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Friday, 26 June 2015


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Where there is oxidation, there is always reduction --Chemistry is essentially a study of redox systems.
Oxidation reaction and reduction reaction—

Oxidation reaction Reduction reaction
Addition of oxygen
C + O2 --- CO2
Removal of oxygen
Fe3O4 + 2C --- 3Fe + CO2
2 Removal of hydrogen .Addition of oxygen.
Removal of electron/s
Zn – 2e- -- Zn+2
Addition of electron/s.
Cl + e- -- Cl-
4. Addition of anion/s i.e Reaction with nonmetal/ electro–ve element Removal of anion/s or electro-ve element
5 Removal of electro+ve element Addition of ectro+ve element
. 6 Oxidation no. increases Oxidation no. decreases
Redox reaction—
Reactions in which both oxidation and reduction reactions take place simultaneously. e.g.
Reactions which involve change in oxidation number of the interacting species.

Oxidizing agent/oxidant i.e. accepter of electrons---
Reactant that oxidize the other reactant and itself that get reduce in given reaction is known as oxidizing agent or oxidant.
Reducing agent/reductant i.e. donar of the electrons-----
Reactant that reduce the other reactant and itself that get oxidized in given reaction is known as reducing agent or reductant.
Oxidation number---
Oxidation number denotes the oxidation state of an element in a compound ascertained according to a set of rules formulated on the basis that electron in a covalent bond belongs entirely to more electronegative element.

Difference between oxidation no. and valency---

Oxidation no. Valency
1. It is due to the partial shifting of electrons It is due to the complete loss of electrons
2. Related to the covalent bond formation. Related to the ionic bond Formation
3. It may be different in different compounds for same element It remain same in all compound for an element. (except transition elements.e.g. Fe e.t.c.)
4 It may be in fraction It always in whole number

Rules to determine the oxidation number---
1. In elements, in the free or the uncombined state, each atom bears an oxidation number of zero. Evidently each atom in H2,O2, Cl2, O3, P4, S8, Na, Mg, Al has the oxidation number zero.

2. For ions composed of only one atom, the oxidation number is equal to the chargeon the ion. Thus Na+ion has an oxidation number of +1, Mg2+ion, +2, Fe3+ion, +3, Clion, –1, O2–ion, –2; and so on. In their compounds all alkali metals have oxidation number of +1, and all alkaline earth metals have an oxidation number of +2. Aluminium is regarded to have an oxidation number of +3 in all its compounds.

3. The oxidation number of oxygen in most compounds is –2. However, we come acrosstwo kinds of exceptions here. One arises in the case of peroxides and superoxides, the compounds of oxygen in which oxygen atoms are directly linked to each other. While in peroxides (e.g., H2O2, Na2O2), each
oxygen atom is assigned an oxidation number of –1, in superoxides (e.g., KO2, RbO2) each oxygen atom is assigned an oxidation number of –(½). The second exception appears rarely, i.e. when oxygen is bonded to fluorine. In such compounds e.g., oxygen difluoride (OF2)and dioxygen difluoride (O2F2), the oxygen is assigned an oxidation number of +2 and +1, respectively. The number assigned to oxygen will depend upon the bonding state of oxygen but this number would now be a positive figure only.

4. The oxidation number of hydrogen is +1, except when it is bonded to metals in binary compounds (that is compounds containing two elements). For example, in LiH, NaH, and CaH2, its oxidation number is –1.

5. In all its compounds, fluorine has an oxidation number of –1. Other halogens (Cl, Br, and I) also have an oxidation number of –1, when they occur as halide ions in their compounds. Chlorine, bromine and iodine when combined with oxygen, for example in oxoacids and oxoanions, have positive oxidation numbers.

6. The algebraic sum of the oxidation number of all the atoms in a compound must be zero. In polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on the ion. Thus, the sum of oxidation number of three oxygen atoms and one carbon atom in the carbonate ion, (CO3)2–must equal

Trend of ox. No./ox.state in the periodic table—
In a period it increases (1to 7) from left to right.
In a group (as we go down in the group) it remains same.
The oxidation number state of a metal in a compound is sometimes presented according to the notation given by German chemist, Alfred Stock. It is popularly known as Stocknotation. According to this, the oxidation number is expressed by putting a Roman numeral representing the oxidation number in parenthesis after the symbol of the metal in the molecular formula. Thus aurous chloride and auric chloride are written as Au(I)Cl and Au(III)Cl3.
Types of redox reactions ----
  1. Combination reactions:Reaction in which two reactants combine together to form a single product is known as combination reaction.
A + B --C e.g.
C(s) + O2(g)----CO2(g)
3Mg(s) + N2(g)------Mg3N2(s)
2. Decomposition reaction;
Decomposition reactions are the opposite of combination reactions. A decomposition reaction leads to the breakdown of a compound into two or more components.
Examples 2H2O (l)-------2H2(g) + O2(g)
2NaH (s)------2Na (s) + H2(g)
2KClO3(s)----------2KCl (s) + 3O2(g)
Note: All decomposition reactions are not redox reactions. For example, decomposition of calcium carbonate is not a redox reaction.
CaCO3 (s)--------CaO(s) + CO2(g)
3. Displacement reactions
In a displacement reaction, an ion (or an atom) in a compound is replaced by an ion (or an atom) of another element. i.e.
X + YZ → XZ + Y
Displacement reactions have two categories:
1.metal displacement and 2. non-metal displacement.
(a) Metal displacement: A metal in a compound can be displaced by another metal in the uncombined state.
Applications: In metallurgical processes in which pure metals are obtained from their compounds in ores. examples ---
Fig. Redox reaction between zinc and aqueous solution of copper nitrate occurring in a beaker.
CuSO4(aq) + Zn (s) → Cu(s) + ZnSO4(aq)
V2O5(s) + 5Ca (s) -----2V (s) + 5CaO (s)
TiCl4 (l) + 2Mg (s) ------Ti (s) + 2 MgCl2(s)
Cr2O3(s) + 2 Al (s) -------Al2O3(s) + 2Cr(s)
In each case, the reducing metal is a better reducing agent than the one that is being reduced which evidently shows more capability to lose electrons as compared to the one that is reduced.
(b) Non-metal displacement: The non-metal displacement redox reactions include hydrogen displacement and a rarely occurring reaction involving oxygen displacement.
All alkali metals and some alkaline earth metals (Ca, Sr, and Ba) which are very good reductants, will displace hydrogen from cold water.
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
Less active metals such as magnesium and iron react with hot water and steam to produce dihydrogen gas:
Mg(s) + 2H2O(l) -----Mg(OH)2(s) + H2(g)
2Fe(s) + 3H2O(l)--------Fe2O3(s) + 3H2(g)
Many metals, such as Cadmium and tin displace the hydrogen from acids are:
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) -----------------------(1
Mg (s) + 2HCl (aq) → MgCl2(aq) + H2(g)
Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g) -----------------------(2
Reactions (1&2) are used to prepare dihydrogen gas in the lab.
Here, the reactivity of metals is reflected in the rate of hydrogen gas evolution, 
which is the slowest for the least active metal Fe, and the fastest for the most reactive metal, Mg. Very less active metals, which may occur in the native state 
such as silver (Ag), and gold (Au) do not react even with hydrochloric acid.

Reactivity of zinc (Zn), copper(Cu) and silver (Ag) is in the order Zn> Cu>Ag. Like metals, activity series also exists for the halogens. The power of these elements as oxidising agents decreases as we move down from fluorine to iodine in group 17 of the periodic table. This implies that fluorine is so reactive that it can replace chloride, bromide and iodide ions in solution. In fact, fluorine is so reactive that it attacks water and displaces the oxygen of water :
2H2O (l) + 2F2(g) → 4HF(aq) + O2(g)

Note: It is for this reason that the displacement reactions of chlorine, bromine and iodine using fluorine are not generally carried out in aqueous solution. On the other hand, chlorine can displace bromide and iodide ions in an aqueous solution as shown below:
Cl2(g) + 2KBr (aq) → 2 KCl (aq) + Br2(l)
Cl2 (g) + 2KI (aq) → 2 KCl (aq) + I2(s)
As Br2and I2 are coloured and dissolve in CCl4, can easily be identified from the colour of the solution. The above reactions can be written in ionic form as:
Hence the above reactions are the basis of identifying Brand Iin the lab and the test popularly known as ‘Layer Test’. Bromine likewise can displace iodide ion in solution:
Br2(l) + 2I (aq) → 2Br(aq) + I2(s)
Fluorine is the strongest oxidizing agent; there is no way to convert F– ions to F2by chemical means. The only way to achieve F2from F– is to oxidise electrolytically.

4. Disproportionation reactions
Disproportionation reactions are a special type of redox reactions. In a disproportionation reaction an element in one oxidation state is simultaneously oxidised and reduced. One of the reacting substances in a
disproportionation reaction always contains an element that can exist in at least three oxidation states.
2H2O2(aq) → 2H2O(l) + O2(g)
Here the oxygen of peroxide, which is present in –1 state, is converted to zero oxidation state in O2 and decreases to –2 oxidation state in H2O.
Phosphorous, sulphur and chlorine undergo disproportionation in the alkaline medium as shown below :
P4(s) + 3OH(aq)+ 3H2O(l) → PH3(g) + 3H2PO2–
S8(s) + 12 OH(aq) → 4S2– (aq) + 2S2O32–(aq) + 6H2O(l)
Cl2(g) + 2 OH(aq) → ClO(aq) + Cl(aq) + H2O (l)
The last reaction describes the formation of household bleaching agents. The hypochlorite ion (ClO–) formed in the reaction oxidises the colour-bearing stains of the substances to colourless compounds.
Same trend followed by bromine and iodine .Fluorine shows deviation from this behaviour when it reacts with alkali.
2 F2(g) + 2OH(aq) → 2 F(aq) + OF2(g) + H2O(l)
Because F is the most electronegative element, it cannot exhibit any positive oxidation state. This means that among halogens, fluorine does not show a disproportionation tendency.

Balancing of Redox Reactions:

Redox Reactions as the Basis for Titrations
In acid-base systems we come across with a titration method for finding out the strength of one solution against the other using a pH sensitive indicator.
Similarly, in redox systems, the titration method can be adopted to determine the strength of a reductant/oxidant using a redox sensitive indicator. The usage of indicators in redox titration is illustrated below:

(i) In one situation, the reagent itself is intensely coloured (known as self indicator), e.g., permanganate ion, MnO4 . Here MnO4acts as the self indicator. The visible end point in this case is
achieved after the last of the reductant (Fe2+or C2O42–)is oxidised and the first lasting tinge of pink colour appears at MnO4concentration as low as 10–6 mol dm–3(10–6mol L–1).

(ii) If there is no dramatic auto-colour change (as with MnO4– titration), there are indicators which are oxidised immediately after the last bit of the reactant is consumed, producing a dramatic colour change. The best example is afforded by Cr2O72–,which is not a self-indicator, but oxidises the indicator substance diphenylamine just after the equivalence point to produce an intense blue colour, thus signalling the end point.

(iii) There is yet another method which is interesting and quite common. Its use is restricted to those reagents which are able to oxidise I– ions, say, for example, Cu(II):
2Cu2+(aq) + 4I(aq) → Cu2I2(s) + I2(aq)
This method relies on the facts that iodine itself gives an intense blue colour with starch and has a very specific reaction with thiosulphate ions (S2O32–), which too is a redox reaction:
I2(aq) + 2 S2O32–(aq)→2I–(aq) + S4O62–(aq)
I2, though insoluble in water, remains in solution containing KI as KI3.
On addition of starch after the liberation of iodine from the reaction of Cu2+ ions on iodide ions, an intense blue colour appears. This colour disappears as soon as the iodine is consumed by the thiosulphate ions. Thus, the
end-point can easily be tracked.

We observed if zinc rod is dipped in copper sulphate solution. The redox
reaction takes place and during the reaction, zinc is oxidised to zinc ions and copper ions are reduced to metallic copper due to direct transfer of electrons from zinc to copper ion.
During this reaction heat is also evolved. Now we modify the experiment in such a manner that for the same redox reaction transfer of electrons takes place indirectly. We can do that by making a separation between Zn metal and copper sulphate solution. 

Preparation of salt bridge----
We connect solutions in two beakers by a salt bridge (a U-tube containing a solution of potassium chloride or ammonium nitrate usually solidified by boiling with agar agar and later cooling to a jelly like substance).
Significance of salt bridge------
This provides an electric contact between the two solutions without
allowing them to mix with each other.
We know that the flow of current is possible only if there is a potential difference between the copper and zinc rods known as electrodes here.
The potential associated with each electrode is known as electrode potential.
If the concentration of each species taking part in the electrode reaction is unity (if any gas appears in the electrode reaction, it is confined to 1 atmospheric pressure) and further the reaction is carried out at 298K, then the potential of each electrode is said to be the Standard Electrode Potential.
By convention, the standard electrode potential (E0 of hydrogen electrode is 0.00 volts). The electrode potential value for each electrode process is a measure of the relative tendency of the active species in the process to remain in the oxidised/reduced form. A negative E0means that the redox couple is a stronger reducing agent than the H+/H2couple. A positive E0means that the redox couple is a weaker reducing agent than the H+/H2couple.

  The standard electrode potentials are very important and we can get a lot of other useful information from them.

Chemistry for class 11.....


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