Silver [100% NCERT Extract]

→Alloys[Mixture]

→brass (mixture of copper and zinc)
→German silver (mixture of copper, zinc

and nickel)

→bronze (mixture of copper and tin)

→Important Mixture

→1 part per million (ppm) of fluoride ions in water

prevents tooth decay, while 1.5 ppm causes the tooth

to become mottled and high concentrations of fluoride

ions can be poisonous (for example, sodium fluoride is

used in rat poison)

→Galvanic Cells

→Galvanic cell is an

electrochemical cell that converts the chemical energy of a spontaneous

redox reaction into electrical energy. 

→In this device the Gibbs energy of

the spontaneous redox reaction is converted into electrical work which

may be used for running a motor or other electrical gadgets like heater,

fan, geyser, etc.

→Daniell cell discussed earlier is one such cell in which the following

redox reaction occurs.

→

→This reaction is a combination of two half reactions whose addition

gives the overall cell reaction:

→
→These reactions occur in two different portions of the Daniell cell.

The reduction half reaction occurs on the copper electrode while the

oxidation half reaction occurs on the zinc electrode. These two portions

of the cell are also called half-cells or redox couples. The copper

electrode may be called the reduction half cell and the zinc electrode,

the oxidation half-cell.

→We can construct innumerable number of galvanic cells on the pattern

of Daniell cell by taking combinations of different half-cells. Each halfcell

consists of a metallic electrode dipped into an electrolyte. The two

half-cells are connected by a metallic wire through a voltmeter and a

switch externally. The electrolytes of the two half-cells are connected

internally through a salt bridge. Sometimes, both

the electrodes dip in the same electrolyte solution and in such cases we

do not require a salt bridge.

→At each electrode-electrolyte interface there is a tendency of metal

ions from the solution to deposit on the metal electrode trying to make

it positively charged. At the same time, metal atoms of the electrode

have a tendency to go into the solution as ions and leave behind the

electrons at the electrode trying to make it negatively charged. At

equilibrium, there is a separation of charges and depending on the

tendencies of the two opposing reactions, the electrode may be positively

or negatively charged with respect to the solution. A potential difference

develops between the electrode and the electrolyte which is called

electrode potential. When the concentrations of all the species involved

in a half-cell is unity then the electrode potential is known as standard

electrode potential. According to IUPAC convention, standard

reduction potentials are now called standard electrode potentials. In a

galvanic cell, the half-cell in which oxidation takes place is called anode

and it has a negative potential with respect to the solution. The other

half-cell in which reduction takes place is called cathode and it has a

positive potential with respect to the solution. Thus, there exists a

potential difference between the two electrodes and as soon as the

switch is in the on position6 the electrons flow from negative electrode

to positive electrode. The direction of current flow is opposite to that of

electron flow

→The potential difference between the two electrodes of a galvanic

cell is called the cell potential and is measured in volts. The cell

potential is the difference between the electrode potentials (reduction

potentials) of the cathode and anode. It is called the cell electromotive

force (emf) of the cell when no current is drawn through the cell. It

is now an accepted convention that we keep the anode on the left and

the cathode on the right while representing the galvanic cell. A galvanic

cell is generally represented by putting a vertical line between metal

and electrolyte solution and putting a double vertical line between

the two electrolytes connected by a salt bridge. Under this convention

the emf of the cell is positive and is given by the potential of the halfcell

on the right hand side minus the potential of the half-cell on the

left hand side i.e.,

E(cell)= E(right) – E(left)

This is illustrated by the following example:

→

→It can be seen that the sum of (3.5) and (3.6) leads to overall reaction

(3.4) in the cell and that silver electrode acts as a cathode and copper

electrode acts as an anode. The cell can be represented as:

→

Electrolytic Cells and Electrolysis

→In an electrolytic cell external source of voltage is used to bring about

a chemical reaction.

→The electrochemical processes are of great importance

in the laboratory and the chemical industry

→One of the simplest electrolytic

cell consists of two copper strips dipping in an aqueous solution of

copper sulphate. If a DC voltage is applied to the two electrodes, then

Cu 2+ ions discharge at the cathode (negatively charged) and the following

reaction takes place.

Cu²⁺(aq) + 2e⁻ → Cu (s)

→Copper metal is deposited on the cathode. At the anode, copper is

converted into Cu2+ ions

→Thus copper is dissolved (oxidised) at anode and deposited

(reduced) at cathode. This is the basis for an industrial process in

which impure copper is converted into copper of high purity. The

impure copper is made an anode that dissolves on passing current

and pure copper is deposited at the cathode. Many metals like Na, Mg,

Al, etc. are produced on large scale by electrochemical reduction of

their respective cations where no suitable chemical reducing agents

are available for this purpose.

→Sodium and magnesium metals are produced by the electrolysis of

their fused chlorides and aluminium is produced (Class XII, Unit 6) by

electrolysis of aluminium oxide in presence of cryolite.

Quantitative Aspects of Electrolysis

→Michael Faraday was the first scientist who described the quantitative

aspects of electrolysis. Now Faraday’s laws also flow from what has

been discussed earlier.

Faraday’s Laws of Electrolysis

→After his extensive investigations on electrolysis of solutions and melts

of electrolytes, Faraday published his results during 1833-34 in the

form of the following well known Faraday’s two laws of electrolysis:

→(i) First Law: The amount of chemical reaction which occurs at any

electrode during electrolysis by a current is proportional to the

quantity of electricity passed through the electrolyte (solution or

melt).

→(ii) Second Law: The amounts of different substances liberated by the

same quantity of electricity passing through the electrolytic solution

are proportional to their chemical equivalent weights (Atomic Mass

of Metal ÷ Number of electrons required to reduce the cation).

→There were no constant current sources available during Faraday’s

times. The general practice was to put a coulometer (a standard electrolytic

cell) for determining the quantity of electricity passed from the amount

of metal (generally silver or copper) deposited or consumed. However,

coulometers are now obsolete and we now have constant current (I)

sources available and the quantity of electricity Q, passed is given by

Q = It

Q is in coloumbs when I is in ampere and t is in second.

→Q is in coloumbs when I is in ampere and t is in second.

The amount of electricity (or charge) required for oxidation or

reduction depends on the stoichiometry of the electrode reaction. For

example, in the reaction:

Ag ⁺(aq) + e⁻→Ag(s)

→One mole of the electron is required for the reduction of one mole

of silver ions.

→We know that charge on one electron is equal to 1.6021× 10⁻¹⁹C.

Therefore, the charge on one mole of electrons is equal to:

→

→This quantity of electricity is called Faraday and is represented by

the symbol F.

For approximate calculations we use 1F Y 96500 C mol⁻¹.

For the electrode reactions:

Mg^2+(l) + 2e– → Mg(s) 

Al^3+(l) + 3e– → Al(s) 

It is obvious that one mole of Mg2+ and Al3+ require 2 mol of

electrons (2F) and 3 mol of electrons (3F) respectively. The charge passed

through the electrolytic cell during electrolysis is equal to the product

of current in amperes and time in seconds. In commercial production

of metals, current as high as 50,000 amperes are used that amounts

to about 0.518 F per second.

Corrosion

→Corrosion slowly coats the surfaces of metallic objects with oxides or

other salts of the metal. The rusting of iron, tarnishing of silver,

development of green coating on copper and bronze are some of the

examples of corrosion. It causes enormous damage

to buildings, bridges, ships and to all objects made

of metals especially that of iron. We lose crores of

rupees every year on account of corrosion.

→In corrosion, a metal is oxidised by loss of electrons

to oxygen and formation of oxides. Corrosion of iron

(commonly known as rusting) occurs in presence of

water and air. The chemistry of corrosion is quite

complex but it may be considered

essentially as an electrochemical

phenomenon. At a particular spot

 of an object made of

iron, oxidation takes place and

that spot behaves as anode and

we can write the reaction

→Electrons released at anodic spot move through the metal and go

to another spot on the metal and reduce oxygen in the presence of H+

(which is believed to be available from H2CO3 formed due to dissolution

of carbon dioxide from air into water. Hydrogen ion in water may also

be available due to dissolution of other acidic oxides from the

atmosphere). This spot behaves as cathode with the reaction

→The ferrous ions are further oxidised by atmospheric oxygen to

ferric ions which come out as rust in the form of hydrated ferric oxide

(Fe2O3. x H2O) and with further production of hydrogen ions.

→Prevention of corrosion is of prime importance. It not only saves

money but also helps in preventing accidents such as a bridge collapse

or failure of a key component due to corrosion. One of the simplest

methods of preventing corrosion is to prevent the surface of the metallic

object to come in contact with atmosphere. This can be done by covering

the surface with paint or by some chemicals (e.g. bisphenol). Another

simple method is to cover the surface by other metals (Sn, Zn, etc.) that

are inert or react to save the object. An electrochemical method is to

provide a sacrificial electrode of another metal (like Mg, Zn, etc.) which

corrodes itself but saves the object.

→Question: Three electrolytic cells A,B,C containing solutions of ZnSO4, AgNO3 and CuSO4,respectively are connected in series. A steady current of 1.5 amperes waspassed through them until 1.45 g of silver deposited at the cathode of cell B.How long did the current flow? What mass of copper and zinc were deposited?

→Question:Predict the products of electrolysis in each of the following:
(i) An aqueous solution of AgNO3 with silver electrodes.
(ii) An aqueous solution of AgNO3 with platinum electrodes.
(iii) A dilute solution of H2SO4 with platinum electrodes.
(iv) An aqueous solution of CuCl2 with platinum electrodes.

Rate of a Chemical Reaction

→Some reactions such as ionic reactions occur very fast, for example,

precipitation of silver chloride occurs instantaneously by mixing of

aqueous solutions of silver nitrate and sodium chloride

→On the other

hand, some reactions are very slow, for example, rusting of iron in

the presence of air and moisture.

→Also there are reactions like inversion

of cane sugar and hydrolysis of starch, which proceed with a moderate

speed.

→speed of a reaction or the rate of a

reaction can be defined as the change in concentration of a reactant

or product in unit time. To be more specific, it can be expressed in

terms of:

(i) the rate of decrease in concentration of any one of the

reactants, or

(ii) the rate of increase in concentration of any one of the products.

Consider a hypothetical reaction, assuming that the volume of the

system remains constant.

R → P

One mole of the reactant R produces one mole of the product P. If

[R]1 and [P]1 are the concentrations of R and P respectively at time t1

and [R]2 and [P]2 are their concentrations at time t2 then,

→

Adsorption

→The accumulation of molecular species

at the surface rather than in the bulk of a solid or liquid is termed

adsorption.

Applications of Adsorption

→The phenomenon of adsorption finds a number of applications.

Important ones are listed here:

→(i) Production of high vacuum: The remaining traces of air can be

adsorbed by charcoal from a vessel evacuated by a vacuum pump

to give a very high vacuum

→(ii) Gas masks: Gas mask (a device which consists of activated charcoal

or mixture of adsorbents) is usually used for breathing in coal

mines to adsorb poisonous gases.

→(iii) Control of humidity: Silica and aluminium gels are used as

adsorbents for removing moisture and controlling humidity.

→(iv) Removal of colouring matter from solutions: Animal charcoal

removes colours of solutions by adsorbing coloured impurities.

→(v) Heterogeneous catalysis: Adsorption of reactants on the solid

surface of the catalysts increases the rate of reaction. There are

many gaseous reactions of industrial importance involving solid

catalysts. Manufacture of ammonia using iron as a catalyst,

manufacture of H2SO4 by contact process and use of finely divided

nickel in the hydrogenation of oils are excellent examples of

heterogeneous catalysis.

→(vi) Separation of inert gases: Due to the difference in degree of adsorption

of gases by charcoal, a mixture of noble gases can be separated by

adsorption on coconut charcoal at different temperatures.

→(vii) In curing diseases: A number of drugs are used to kill germs by

getting adsorbed on them.

→(viii) Froth floatation process: A low grade sulphide ore is concentrated

by separating it from silica and other earthy matter by this method

using pine oil and frothing agent (see Unit 6).

→(ix) Adsorption indicators: Surfaces of certain precipitates such as

silver halides have the property of adsorbing some dyes like eosin,

fluorescein, etc. and thereby producing a characteristic colour at

the end point.

→(x) Chromatographic analysis: Chromatographic analysis based on

the phenomenon of adsorption finds a number of applications in

analytical and industrial fields.

Intext.

Colloids

→A colloid is a heterogeneous system in which one substance is

dispersed (dispersed phase) as very fine particles in another substance

called dispersion medium.

The essential difference between a solution and a colloid is that of

particle size. While in a solution, the constituent particles are ions or

small molecules, in a colloid, the dispersed phase may consist of

particles of a single macromolecule (such as protein or synthetic

polymer) or an aggregate of many atoms, ions or molecules. Colloidal

particles are larger than simple molecules but small enough to remain

suspended. Their range of diameters is between 1 and 1000 nm

(10^–9 to 10^–6 m).

Preparation of Colloids

→A few important methods for the preparation of colloids are as follows:

→(a) Chemical methods

Colloidal dispersions can be prepared by chemical reactions leading

to formation of molecules by double decomposition, oxidation,

reduction or hydrolysis. These molecules then aggregate leading to

formation of sols.

→

→(b) Electrical disintegration or Bredig’s Arc method

This process involves dispersion as well as

condensation. Colloidal sols of metals such as gold,

silver, platinum, etc., can be prepared by this

method. In this method, electric arc is struck

between electrodes of the metal immersed in the

dispersion medium .The intense heat

produced vapourises the metal, which then

condenses to form particles of colloidal size.

→(c) Peptization

Peptization may be defined as the process of converting a precipitate

into colloidal sol by shaking it with dispersion medium in the presence

of a small amount of electrolyte. The electrolyte used for this purpose

is called peptizing agent. This method is applied, generally, to convert

a freshly prepared precipitate into a colloidal sol.

During peptization, the precipitate adsorbs one of the ions of the

electrolyte on its surface. This causes the development of positive or

negative charge on precipitates, which ultimately break up into smaller

particles of the size of a colloid. You will learn about the phenomenon

of development of charge on solid particles and their dispersion in next

Section  under the heading “Charge on collodial particles”.

→Medicines: Most of the medicines are colloidal in nature.
For example, argyrol is a silver sol used as an eye lotion.
Colloidal antimony is used in curing kalaazar. Colloidal
gold is used for intramuscular injection. Milk of
magnesia, an emulsion, is used for stomach disorders.

→Photographic plates and films: Photographic plates or films areprepared by coating an emulsion of the light sensitive silver bromidein gelatin over glass plates or celluloid films.

→The ‘Seven metals

of antiquity’, as they are sometimes called, are gold,

copper, silver, lead, tin, iron and mercury.

→Harappans also used gold and silver, as well as their joint alloy

electrum.

→Early gold and silver

ornaments have been found from Indus Valley sites such as

Mohenjodaro (3000 BCE).

→Mauryan era, which has much data on

prevailing chemical practices in a long section on mines and minerals

including metal ores of gold, silver, copper, lead, tin 

→
→In the metallurgy of silver and gold, the respective metal is

leached with a dilute solution of NaCN or KCN in the presence

of air, which supplies O2. The metal is obtained later by

replacement reaction

→extraction of gold and silver involves leaching the

metal with CN–. This is also an oxidation reaction (Ag ® Ag+ or Au ® Au+).

The metal is later recovered by displacement method.

→Impurities from the blister copper deposit as anode mud which

contains antimony, selenium, tellurium, silver, gold and platinum;

recovery of these elements may meet the cost of refining. Zinc may

also be refined this way.

→Zinc is used for galvanising iron. It is also used in large quantities

in batteries. It is constituent of many alloys, e.g., brass, (Cu 60%, Zn

40%) and german silver (Cu 25-30%, Zn 25-30%, Ni 40–50%). Zinc

dust is used as a reducing agent in the manufacture of dye-stuffs,

paints, etc

→The acids which contain P–H bond have strong reducing properties.

Thus, hypophosphorous acid is a good reducing agent as it contains

two P–H bonds and reduces, for example, AgNO3 to metallic silver.

→Iron, copper, silver and gold are among the transition elements that

have played important roles in the development of human civilisation.

The inner transition elements such as Th, Pa and U are proving

excellent sources of nuclear energy in modern times.

→Alfred Werner (1866-1919), a Swiss chemist was the first to formulate

his ideas about the structures of coordination compounds. He prepared

and characterised a large number of coordination compounds and

studied their physical and chemical behaviour by simple experimental

techniques. Werner proposed the concept of a primary valence and

a secondary valence for a metal ion. Binary compounds such as

CrCl3, CoCl2 or PdCl2 have primary valence of 3, 2 and 2 respectively.

In a series of compounds of cobalt(III) chloride with ammonia, it was

found that some of the chloride ions could be precipitated as AgCl on

adding excess silver nitrate solution in cold but some remained in

solution.

→Some important extraction processes of metals, like those of silver and

gold, make use of complex formation. Gold, for example, combines with

cyanide in the presence of oxygen and water to form the coordination

entity [Au(CN)2]– in aqueous solution. Gold can be separated in metallic

form from this solution by the addition of zinc

→Articles can be electroplated with silver and gold much more

smoothly and evenly from solutions of the complexes, [Ag(CN)2]–

and [Au(CN)2]– than from a solution of simple metal ions.

→Let us extend electron transfer reaction now

to copper metal and silver nitrate solution in

water and arrange a set-up as shown in

Fig. 8.2. The solution develops blue colour due

to the formation of Cu2+ ions on account of the

reaction:

→Test for Halogens

The sodium fusion extract is acidified with nitric

acid and then treated with silver nitrate. A white

precipitate, soluble in ammonium hydroxide

shows the presence of chlorine, a yellowish

precipitate, sparingly soluble in ammonium

hydroxide shows the presence of bromine and