MUHAMMAD FAHAD ANSARI 12 IEEM 14

1. Electrochemical Reactions
By
MUHAMMAD FAHAD ANSARI
12 IEEM 14

2. Electrochemical Reactions
Oxidation-Reduction
Reactions

3. 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:
1- Those that produce electrical energy and
2- 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.

4. Spontaneous Chemical Reactions
Produce 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.

5. Non spontaneous Chemical Reaction
Produced by Electrical-Energy
The 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.

6. 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.

7. 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

8. 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).

9. 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.

10. • 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.

11. Reactions
Reactants Product
2Na + Cl2 → 2 Na+ Cl-
Loses Gains
Oxidized Reduced
Reducing agent Oxidizing agent

12. Electron transfer or oxidation-reduction reactions
Or 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)

13. 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.

14. 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, O20, and so on.
• The oxidation number of a mono-atomic (one-atom) cation
(positive ion) or anion (negative ion) is equal to its charge. In
other words, 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.

15. • 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+O48-, H22+S6+O48-, and so on.
• In complex ions such as SO42-, PO43-, ClO31-, the algebraic sum of
the oxidation numbers of the individual atoms in the ion equals
the charge on the ion.

16. Examples
The oxidation number of S in SO42- , Cl in ClO31- and P in PO43- are
• SO42- x +4(-2) = -2 (net charge on ion)
x -8 = -2
x = 8 -2 = 6
• ClO31- x + 3(-2) = -1 (net charge on ion)
X-6=-1
X=6–1=5
• PO43- x + 4(-2) = -3 (net charge on ion)
X-8=-3
X= 8 – 3 = 5

17. Table-Oxidation States of Some Chemical Elements
Element Oxidation state Species Formula
-2 Hydrogen sulfide H2S
0 Elemental sulfur S
Sulfur
+4 Sulfur dioxide SO2
+6 Sulfate ion SO42-
-4 Methane CH4
0 Soot, graphite C
Carbon
+2 Carbon monoxide CO
+4 Carbon dioxide CO2
-3 Ammonia NH3
0 Nitrogen gas N2
+2 Nitric oxide NO
Nitrogen
+3 Nitrite ion NO2-
+4 Nitrogen dioxide NO2
+5 Nitrate ion NO3-
-2 Almost all compounds -
Oxygen -1 Hydrogen peroxide H2O2
0 Oxygen gas O2
0 Hydrogen gas H2
Hydrogen
+1 Hydrogen ion H+
-1 Chlorine ion Cl-
0 Chlorine gas Cl2
Chlorine
+1 Hypochlorous acid HOCl
+7 Perchloric acid HClO4

18. 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.

19. 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

20. 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.

21. 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.

22. 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+

23. 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.

24. Nitrate Reduction or Denitrification
When 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:

25. 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)

26. 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.

27. 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.

28. 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)
CH2O + 2MnO2 + 3CO2 + H2O ↔ 2Mn2+ + 4HCO3–
reduction
Iron (III) reduction CH2O + 4Fe(OH)3 + 8H+ ↔ 4Fe2+ + 8HCO3– + 3H2O
Sulphate reduction 2CH2O + SO42– + H+ ↔ H2S + 2HCO3–
Methane fermentation 2CH2O ↔ CH4 + CO2

29. Thank you,