentials

Ecell° = Ered° (cathode) − Ered° (anode)

= +0.34 V − (−0.76 V)

= +1 10 V= +1.10 V

As a generalization, for most common voltaic cells, the cellAs 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?Why Study Electrochemistry?Why Study Electrochemistry?Why Study Electrochemistry?•• BatteriesBatteries•• CorrosionCorrosion•• CorrosionCorrosion•• Industrial production Industrial production

of chemicalsof chemicals suchsuchof chemicalsof chemicals such such as as ClCl22, , NaOHNaOH, F, F22and Aland Aland Aland Al

•• Biological Biological redoxredoxreactionsreactionsreactionsreactions

The heme groupThe heme group

BATTERIESBATTERIESPrimary, Secondary, and Fuel CellsPrimary, Secondary, and Fuel Cells

BatteriesSince most batteries only produce 1.5 V, batteries are combined to produce higher voltages

Dry Cell BatteryDry Cell BatteryD y yD y yPrimary batteryPrimary battery —— uses redox reactions uses redox reactions that cannot be restored by recharge.that cannot be restored by recharge.

Anode (Anode ( ))

that cannot be restored by recharge.that cannot be restored by recharge.

Anode (Anode (--))

ZnZn ZnZn2+2+ + 2e+ 2e--Zn Zn ZnZn 2e 2e

Cathode (+)Cathode (+)

2 NH2 NH44++ + 2e+ 2e--2 NH2 NH33 + H+ H22

Alkaline BatteryAlkaline BatteryNearly same reactions as in common dry Nearly same reactions as in common dry

ll b t d b i ditill b t d b i diti

yy

cell, but under basic conditions.cell, but under basic conditions.

Anode (Anode (‐‐):): Zn + 2 OHZn + 2 OH‐‐ ZnOZnO + H+ H22O + 2eO + 2e‐‐Cathode (+):Cathode (+): 2 MnO2 MnO + H+ H O + 2eO + 2e‐‐ MnMn OO + 2 OH+ 2 OH‐‐Cathode (+): Cathode (+): 2 MnO2 MnO22 + H+ H22O + 2eO + 2e‐‐ MnMn22OO33 + 2 OH+ 2 OH

Alkaline BatteriesAlkaline Batteries

Lead Storage BatteryLead Storage Battery

•• Secondary Secondary batterybatterybatterybattery

•• Uses Uses redoxredoxreactions thatreactions thatreactions that reactions that can be reversed.can be reversed.

•• Can be restoredCan be restoredCan be restored Can be restored by rechargingby recharging

Lead Storage BatteryLead Storage BatteryAnode (Anode (--)) EEoo = +0.36 V= +0.36 VPbPb + HSO+ HSO44

-- PbSOPbSO44 + H+ H++ + 2e+ 2e--PbPb + HSO+ HSO44 PbSOPbSO44 + H+ H + 2e+ 2eCathode (+) Cathode (+) EEoo = +1.68 V= +1.68 VPbOPbO + HSO+ HSO -- + 3 H+ 3 H++ + 2e+ 2e-- PbSOPbSO + 2 H+ 2 H OOPbOPbO22 + HSO+ HSO44 + 3 H+ 3 H + 2e+ 2e-- PbSOPbSO44 + 2 H+ 2 H22OO

NiNi--Cad BatteryCad BatteryyyAnode (Anode (--))CdCd + 2 OH+ 2 OH-- CdCd(OH)(OH)22 + 2e+ 2e--CdCd + 2 OH+ 2 OH CdCd(OH)(OH)22 + 2e+ 2eCathode (+) Cathode (+) NiONiO(OH) + H(OH) + H22O + eO + e-- Ni(OH)Ni(OH)22 + OH+ OH--( )( ) 22 ( )( )22

Fuel Cells: HFuel Cells: H22 as a Fuelas a FuelFuel Cells: HFuel Cells: H22 as a Fuelas a Fuel••Fuel cellFuel cell ‐‐ reactants are reactants are

supplied continuously from supplied continuously from

an external source.an external source.an external source.an external source.

••Cars can use electricity Cars can use electricity

generated by Hgenerated by H /O/O fuelfuelgenerated by Hgenerated by H22/O/O22 fuel fuel

cells.cells.

••HH carried in tanks orcarried in tanks or••HH22 carried in tanks or carried in tanks or

generated from generated from

h d bh d bhydrocarbons.hydrocarbons.

HydrogenHydrogen——Air Fuel CellAir Fuel Cell

See Figure 20.12See Figure 20.12

Hydrogen Fuel CellsHydrogen Fuel Cells

HH22 as a Fuelas a FuelHH22 as a Fuelas a Fuel

Comparison of the volumes of substances required Comparison of the volumes of substances required to store 4 kg of hydrogen relative to car to store 4 kg of hydrogen relative to car size.size.

Storing HStoring H22 as a Fuelas a FuelStoring HStoring H22 as a Fuelas a Fuel

One way to store HOne way to store H22 is to adsorb the gas onto a is to adsorb the gas onto a metal or metal alloy. metal or metal alloy.

ElectrolysisElectrolysisyyUsing electrical energy to produce chemical change.

Sn2+( ) + 2 Cl‐( ) Sn( ) + Cl2( )Sn (aq) + 2 Cl (aq) Sn(s) + Cl2(g)

Electrolysis of Aqueous NaOHElectrolysis of Aqueous NaOH

Anode (+)Anode (+)

Electric Energy Electric Energy ffChemical ChangeChemical Change

AnodeAnode CathodeCathodeAnode (+)Anode (+)4 OH4 OH-- OO2(g2(g) ) + 2 H+ 2 H22O + 4eO + 4e--

AnodeAnode CathodeCathode

Cathode (Cathode (--) ) 4 H4 H22O O + 4e+ 4e-- 2 2 HH22 + 4 + 4 OHOH--

EEoo for cell = for cell = --1.23 V1.23 V

ElectrolysisElectrolysisyyElectric Energy Electric Energy ff Chemical ChangeChemical Change

electronselectrons

BATTERY

electrons

BATTERY

electrons•• Electrolysis of molten Electrolysis of molten

NaClNaCl..

•• Here a battery “pumps”Here a battery “pumps” BATTERY

+

A d

BATTERY

+

A d

•• Here a battery “pumps” Here a battery “pumps”

electrons from electrons from ClCl‐‐ to Nato Na++..

••NOTE:NOTE: Polarity ofPolarity of Anode CathodeAnode CathodeNOTE:NOTE: Polarity of Polarity of

electrodes is reversed electrodes is reversed

from batteries.from batteries.

Na+Cl- Na+Cl-

Electrolysis of Molten NaClElectrolysis of Molten NaClyy

See Figure 20.18See Figure 20.18

Electrolysis of Molten NaClElectrolysis of Molten NaClElectrolysis of Molten NaClElectrolysis of Molten NaClAnode (+) Anode (+) electronselectrons

2 2 ClCl-- ClCl22(g) + 2e(g) + 2e--Cathode (Cathode (--) )

BATTERY

+

Anode Cathode

BATTERY

+

Anode Cathode (( ))

NaNa++ + e+ e-- NaNaNa+Cl- Na+Cl-

EEoo for cell (in water) = E˚for cell (in water) = E˚cc ‐‐ E˚E˚aa2 71 V2 71 V (+1 36 V)(+1 36 V)= = ‐‐ 2.71 V 2.71 V –– (+1.36 V)(+1.36 V)

= = ‐‐ 4.07 V (in water)4.07 V (in water)

E t l d d b EE t l d d b Eoo i (i ( ))External energy needed because EExternal energy needed because Eoo is (is (‐‐). ).

Electrolysis of Aqueous NaClElectrolysis of Aqueous NaCly qy qCells like these are the source of NaOH and ClCells like these are the source of NaOH and Cl22..In 1995: 25 1 x 10In 1995: 25 1 x 1099 lb Cllb Cl and 26 1 x 10and 26 1 x 1099 lb NaOHlb NaOHIn 1995: 25.1 x 10In 1995: 25.1 x 10 lb Cllb Cl22 and 26.1 x 10and 26.1 x 10 lb NaOHlb NaOH

Also the source of NaOCl for use in bleach.Also the source of NaOCl for use in bleach.

Electrolysis of Aqueous NaIElectrolysis of Aqueous NaI

Anode (+):Anode (+): 2 I2 I-- II22(g) + 2e(g) + 2e--Cathode (Cathode (--): ): 2 H2 H22O + 2eO + 2e-- HH22 + 2 OH+ 2 OH--


Electrolysis of Aqueous CuClElectrolysis of Aqueous CuCl22Anode (+) Anode (+)

electronselectrons2 2 ClCl-- ClCl22(g) + 2e(g) + 2e--

Cathode (Cathode (--) ) BATTERYBATTERY

(( ))

CuCu2+2+ + 2e+ 2e-- CuCu+

Anode Cathode

+

Anode Cathode


Note that Cu is more Note that Cu is more Cu2+Cl-

H2OCu2+Cl-

H2Oeasily easily reduced than reduced than either Heither H22O or NaO or Na++

H2OH2O

either Heither H22O or NaO or Na . .

Electrolytic Refining of CopperElectrolytic Refining of CopperElectrolytic Refining of CopperElectrolytic Refining of Copper

Impure copper is oxidized to CuImpure copper is oxidized to Cu2+2+ at the anode. The aqueous at the anode. The aqueous CuCu2+2+ ions are reduced to Cu metal at the cathodeions are reduced to Cu metal at the cathode..The copper formed at the cathode is over 99% pureThe copper formed at the cathode is over 99% pure

Producing AluminumProducing Aluminumgg2 Al2 Al22OO33 + 3 C + 3 C 4 Al + 3 CO4 Al + 3 CO22

Charles Hall (1863Charles Hall (1863 1914) developed electrolysis process1914) developed electrolysis processCharles Hall (1863Charles Hall (1863‐‐1914) developed electrolysis process. 1914) developed electrolysis process. Founded Alcoa.Founded Alcoa.

Corrosion and…Corrosion and…

…Corrosion Prevention…Corrosion Prevention

The zinc protects the iron from oxidizing (rusting)

oxidation reduction and el h ielectrochemistry to redox 130.pdf · 2008. 12. 1. ·...

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