Electrochemistry

• Muhammad Fahad Ansari

• 12IEEM14

Electrochemistry

• One of the major classes of chemical reactions involves an electron exchange between the reactants.

• These reactions are called oxidation-reduction reactions or electrochemical, reactions. they can be further subdivided into two types or classes:Reactions those that produce electrical energy and those that are produced by electrical energy.One-type of reaction is spontaneous and may be used to produce electricity;The-second must be forced to occur and requires electricity to make it take place

Spontaneous Chemical ReactionsProduce Electrical-Energy

For example:• The electrical energy used to start a car is

produced by an oxidation-reduction reaction in the car’s battery.

• Similarly, the electricity that causes the light in a flash light comes from electrochemical reactions in the dry cell of the flashlight.

Non spontaneous Chemical Reaction Produced by Electrical-Energy

The type of electrochemical reactions require electrical energy, or the use of electrical energy to bring about chemical reactions, which is called electrolysis.Electroplating is another important application in which electricity is used to bring about a chemical reaction.

Electrolysis

Many of the active elements (those with strong tendency to lose or gain electrons) are available only as the result of electrolysis.

For example: Na and Cl found in nature only in the form of their ions (Na+, Cl-) in compounds.

• Electricity is used to take an electron away from the Cl- ion and their by re-form the atom, (Cl).

• Similarly, electricity is required to force the Na+ to gain an electron to become a sodium (Na) again.

Electroplating

• Electroplating involves electrolysis that is carried out so that the ions of metal are converted into atoms on the surface to be plated.

• Cu (s) Cu (aq) + 2e- (half equation/reaction)• Ag+ (aq) + e- Ag (s) (half equation/reaction)

• Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

• Cu(s)RA + 2Ag+(aq)OA <===> Cu2+(aq)OA + 2Ag(s)RA

Redox Reactions

• Some common redox reactions are those that occur in batteries, when metals rust, when metals are plated from solutions, and

• Combustion of organic molecules such as hydrocarbons (like methane and gasoline) and carbohydrates (like wood).

Basic Oxidation-Reduction Reactions Concepts

Chemical reactions involves an electron exchange between the reactants, in which one substance loses electrons and another substance gains electrons.

• Oxidation: is the loss of electrons by atoms, ions, or molecules.

• Reduction: is the gain of electrons by atoms, ions, or molecules.

• Oxidizing agent: Substance that causes another substance to lose electrons

• Reducing agent: Substance that loses electrons in a chemical reaction; it reduces another substance by losing electrons to it.

Reactions

Reactants Product

2Na + Cl2 → 2 Na+ Cl-

Loses GainsOxidized ReducedReducing agent Oxidizing agent

Electron transfer or oxidation/reduction reactionsOr transfer of charge - an electron - from one

species to another

• 2Mg(s) + O2(g) 2MgO(s)

• Fe(s) + O2(g) Fe oxides (s)

• C(s) + O2(g) CO2(g)

CH4 + O2 H2O + CO2

Since H has an oxidation # of 1+, the oxidation # of C in CH4 is 4-,

while in CO2 it is 4+.

Clearly C has been oxidized by the oxidizing agent O2.

O2 has been reduced to form both products.

CH4 + O2 H2O + CO2

the oxidation numbers for the central C in each molecule are 4-, 2-, 0, 2+, and 4+ as you proceed from left to right

Step-wise oxidations of the carbon.

Oxidation Number or Oxidation State

Simple and Arbitrary Rules:• The oxidation number of an element in its free or

uncombined form is zero. For example: Na0, Mg0, S0, O2

0, and so on.

• The oxidation number of mono-atomic ions equals the number of electrons it has lost or gained.For example: Na1+, Mg2+, Al3+, Cl1-, S2-, and so on.

• The oxidation number of oxygen in compounds is usually -2. The exceptions are the peroxides, such as H2O2, and compounds of oxygen and fluorine, such as OF2.

• The oxidation number of hydrogen in compounds is usually +1. The hydrides, such as NaH, are exceptions.

• In the formula for a compound, the sum of the positive oxidation numbers must equal the sum of the negative oxidation numbers.For example: Mg2+S2-, Na1+Cl1-, K1+Mn7+O4

8-, H22+S6+O4

8-, and so on.

• In complex ions such as SO42-, PO4

3-, ClO31-, the algebraic sum of

the oxidation numbers of the individual atoms in the ion equals the charge on the ion.

Examples

The oxidation number of S in SO42- , Cl in ClO3

1- and P in PO43- are

• SO42- x +4(-2) = -2 (net charge on ion)

x -8 = -2x = 8 -2 = 6

• ClO31- x + 3(-2) = -1 (net charge on ion)

X - 6 = - 1X = 6 – 1 = 5

• PO43- x + 4(-2) = -3 (net charge on ion)

X - 8 = - 3 X= 8 – 3 = 5

Table-Oxidation States of Some Chemical ElementsElement Oxidation state Species Formula

Sulfur

-2 Hydrogen sulfide H2S

0 Elemental sulfur S+4 Sulfur dioxide SO2

+6 Sulfate ion SO42-

Carbon-4 Methane CH4

0 Soot, graphite C+2 Carbon monoxide CO+4 Carbon dioxide CO2

Nitrogen

-3 Ammonia NH3

0 Nitrogen gas N2

+2 Nitric oxide NO+3 Nitrite ion NO2

-

+4 Nitrogen dioxide NO2

+5 Nitrate ion NO3-

Oxygen-2 Almost all compounds --1 Hydrogen peroxide H2O2

0 Oxygen gas O2

Hydrogen0 Hydrogen gas H2

+1 Hydrogen ion H+

Chlorine-1 Chlorine ion Cl-

0 Chlorine gas Cl2

+1 Hypochlorous acid HOCl+7 Perchloric acid HClO4

Microbial Redox Process

• Important redox reactions that are carried out by microorganisms are summarized here:

The notation [CH2O] is used to denote a fragment of an arbitrary carbohydrate.

Photosynthetic Production of Biomass

• Photosynthetic microorganisms (algae and some bacteria) carryout photosynthesis reactions, in these reactions, energy-rich carbohydrate molecules are produced by combining carbon dioxide and water, using energy derived from sunlight.

• From a Redox Perspective:

CO2 + H2O [CH2O] + O2

Carbon is reduced from oxidation state +4 to 0, and oxygen is oxidized from -2 to 0

Aerobic Respiration:In the presence of oxygen, microorganisms degrade biomass to form carbon dioxide and water. Chemical energy that is released can be used by the organisms.

[CH2O] + O2 CO2 + H2O

This process is the reverse of photosynthesis, carbon is oxidized and oxygen is reduced.

Sources of Nitrogen

• Atmosphere N2 gas• Water NH3, NH4, NO2, NO3 ions• Soil NH4+ , NO2, NO3 ions• Plants Proteins contains

(-NH2) Amino group(O=C-OH) Carboxyl group

• Animals Proteins, Urea, and AminesUrea contains (NH2-C=O-NH2)

Nitrogen Fixation• In the atmosphere, nitrogen is almost entirely in the

form of N2 and is in oxidation state 0.

• The nitrogen in biological system is mostly in the form of an amine –NH2, Which is very closely related to ammonia (NH3) and ammonium ion (NH4

+) here nitrogen is in oxidation state -3.

• Nitrogen in water and soil is in form of nitrate (NO3-)

in which nitrogen is in oxidation state +5.

Microorganisms play an essential role in the movement of nitrogen among these oxidation states.

Compounds such as ammonia and nitrate contain a single nitrogen atom as fixed nitrogen species.

Certain groups of bacteria are capable of converting gaseous nitrogen to fixed nitrogen, in the form of the ammonium ion.

Energy from the oxidation of biomass to CO2 is used to reduce the nitrogen in N2 to ammonium.

3[CH2O] + 2N2 + 3H2O + 4H+ 3CO2 + 4NH4+

Nitrification In the nitrification, nitrogen in the ammonium ion is oxidized from -3 to +5, with oxygen as oxidizer.

NH4+ + 2O2 NO3

- + 2H+ + H2O

Plants absorb nitrogen more efficiently in the form of nitrate than an ammonium, so redox reaction can enhance the effectiveness of ammonia-based agricultural fertilizers.

Nitrate Reduction or DenitrificationWhen oxygen is not available as the oxidizer to degrade biomass, microorganisms can use nitrate as the oxidizer (electron acceptor).

Nitrate Reduction is used in some wastewater treatment systems to convert fixed nitrogen to N2 gas, which can then be safely released to the atmosphere. This process is called denitrification, since nitrogen is removed from the aqueous system.

Nitrogen in municipal wastewater begins in a reduced state (-3), the overall process involved two steps:

Nitrification in an aerobic reactor, followed by denitrification in an anaerobic reactor, four nitrogen atoms, being reduced from +5 to 0, can fully oxidize five carbon atoms from 0 to +4

NH4+ + 2O2 NO3

- + 2H+ + H2O (aerobic reaction)

5[CH2O] + 4NO3- + 4H+ 5CO2 + 7H2O + 2N2

(anaerobic reaction)

Sulfate Reduction• Some environments that contain biodegradable materials

lake both oxygen and nitrate to serve as the oxidizing agent, in such cases, sulfate may serve that role.

• The conversion of one sulfur atom from +6 in sulfate to -2, in hydrogen sulfide oxidizes two carbon atoms from 0 to +4 oxidation states.

2[CH2O] + 2H+ + SO42- 2CO2 + 2H2O + H2S

• This reaction can occur in stagnant anaerobic marine sediments that are supplied with decaying biomass, algae or seaweed accumulation.

Methane Formation (Methanogenesis)• In the absence of oxygen, nitrate, and sulfate, biomass can

still be converted to carbon dioxide as:

2[CH2O] CO2 + CH4

• This is an interesting redox reaction, since the two carbon atoms begins in oxidation state zero (0). One carbon atom is oxidized to +4, and the other is reduced to -4.

• Methane generation process is exploited in seawater treatment to convert excess microbiological material to gases, which are more easily handled for disposal.

Sequence Of Redox Reactions (CH2O = Unidentified Organic Matter With Zero-valent Carbon)

Aerobic respiration CH2O + O2 ↔ CO2 + H2O

Denitrification 5CH2O + 4NO3– ↔ 2N2 + 4HCO3– + CO2 + 3H2O

Manganese (IV) reduction

CH2O + 2MnO2 + 3CO2 + H2O ↔ 2Mn2+ + 4HCO3–

Iron (III) reduction CH2O + 4Fe(OH)3 + 8H+ ↔ 4Fe2+ + 8HCO3– + 3H2O

Sulphate reduction 2CH2O + SO42– + H+ ↔ H2S + 2HCO3

–

Methane fermentation 2CH2O ↔ CH4 + CO2

Thank you,well come for questions