Topic#6.
Redox reaction
chemistry
Redox reaction
Oxidation–reduction reactions (redox reactions) reactions that involve the transfer of
electrons from one species to another
Oxidation – is loses of electrons
Reduction – is gain of electrons
Early chemists saw “oxidation” reactions only as the combination of a material with
oxygen to produce an oxide.
A process called “reduction” is the opposite of oxidation, and originally meant the loss of
oxygen from a compound .
Oxidation and reduction always occur simultaneously
The substance gaining oxygen (or losing electrons) is oxidized, while the substance losing
oxygen (or gaining electrons) is reduced
Oxidation and Reduction (Redox)
Today, many of these reactions may not even involve oxygen
Redox currently says that electrons are transferred between reactants
Mg + S→ Mg2+ + S2- (MgS)
The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2
electrons, and is oxidized to Mg2+
The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons,
and is reduced to S2-
Oxidation number
Oxidation number (oxidation state) – is the total number of electrons that an atom
either gains or loses in order to form a chemical bond with another atom.
Rules for assigning oxidation numbers
1. The oxidation number of the atoms in any free, uncombined element is zero. This
includes polyatomic elements such as H2, O2, O3, P4, and S8.
H20 O20
oxidation number
2. The oxidation number of an element in a simple (monatomic) ion is equal to the
charge on the ion.
S-2 Cl- Br-
oxidation number
Rules for assigning oxidation numbers
Fluorine has an oxidation number of -1 in its compounds.
Hydrogen has an oxidation number of +1 in compounds unless it is combined with
metals, in which case it has an oxidation number of -1. Examples of these exceptions are
NaH and CaH2.
Oxygen usually has an oxidation number of -2 in its compounds. There are some
exceptions:
a. Oxygen has an oxidation number of -1 in hydrogen peroxide, H2O2, and in
peroxides, which contain the O22- ion; examples are CaO2 and Na2O2.
b. Oxygen has an oxidation number of - in superoxides, which contain the O2 ion;
examples are KO2 and RbO2.
c. When combined with fluorine in OF2, oxygen has an oxidation number of +2.
Rules for assigning oxidation numbers
The position of the element in the periodic table helps to assign its oxidation
number:
a. Group IA elements have oxidation numbers of +1 in all of their compounds.
b. Group IIA elements have oxidation numbers of +2 in all of their compounds.
c. Group IIIA elements have oxidation numbers of +3 in all of their compounds,
with a few rare exceptions.
d. Group VA elements have oxidation numbers of -3 in binary compounds with
metals, with H, or with NH4+.
Exceptions are compounds with a Group VA element combined with an element to
its right in the periodic table; in this case, their oxidation numbers can be found by using
rules 3 and 4.
e. Group VIA elements below oxygen have oxidation numbers of -2 in binary
compounds with metals, with H, or with NH4+. When these elements are combined with
oxygen or with a lighter halogen, their oxidation numbers can be found by using rules 3
and 4.
Rules for assigning oxidation numbers
The sum of the oxidation numbers of all atoms in a compound is zero.
H21+ O 2-
So 1 hydrogen atom loses 1 electron, 2 hydrogen atoms – 2e, 1 oxygen atom
gains 2 e
2٠ (+1) + 1٠(-2) =0
2 hydrogen atoms oxidation number of
oxygen
oxygen atom
oxidation number of
hydrogen
Rules for assigning oxidation numbers
In a polyatomic ion, the sum of the oxidation numbers of the constituent atoms is
equal to the charge on the ion.
SO42-
Overall charge of sulfate-ion is 2-
Oxidation number of S is 6+, of O – 2-
1٠(+6) + 4 ٠(-2) = -2
Find oxidation number of Cr in K2Cr2O7
+1 x -2
K2Cr2O7
2 (+1) + 2x + 7 (-2)= 0
2+2x-14= 0
2x=12
x=6
Exercise #1
1. Determine the oxidation numbers of nitrogen in the following species: N2O4,
NH3, HNO3, NO3-,N2.
2. Determine the oxidation numbers of elements in following compounds and ions:
MnO4-, CaH2, KClO3, Cr2O72-, MgCl2.
3. Determine the oxidation numbers of following elements in compounds:
a) sulfur in SO2, H2SO4, H2S, K2SO4.
b) halogen in KI, KClO3, HClO, HCl.
c) manganese in KMnO4, K2MnO4, MnO2, MnSO4.
d) chromium in K2Cr2O7, K3[Cr(OH)6], Cr2(SO4)3, K2CrO4.
e) metal in PbO2, K[Al(OH)4], K2[Zn(OH)4], CuSO4, FeS, Fe3(NO3)3.
f) nitrogen in HNO2, NO, NO2, NH4Cl.
Reducing and oxidizing agent
Oxidation – is loses of electrons
Reduction – is gain of electrons
reducing agent:
reducing agents decrease the ox. no. of another atom
an atom in the reducing agent increases in ox. no.
the reducing agent is the substance that gets oxidized – it loses electrons.
oxidizing agent:
oxidizing agents increase the ox. no. of another atom
an atom in the oxidizing agent decreases in ox. no.
the oxidizing agent is the substance which gets reduced – it gains electrons.
Examples of oxidizing agents
1.Most of non-metals (oxygen-O2; halogens: fluorine-F2, chlorine-Cl2, Br2-bromine),
except carbon –C, hydrogen –H and phosphorous- P;
2.Oxides and higher oxides of metals – copper oxide - CuO, manganese oxide –
MnO2, lead oxide - PbO2, chromium oxide – CrO3 etc;
3.Oxyacids - nitric acid - HNO3, hypochlorous acid - HClO, sulfuric acid - H2SO4,
etc;
4.Some salts of metals – FeCl3, CuCl2, HgCl2;
5.Peracids and their salts KMnO4, HClO4;
6.Chromic and Dichromic acids and their salts- H2CrO4 , K2Cr2O7 etc.
Examples of reducing agents
All metals (Na, K, Zn, etc.)
All non-metals such as FeCl2, FeSO4, CrCl2, etc
Hydrides of metals and non-metals such as NaH, NH3 , H2S, HI, etc;
Carbon –C, hydrogen –H and phosphorous- P.
There are some substances which are both oxidizing and reducing
agents
SO2,
nitrous acid - HNO2,
NO2,
hydrogen peroxide - H2O2,
sodium nitrite - NaNO2,
sodium sulfite - Na2SO3.
Exercise #2
1. Determine redox reactions among following:
a. Na2CrO4+H2SO4→Na2Cr2O7+Na2SO4+H2O
b. HClO4+SO2+H2O → HCl +H2SO4
c. Ag + H2SO4→Ag2SO4 +SO2 +H2O
d. NaOH + Cr2(SO4)3 →Na3[Cr(OH)6] + Na2SO4
2. Find reducing and oxidizing agents:
a. HCl+CrO3→Cl2+CrCl3+H2O
b. Fe+KNO3→KFeO2+ N2+ H2O
c. H2S +H2O→H2SO4 + H2O
d. Mn2(SO4)3 →MnSO4 + O2 +SO3
e. PbS+H2O2 →PbSO4 +H2O
f. KClO3 + KOH + MnO2→K2MnO4 + H2O + KCl
g. (NH4)2CrO4→Cr2O3 +N2+H2O+NH3
Types of the Oxidation-Reduction reactions
Intermolecular reaction: It is the reaction in which oxidizing and reducing agents are
localized in the different molecules.
Intramolecular reactions: It is the reaction in which oxidizing and reducing agents are
localized in one molecule
Dis-proportionation reactions (auto oxidation-reduction) in which the oxidation
number of atoms of the same element changes in both directions.
Balancing of redox reaction
Balance following equations:
1. C+HNO3=CO2+NO+ H2O
2. H2S+K2Cr2O7 +H2SO4=S +Cr2(SO4)3+K2SO4 + H2O
3. Ag + HNO3 = AgNO3 + NO + H2O
4. Ca +H2SO4 = CaSO4 + H2S + H2O
5. Be + HNO3 = Be(NO3)2 + NO + H2O
6. Fe2O3 + Al=Al2O3 +Fe
7. Mn2O3 + Si = SiO2+Mn
8. V2O5 + Ca+=CaO +V
9. NH3 + O2 =NO + H2O
10. P2O5+ C= P + CO
11. KClO3 + S= KCl + SO2
12. KNO2 + KClO3 = KCl + KNO3
Balance following equations:
13.SO2 + HNO3 + H2O = H2SO4 + NO
14.NaI+NaIO3 + H2SO4= I2+Na2SO4+ H2O
15.Cu+ HNO3 = Cu(NO3)2+ NO + H2O
16.Mg+ HNO3=Mg(NO3)2+ N2 +H2O
17.Mg+ HNO3=Mg(NO3)2+ NH4NO3 +H2O
18.Fe + HNO3= Fe(NO3)3+NO2+H2O
19.S + KClO3+H2O=Cl2+K2SO4+H2SO4
20.HNO3=NO2+O2+ H2O
21. KI +Cu(NO3)2 = CuI + I2 + KNO3
22.H3PO3 +KMnO4 +H2SO4=H3PO4 +MnSO4 +K2SO4+ 3H2O
23.(NH4)2Cr2O7 =Cr2O3+N2+H2O
24.Na2S2O3 +H2SO4 =Na2SO4+SO2+S+H2O
Balance following equations:
25. H2S + H2O2 = H2SO4 + H2O
26. Na2S2O3 +Br2 +NaOH = NaBr + Na2SO4 + H2O
27. Mn(NO3)2+NaBiO3 +HNO3 = HMnO4 + BiONO3 + NaNO3 +H2O
28. SnSO4 + KMnO4 + H2SO4 = Sn(SO4)2 + MnSO4 +K2SO4 +H2O
29. Na2SO3 + KIO3 + H2SO4 = I2 +Na2SO4 +K2SO4 +H2O
30. K2MnO4 + CO2 = KMnO4 +MnO2 +K2CO3
Oxidants and antiseptics
1. Hydrogen peroxide (H2O2) in the form of the 3% aqueous solution is used in
medicine as a topical antiseptic to treat minor cuts abrasions on the certain parts of
the body.
2. Potassium permanganate (KMnO4) in concentration rang from 0.01% to 0.2%, can
be used as a topical antiseptic.
3. Solutions of iodine are frequently used as antiseptics. In preparation for surgery,
skin may be painted with the brown antiseptic solutions.
Oxidizing agents are used as disinfectants
•Calcium hypochlorite [Ca(OCl)2], is used to disinfect clothing
and bedding
•Chlorine (Cl2) is used to kill pathogenic microorganisms in
the drinking water. Waster water is usually treated with
chlorine also before it is returned to the stream or lake.
•Ozone has also been used to disinfect the drinking water.
Ozone is more expensive than chlorine, but less amount of
first one is needed, ozone kills viruses on which chlorine
has little effect.
