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CHEMICAL SCIENCES
SUBJECT CODE – 102
INORGANIC CHEMISTRY
UNIT - III
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[1]
Chemistry
Ch. No.
Contents
Pg. Nos.
1
Main Group Element
3 – 46
2
Co-ordination Compound
1 - 33
[2]
Chemistry
1.
The correct order of stability of difluorides is :
[June 2011]
(a) GeF2>SiF2>CF2
(b) CF2>SiF2 > GF2
(c) SiF2> GeF2>CF2
(d) CF2> GF2>SiF2
Ans. (a) Hint: Inter pair effect (As we move down the group there is a reluctance of
ns electrons to participate in bond formation) Therefore +2 state stability will be in the
order:- Ge>Si>C. Therefore the correct order of stability of difluorides is therefore
GeF2>SiF2>CF2
2.
Lewis acidity of BCl3, BPh3 and BMe3 with respect to pyridine follows the order [Dec. 2011]
(a) BCl3>BPh3>BMe3
(b) BMe3>BPh3>BCl3
(c) BPh3>BMe3>Bcl3
(d) BCl3>BMe3>BPh3
Ans. (b) Lewis acidity of BCl3, BPh3 and BMe3 with respect to pyridine follows the
order. –BMe3> BPh3>BCl3 Among BCl3, BMe3 & PPh3 & BCl3 is weak Lewis acidic
due 2 pB  3 pB  3 pCl  back bonding. Among BMe3, and BPh3, BMe3 is strong
Lewis acid with respect to pyridine due to steric factor in BPh3.
3.
The material that exhibits the highest electrical conductivity among the following sulfur –
nitrogen compounds is :
[Dec 2011]
(a) S4N4
(b) S7NH
(c) S2N2
(d) (SN)x
Ans. (d) Sulfur forms a variety of covalent binary nitrides, but the two most
interesting ones are tetrasulfur & tetranitride S4N4, and disulfur dinitirde, S2N2,
because they are precursors to an unusual polymer called polythiazyl, (SN)x. This
polymeric sulfur nitride is unusual because, even though it is composed solely of two
nonmetals, if exhibits some properties normally associated only with metals like
electrical and thermal conductivity & It becomes a superconductor at 0.26 kelvin.
[June 2012]
(a) Cl>S>P>Si
(b) Cl>P>S>Si
(c) P>S>Si>Cl
(d) Si>P>S>Cl
Ans. (d) Hint: The size of the s, p and d orbitals generally decreasses with decrease
in size of an element And as we know in periods, the atomic radius decreases across
the period from left to right. Therefore order of the sizes of d orbitals is Si>P>S>Cl
4.
The size of the d orbitals in Si, P, S and Cl follows the order.
5.
The electronegativity differences is the highest for the pair
6.
Which of the following pairs has the highest difference in their first ionization energy ?
[June 2013]
[Dec. 2012]
(a) Li, Cl
(b) K,F
(c) Na, Cl
(d) Li, F
Ans. (b) Among metal, order of electronegativity is Li>Na>K & among halide order
of electronegativity is F>Cl. Therefore an electronegativity difference is the highest
for the K and F.
(a) Xe, Cs
(c) Ar, K
(b) Kr, Rb
(d) Ne, Na
Ans. (d) Order of size of alkali metal is Na<K<Rb<Cs & first ionizatioin energy is
Na>K>Rb>Cs while that of Nobel gases, order of size is Ne<Ar<Kr<Xe & first
ionization energy is Ne> Ar>Kr>Xe. Difference in first ionization energy betweend
[3]
Chemistry
alkali metal & nobel gases of same period is decreases top to bottom. Thus the Na, Ne
pairs has the highest difference in their first ionization energy.
[Dec. 2013]
(a) BF3OEt2
(b) SF4

(c) PFe6 
(d) Sb2S3
Ans. (c) A hypervalent molecule (expanded octet) is a molecule that contains one or
more main group elements with more than eight electrons in their valence shells e.g.
PFe6 
7.
Among the following, an example of a hypervalent species is
8.
The strength of p  d bonding in E-O (E=Si, P, S and C) follows the order
[June 2012]
(a) Si-O>P-O>S-O>Cl-O
(b) P-O>Si-O>S-O>Cl-O
(c) S-O>Cl-O>P-O>Si-O
(d) Cl-O>S-O>P-O>Si-O
Ans. (d) As the size increases, strength of p  d bonding decreases because inter
nuclear distance between two atoms increases & strength of  bond decreases due to
poor overlap as  bond is formed by side by side overlap. As increasing order dof size
is Si>P>S>Cl & therefore order of strength of p  d bonding is Cl-O>S-O>P-O>SiO
[June 2014]
(a) As>Al>Ca>S
(b) S>As>Al>Ca
(c) Al>Ca>S>As
(d) S>Ca>As>Al
Ans. (b) The correct order of decreasing electronegativity is : S(2.58)>As(2.18)>
Al(1.61)>Ca(1). In periodic table electronegativity increases from left to right in
period & decreases top to bottom in group.
9.
The correct order of decreasing electronegativity of the following atoms is :
10.
The correct order of the size of S, S2-,S2+ and S4+ species is :
[June 2014]
(a) S  S 2 S 4  S 2
(b) S 2 S 4  S 2  S
(c) S 2 S  S 2  S 4
(d) S 4  S 2  S  S 2
Ans. (c) The correct order of the size of . S 2  S  S 2  S 4 General order of size:-
Anoinic species>Neutral atom>Cation species.
11.
The series with correct order of decreasing ionic size is :
[GATE 2006]

2
2

2


2
(a) K  Ca  S  Cl
(b) S  Cl  K  Ca


2
2
(c) K  Cl  Ca  S
(d) Cl   K   S 2  Ca 2
Ans.
(b) The series with correct order of decreasing ionic size is
S  Cl  K   Ca 2 General order of size: Anoinic species > neutal atom>Cation
species
2
12.
1
[GATE 2007]
The Lewis acid character of BF3, BCl3 follows the order :
(a) BF3<BBr3BCl3
(c) BF3<BCl3<BBr3
(b) BCl3<BBr3<BF3
(d) BBr3<BCl3<BF3
Ans.
(c) The Lewis acidity BF3 is less than BCl3 because of stronger
2 pB   2 pF  back bonding in case of BF3 which is not effective in case of BCl3
due to weaker 2 pB   3 pCl  back bonding and this  back bonding became worse
[4]
Chemistry
in case of BBr3; due to poorer 2 pB   4 pBr  back bonding. Therefore Lewis
acidity order is BF3  BCl3  BBr3  Bl3
13.
[GATE 2009]
Among the following, the isoelectronic and isostructural pair is :
(a) CO2 and SO2
(c) NO2 and TeO2
(b) SO3 and SeO3
(d) SiO44 and PO43
Ans. (d) Among all SiO44 & PO43 is the isoelectronc and isostructural pair.
Pair Molelcular/ion Total electrons
(a) CO2
22
(b)
(c)
(d)
14.
Structure according to VSEPR
linear
SO2
32
Bent or V shaped
SO3
40
trigonal planar
SeO3
58
trigonal planar
NO2
TeO2
24
bent or V shaped
68
bent or V shaped
SiO4
50
tetrahedral
PO43
50
tetrahedral
Among the following, the group of molecules that undergoes rapid hydrolysis is
[GATE 2011]
(a) SF6, Al2Cl6, SiMe4
(b) BCl3, SF6. SiCl4
(c) BCl3, SiCl4, PCl5
(d) SF6, Al2Cl6, SiCl4
Ans. (c) Among all SF6, Al2Cl6 doesn’t undergo hydrolysis, In SF6S is in high
oxidation state (VI) & as F is bad leaving group) nucleophilic atack of water is not
possible while Al2Cl6 is dimeric structure therefore does’t undergo hydrolysis. In case
of BCl3 SiCl4 & PCl5 water can attack on central atom & undergoes rapid hydrolysis.
15.
The reaction of solid XeF 2 with AsF6 in 1:1 ratio affords
(a) Al 2Cl6 , SiMe4
(b) BCl3 , SF6 , SiCl 4
(c) XeF  ASF6 

[GATE 2011]
(d) Xe2 F3  AsF6 



Ans. (c) The reaction of solid XeF2 with AsF5 in 1 : 1 ratio affords XeF  and AsF6 
Hint : BF3, AsF5 and SbF5 are the F- acceptor & it leads to the formation nof a lewis
base due to electron withdrawing nitro group while in case of (B) & (D) it leads to
steric crowding in adduct.

[5]
Chemistry

CHAPTER-1
MAIN GROUP ELEMENTS
HYDROGEN AND THE HYDRIDES:
ELECTRONIC STRUCTURE :
Hydrogen has the simple atomic structure of all the elements, and consists of a nucleus containing one
proton with a charge +1 and one orbital electron. The electronic structure may be written as 1s1.
1. By forming an electron pair (covalent) bond with another atom Non-metals typically form this
type of bond with hydrogen, for example H 2 , H 2O, HCl gasor CH 4 and many metals do so
too.
2. By losing an electron to form H+.
A proton is extremely small (radius approximately 1.5 105 Å for most atoms). Because H+
is so small, it has a very high polarizing power, and therefore distorts the electron cloud on
other atoms. Thus protons are always associated with other atoms or molecules. For example,
in water or aqueous solutions of HCl and H2SO4 , protons exist as H 3O , H 9O4 H H 2O n
ions. Free protons do not exist under ‘normal conditions’, though they are found in low
pressure gaseous beams, for example in a mass spectrometer.
3. By gaining an electron to form HCrystalline solids such as LiH contain the H  ion and are formed by highly electropositive
metals (all of Group 1, and some of Group2) However, H  ions are uncommon.
Since hydrogen has an electro negativity of 2.1, it may use any of the three methods, but the
most common way is forming covalent bonds.



PREPARATION OF HYDROGEN
Hydrogen is manufactured on a large scale by a variety of methods.
1. Hydrogen is made cheaply, and in large amounts, by passing steam over red hot coke. The
product is water gas, which is a mixture of CO and H2. This is an important industrial fuel
since it is easy to make and it burns, evolving a lot of heat.
C
C  H 2O 1000

CO  H 2



o
water gas
CO  H 2  O2  CO2  H 2O  heat
It is difficult to obtain pure H 2 from water gas, since CO is difficult to remove. The CO may
be liquefied at a low temperature under pressure, thus separating it from H 2 .Alternatively
the gas mixture can be mixed with steam, cooled to 4000oC and passed over iron oxide in a
shift converter, giving H 2 and CO2 . The CO2 so formed is easily removed either by
dissolving in water under pressure, or reacting with
thus giving H 2 gas
K 2CO3 solution, giving KHCO3 , and
H 2O
CO  H 2 
 2 H 2  CO2


 450o C
water gas
Fe 2O3
2. Hydrogen is also made in large amounts by the steam reformer process. The hydrogen
produced in this way is used in the Haber process to make NH 3 and for hardening oils . Light
hydrocarbons such as methane are mixed with steam and passed over a nickel catalyst at 800900oC. These hydrocarbons are present in natural gas, and are also produced at oil refineries
when ‘cracking’ hydrocarbons.
CH 4  H 2O  CO  3H 2
CH 4  2 H 2O  CO2  4 H 2
[6]
Chemistry
The gas emerging from the reformer contains CO, CO2 .H 2 and excess steam. The gas
mixture is mixed with more steam, cooled to 400oC and passed into a shift converter. This
contains an iron/ copper catalyst and CO is converted into CO2.
CO  H 2O  CO2  H 2
3. Pure hydrogen (99.9%pure) is made by electrolysis of water or solutions of NaOH or KOH.
This is the most expensive method. Water does not conduct electricity very well. So it is usual
to electrolyse aqueous solutions of NaOH or KOH in a cell with nickel anodes and iron
cathodes. The gases produced in the anode and cathode compartments must be kept separate.
2OH   H 2O  ½O2  2e
2H 2O  2e  2OH   H 2
Anode
Cathode
Overall
H 2O  2OH  H 2
4. The laboratory preparation is the reaction of dilute acids with metals.
Zn  H 2 SO4  ZnSO4  H 2
2 Al  2 NaOH  6 H 2O  2 NaAl OH 4   3H 2
PROPERTIES OF MOLECULAR HYDROGEN
Hydrogen is the lightest gas known, and because of its low density, it is used instead of helium to fill
balloons for me meteorology. It is colourless, odourless and almost insoluble in water, Hydrogen
forms diatomic molecules H 2 , and the twos atoms are joined by a very strong covalent bond (bond
1
energy 435.9 kJ mol ) .
Hydrogen is not very reactive under normal conditions. The lack of reactivity is due to kinetics rather
than thermodynamics, and relates to the strength of the H  H bond to produce atoms of hydrogen.
This requires 435.9kJ mol 1 .hence there is a high activation energy to such reactions . Consequently
many reactions are slow or require high temperatures, or catalysts (often transition metals). Many
important reactions of hydrogen involve heterogeneous catalysis, where the catalyst first reacts with
H 2 and either breaks or weakens the H  H bond, and thus lowers the activation energy, Examples
include :
1. The Haber process for the manufacture of NH 3 from N 2 and H 2 using a catalyst of
activated Fe at 380  4500 C and 200 atmospheres pressure.
2. The hydrogenation of a variety of unsaturated organic compounds, (including the hardening
of oils), using finely divided Ni, Pd or Pt as catalysts.
3. The production of methanol by reducing CO with H 2 over a Cu / Zn catalyst at 300oC.
Thus hydrogen will react directly with most elements under the appropriate conditions.
Hydrogen burns in air or dioxygen, forming water, and liberates a large amount of energy.
This is used in the oxy-hydrogen flame for welding and cutting metals . Temperatures of
almost 3000oC can be attained. Care should be taken with these gases since mixtures of H 2
and O2 close to a 2 : 1 ratio are often explosive.
2H 2  O2  2H 2O H  485kJ mol 1
Hydrogen reacts with the halogens. The reaction with fluorine is violent, even at low
temperatures. The reaction with chlorine is slow in the dark, but the reaction is catalysed by
light (photocatalysis), and becomes faster in daylight, and explosive in sunlight. Direct
combination of the elements is used to produce HCl.
H 2  F2  2 HF
H 2  Cl2  2 HCl
A number of metals react with H2, forming hydrides. The reaction are not violent, and usually
require a high temperature. These are described in a later section
[7]
Chemistry
Large quantities of H2 are used in the industrial production of ammonia by the Haber process.
The reaction is reversible, and the formation of NH3 is favoured by high pressure, the
presence of a catalyst (Fe), and a low temperature . In practice a high temperature of 380450oand a pressure of 200 atmosphere are used to get a reasonable conversion in a reasonable
time.
N 2  3H 2
g 298k  33.4Kj mol 1
Large amounts of H2 are used for hydrogenation reactions, in which hydrogen is added to a
double bond in an organic compound. An important example is the hardening of fats and ioils.
Unsaturated fatty acids acids which have higher melting points. By removing double bonds in
the carbon chain in this way, edible oils which are liquid at room temperature may be
converted into fats which are solid at room temperature. The reason for doing this is that solid
fats are more useful than oils, for example in the manufacture of margarine.
CH 3.CH 2 n .CH  CH .COOH  H 2  CH 3.CH 2 n .CH 2 .COOH
Hydrogen is also used to reduce nitrobenzene to aniline (dyestuffs industry), and in the
catalytic reduction of benzene (the first step in the production of nylon-66). It also reacts with
CO to form methyl alcohol.
CO  2H 2 catalyst
 CH 3OH
The hydrogen molecule is veryd stable, and has little tendency to dissociate at normal
temperatures, since the dissociation reaction is highly endothermic.
H 2  2H
H  435.9kJ mol 1
However, at high temperatures, inm an electric arc, or under ultraviolet light, H2 does
dissociate. The atomic hydrogen produced exists for less than half a second, after which it
recombines to give molecular hydrogen Atomic hydrogen is a strong reducing agent, and is
commonly prepared in solution by moons of a zinc-copper couple or a mercury-aluminium
couple.
ORTHO AND PARA HYDROGEN
The hydrogen molecule H2 exists in two different forms known as ortho. And para hydrogen. The
nucleus of an atom has nuclear spin, in a similar way to electrons having a spin. In the H 2 molecule,
the two nuclei may be spinning in either the same direction, or in opposite directions. This gives rise
to spin isomerism, that is two different forms of H2 may exist. These are called ortho and para
hydrogen. Spin isomerism is also found in other symmetrical molecules whose nuclei have spin
momenta, e.g. D2,N2, F2, Cl2, There are considerable differences between the physical properties(e.g.
boiling points, specific heats and thermal conductivities) of the ortho and para forms, because of
differences in their internal energy. There are also differences in the band spectra of trhe ortho and
para forms of H2.
The para form has the lower energy, and at absolute zero the gas contain 100%of the para form. As
the temperature is raised, some of the para form changes into the ortho form. At high temperatures the
gas contains about 75% ortho hydrogen.
(a)
(b)
Ortho and para hydrogen,(a) ortho, parallel spins;(b) para, opposite
Para hydrogen is usually prepared by passing a mixture of the two forms of hydrogen through a tube
packed with charcoal cooled to liquid air room temperature in a glass vessel, because the ortho-para
conversion is slow in the absence of catalysts. Suitable catalysts include activated charcoal, atomic
[8]
Chemistry
hydrogen, metals such as Fe, Ni, Pt and W and paramagnetic substances or ions (which contain
unpaired electrons) such as O2, NO, NO2, Co2+and Cr2O3.
Electronic Structure
Group 1 elements all have one valency electron in their outer orbital-an s electron which
occupies a spherical orbital. Ignoring the filled inner shells the electronic structures may be written:
2s1, 3s1, 4s1, 5s1, 6s1 and 7s1. The single valence electron is a long distance from the nucleus, is only
weakly held and is readily removed. In contrast the remaining electrons are closer to the nucleus,
more tightly held, and are removed only with great difficulty. Because of similarities in the electronic
structures of these elements, many similarities in chemical behaviour would be expected.
Electronic Structure
Element
Atomic
Symbol
Electronic Structure
Number
Lithium
3
Li
1s2 2s1
or [He] 2s1
2
2
6
1
Sodium
11
Na
1s 2s 2p 3s
or [Ne] 3s1
Potassium
19
K
1s2 2s2 2p6 3s2 3p6 4s1
or [Ar] 4s1
2
2
6
2
6
10
2
6
1
Rubidium
37
Rb
1s 2s 2p 3s 3p 3d 4s 4p 5s
or [Kr] 5s1
2
2
6
2
6
10
2
6
10
2
6
1
Caesium
35
Cs
1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p 6s
or [Xe] 6s1
Francium
87
Fr
or [Rn] 7s1
Size of Atoms and Ions
Group 1 atoms are the largest in their horizontal periods in the periodic table. When the outer
electron is removed to give a positive ion, the size decreases considerably. There are two reasons for
this.
1. The outermost shell of electrons has been completely removed.
2. Having removed an electron, the positive charge on the nucleus is now greater than the charge
on the remaining electrons, so that each of the remaining electrons is attracted more strongly
towards the nucleus. This reduces the size further.
Positive ions are always smaller than the parent atom. Even so, the ions are very large and
they increase in size from Li+ to Fr+ as extra shells of electrons are added.
The Li+ is much smaller than the other ions. For this reason, Li only mixes with Na above
o
380 C and it is immiscible with the metals K, Rb and Cs, even when molten; nor will Li form
substitutional alloys with them. In contrast the other metals Na, K, Rb and Cs are miscible with each
other in all proportions.
Density
The atoms are large, so Group 1 elements have remarkably low densities. Lithium metal is
only about half as dense as water, whilst sodium and potassium are slightly less dense than water. It is
unusual for metals to have low densities, and in contrast most of the transition metals have densities
greater than 5g cm-3 for example iron 7.9 g cm-3, mercury 13.6 g cm-3, and osmium and iridium (the
two most dense elements) 22.57 and 22.61 g cm-3 respectively.
Table: Size and density

Ionic radius M+
Density
Metallic radius( A )

(gcm-3)
six-coordinate( A )
Li
1.52
0.76
0.54
Na
1.86
1.02
0.97
K
2.27
1.38
0.86
Rb
2.48
1.52
1.53
Cs
2.65
1.67
1.90
Ionization Energy
The first ionization energies for the atoms in this group are appreciably lower than those for
any other group in the periodic table.
[9]
Chemistry
The second ionization energy-that is the energy to remove a second electron from the atoms-is
extremely high. The second ionization energy is always larger than the first, often by a factor of two,
because it involves removing an electron from a smaller positive ion, rather than from a larger neutral
atom. The difference between first and second ionization energies is much larger in this case since in
addition it corresponds to removing an electron from a closed shell.A second electron is never
removed under normal conditions, as the energy required is greater than that needed to ionize the
noble gases. The elements commonly form M* ions.
Ionization energies
Li
Na
K
Rb
Cs
First ionization energy
(kJ mol-1)
520.1
495.7
418.6
402.9
375.6
Second ionization energy
(kJ mol-1)
7296
4563
3069
2650
2420
Electronegativity and Bond Type
The electronegativity values for the elements in this group are very small-in fact the smallest
values of any element. Thus when these elements react with other elements to form compounds, a
large electronegativity difference between the two atoms is probable, and ionic bonds are formed.
Na electronegativity
0.9
Cl elect ronegativity
3.0
Electronegativity difference
2.1
Electronegativity values
Pauling's electronegativity
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Cs
0.7
An electronegativity difference of approximately 1.7-1.8corresponds to 50% ionic character.
TRhe value 2.1 exceeds this, so the bonding in NaCl is predominantly ionic. Similar arguments apply
to other compounds: for example, the electronegativity difference in Lif is 3.0 and in KBr is 2.0 , and
both compounds are ionic.
The chemistry of the alkali metals is largely that of their ions.
Born-Haber Cycle:Enthalpy changes in the Formation of Ionic Compounds
When elements react to form compounds, G (the free energy of formation) is negative. For
a reaction to proceed spontaneously, the free energy of the products must be lower than that of the
reactants.
Usually the energy changes are measured as enthalpy values H, and G is related to H
by the equation.
G = H - TS
In many cases enthalpy values are used instead of free energy values, and the two are almost
the same if the term TS is small, At room temperature T is almost 300K, so G and H are only
similar when the change in entropy S is very small. Entropy changes are large when there is a
change in physical state, e.g. solid to liquid, or liquid to gas, but otherwise entropy changes are
usually small.
There is no direct way of obtaining electron affinity values, and these have been calculated
from this type of energy cycle.
Hess’s law states that the energy change occurring during a reaction depends only on the
energy of the initial reactants and the energy of the final products and not on the reaction mechanism,
or the route taken. Thus, by Hess’s law, the energy change for the reaction of solid sodium and
[10]
Chemistry
chlorine gas to from a sodium chloride crystal by the direct route (measured as the enthalpy of
formation) must be the same as the sum of all the energy changes going round the cycle by the long
route, i.e. by producing first gaseous almost of the all the energy changes going round the cycle by the
long route, i.e. by producing first gaseous atoms of the elements, then gaseous ions, and finally
packing these to give the crystalline solid. This may be expressed elements, then gaseous ions, and
finally packing these to give the
crystalline solid. This may be expressed
as :
Hf = +Hs + 1+ ½ Hd + E + U
All the halides MCI have
negative enthalpies of formation, which
indicates that thermodynamically (that
is in terms of energy) it is feasible to
form the compounds MCI from the
elements.
1. The most negative enthalpies of
formation occur with t1e
fluorides. For any given metal,
the values decrease in the
sequence fluoride > chloride >
bromide > iodide. Thus the
fluorides are the most stable,
and the iodides the least stable.
2. The enthalpies of formation for the chlorides, bromides and iodides become more negative on
descending the group. This trend is observed with most salts, but the opposite trend is found
in the fluorides.
Ionic compounds may also be formed in solution, when a similar cycle of energy changes
must be considered, but the hydration energies of the positive and negative ions must be substituted
for the lattice energy.
Structures of Metals, Hardness and Cohesive Energy
At normal temperatures all the Group 1 metals adopt a body-centred cubic type of lattice with
a coordination number of 8. However, at very low temperatures lithium forms hexagonal close-packed
structure with a coordination number of 12.
The metals are all very soft, and can be cut quite easily with a knife. Lithium is harder than
the others, but is softer than lead.
The cohesive energy is the force holding the atoms or ions together in the solid. (This is the
same in magnitude, but the opposite in sign, to the enthalpy of atomization, which is the energy
required to break the solid up into gaseous atoms.) The cohesive energies of Group l-metals are about
half of those for Group 2, and one third of those for Group 13 elements. The magnitude of the
cohesive energy determines the hardness, and it depends on the number of electrons that can
participate in bonding and on the strength of the bonds formed. The softness, low cohesive energy and
weak bonding in Group 1 elements are consequences of these metals having only one valency electron
which can participate in bonding (compared with two or more electrons in most other metals), and of
the large size and diffuse nature of the outer bonding electron. The atoms become larger on
descending the group from lithium to caesium, so the bonds are weaker, the cohesive energy
decreases and the softness of the metals increases.
Table: Cohesive energy
Li
Na
Cohesive energy
(enthalpy of atomization) (kJ mol-1)
161
108
[11]
Chemistry
K
Rb
Cs
90
82
78
Melting and Boiling Point
The generally low values for cohesive energy are reflected in the very low values of melting
and boiling points in the group. The cohesive energy decreases down the group, and the melting
points decrease corresponding.
The melting points range from lithium 181oC to caesium 28.5oC. These are extremely low
values for metals, and contrast with the melting points of the transition metals, most of which are
above 1000oC.
The melting point of lithium is nearly twice as high (in oC) as that for sodium.
Melting and Boiling points
Melting point
Li
Na
K
Rb
Cs
Boiling point
(oC)
1347
881
766
688
705
181
98
63
39
28.5
Flame Coloration: The colour actually arises from electronic transitions in short-lived
species which are formed momentarily in the flame. The flame is rich in electrons, and in the case of
soldium the ions are temporaliy reduced to atoms.
Na  Na   e
The sodium D-line (which is actually a doublet at 589.0 nm and 589.6nm)arises from the
electronic transition
3s1  3 p1 in sodium atoms formed in the flame.
Colour
Li
Na
K
Rb
Cs
Wavelength(nm)
670
589
766
780
455
Crimson
Yellow
Lilac
Red-violet
Blue
Flame colours and wavelengths
Lattice Energies: Salts of alkali metals consist of cations and anions only and are, therefore,
called ionic solids. The lattice energy of an ionic solid is defined as the energy relased when cations
and anions in gaseous state are brought from infinity to their respective crystal sites to form one mole
of the ionic crystal.
Li  g   Cl  g   Li  Cl  s ; H  Lattice energy of Li  Cl 
Cs  g   Cl  g   Cs  Cl  s ; H  Lattice energy of Cs  Cl 
Reaction with water
Group 1 metals all react with water, liberating hydrogen and forming the hydroxides. The
reaction becomes increasingly violent on descending the group. Thus lithium reacts gently, sodium
melts on the surface of the water and the molten metal skates about vigorously and may catch fire and
potassium melts and always catches fire.
2Li + 2H2O → 2LiOH + H2
2Na + 2H2O → 2NaOH + H2
[12]
Chemistry
2K + 2H2O → 2KOH + H2
The standard electrode potentials Eo are Li+ | Li = -3.05 volts, Na+ | Na = -2.71, K+ | K = -2.93,
Rb | Rb = 2.92, Cs+ | Cs = -2.92. Lithium has the most negative standard electrode potential of any
element in the periodic table, largely because of its high hydration energy. Standard electrode
potentials Eo and Gibbs free energy G are related by the equation.
G  nFE o
Where n is the number of electrons removed from the metal to produce the ion, and F is the
Faraday constant.
The reaction Li+ + e → Li has the largest negative Eo value, and hence the largest positive
G value. Thus the reaction does not occur. However, the reverse reaction Li+ → Li+ + e has a large
negative value of G, so lithium liberates more energy than the other metals when it reacts with
water. In view of this it is at first sight rather surprising that lithium reacts gently with water, whereas
potassium, which liberates less energy, reacts violently and catches fire. The explanation lies in the
kinetics (that is the rate at which the reaction proceeds) rather than in the thermodynamics (that is the
total amount of energy liberated). Potassium has a low melting point, and the heat of reaction is
sufficient to make it melt, or even vaporize. The molten metal spreads out, and exposes a larger
surface to the water, so it reacts even faster, gets even hotter and catches fire.
+
Reaction with dinitrogen
Lithium is the only element in the group that reacts with dinitrogen to form a nitride. Lithium
nitride, Li3N, is ionic (3Li+ and N3-), and is ruby red. Two reactions of the nitride are of interest. First,
on heating to a high temperature it decomposes to the elements, and second, it reacts with water,
giving ammonia.
2Li3N
6Li + N2
Li3N + 3H2O → 3LiOH + NH3
OXIDES, HYDROXIDES, PEROXIDES AND SUPEROXIDES
Reaction with air
The metals all burn in air to form oxides, though the product varies depending on the metal.
Lithium forms the monoxide Li2O (and some peroxide Li2O2), sodium forms the peroxide Na2O2 (and
some monoxide Na2O), and the others form superoxides of the type MO2.
All five metals can be induced to form the normal oxide, peroxide or superoxide by
dissolving the metal in liquid ammonia and bubbling in the, appropriate amount of dioxygen.
Normal oxides - monoxides
The monoxides are ionic, for example 2Li+ and O2-Li2O and Na2O are pure white solids as
expected, but surprisingly K2O is pale yellow, Rb2O is bright yellow and Cs2O is orange. Metallix
oxides are usually basic. The typical oxides M2O are strongly basic oxides, and they react with water,
forming strong bases.
Li2O + H2O → 2LiOH
Na2O + H2O → 2NaOH
K2O + H2O → 2KOH
The crystal structures of Li2O, Na2O, K2O and Rb2O are anti-fluorite structures. The antifluorite structure is like that for fluorite CaF2, except that the positions of the positive and negative
ions are interchanged. Thus Li+ fill the sites occupied by F- and O2- fill sites occupied by Ca2+. Cs2O
has an anti-CdCl2 layer structure.
[13]
Chemistry
Hydroxides
Sodium hydroxide NaOH is often called caustic soda, and potassium hydroxide is called
caustic potash, because of their corrosive properties (for example on glass or on skin). These caustic
alkalis are the strongest bases known in aqueous solution. The hydroxides of Na. K, Rb and Cs are
very soluble in water, but LiOH is much less soluble.
Solubility of Group 1 hydroxides
Element
Solubility (g/100g H O)
Li
13.0
Na
108.3
K
112.8
Rb
197.6
Cs
385.6
At 25oC a saturated solution of NaOH is about 27 molar, whilst saturated LiOH is only about
5 molar. The bases react with acids to form salts and water, and are used for many neutralizations.
The bases also react with CO2 even traces in the air, forming the carbonate. LiOH is used to
absorb carbon dioxide in closed environments such as space capsules (where its light weight is an
advantage in reducing the launching weight).
2NaOH + CO2 → Na2CO3 + H2O
They also react with the amphoteric oxides, Al2O3 forming aluminates, SiO2 (or glass)
forming silicates, SnO2 forming stannates, PbO2 forming plumbates and ZnO forming zincates.
The bases liberate ammonia from both ammonium salts and coordination complexes where
ammonia is attached to a transition metal ion (ammine complexes).
NaOH + NH4Cl → Nh3 + NaCl + H2O
6NaOH + 2[Co(NH3)6] Cl3 → 12Nh3 + Co2O3 + 3NaCl + 3H2O
Hexamine
cobalt(III) chloride
NaOH reacts with H2S to form sulphides S 2 , and hydrogen sulphides SH-, and it is used to
remove mercaptans from petroleum products.
NaOH + H2S → NaSH → Na2S
The hydroxides react with alcohols, forming alkoxides.
NaOH + EtOH → NaOEt + H2O
Sodium ethoxide
KOH resembles NaOH in all its reactions, but as KOH is much more expensive it is seldom
used. However, KOH is much more soluble in alcohol, thus producing Oc2H5- ions by the equilibrium.
C2H5OH + OH- OC2H5- + H2O
Peroxides and superoxides
The peroxides all contain the [−O−O−]2- ion. They are diamagnetic (all the electrons are
paired), and are oxidizing agents. They may be regarded as salts of the dibasic acid H2O2 and they
react with water and acid giving hydrogen peroxide H2O2.
Na2O2  2H 2O  2 NaOH  H 2O2
Na2O2 is pale yellow in colour: It is used industrially for bleaching wood pulp, paper and
fabrics such as cotton and linen. It is a powerful oxidant, and many of its reactions are dangerously
violent, particularly with materials that are reducing agents such as aluminium powder, charcoal,
sulphur and many organic liquids. Because it reacts with CO2 in the air it has been used to purify the
air in submarines and confined spaces, as it both removes CO 2 and produces O2. Potassium
superoxide KO2 is even better for this purpose. Some typical reactions are.
Na2O2 + Al → Al2O3
Na2O2 + Cr3+ → CrO24Na2O2 + Co → Na2CO3
2Na2O2 + 2Co2 → Na2CO3 + O2
[14]
Chemistry
The industrial process for forming sodium peroxide is a two-stage reaction in the presence of
excess air.
1
2 Na  O2  Na2O
2
1
Na2O  O2  Na2O2
2
The superoxides contain the ion [O2]-, which has a unpaid electron and hence they are
paramagnetic and are all coloured (LiO2 and NaO2 yellow, KO2 orange, RbO2 brown and CsO2
orange).
The normal oxides contain O2- ion, the peroxides contain O22- ion and superoxides contain O-2
ion. The peroxides and superoxides become more stable with increase in atomic number of the alkali
metal. The formation and stability of these oxides can be explained on the basis of their lattice energy
Li+ ion being a small ion has a strong positive field around it and can stabilize only a small anion, O 2-,
whereas Na+ ion being a large cation can stabilize a large anion and so on. Thus, smaller cation can
stabilize smaller anion and larger cation can stabilize larger anion.
The higher oxides, viz., peroxides and superoxides are important oxidizing agents. They react
with dilute acids forming hydrogen peroxide and oxygen, respectively.
Na2O2 + H2SO4 → Na2SO4 + H2O2
4KO2 + 2H2SO4 → 2K2SO4 + 2H2O + 3O2
Superoxides are even stronger oxidizing agents than peroxides, and give both H2O2 andO2
with either water or acids.
1
KO2 + 2H2O → KOH + H2O2 +
O2
2
KO2 is used in space capsules, submarines, and breathing masks, because it both produces
dioxygen and removes carbon dioxide. Both functions are important in life support systems.
4KO2 + 2CO2 → 2K2CO3 + 3O2
4KO2 + 4CO2 + 2H2O
4KHCO3 + 3O2
Sodium superoxide cannot be prepares by burning the metal in dioxygen at atmospheric
pressure, but it is made commercially and in good yields by reacting sodium peroxide with dioxygen
at a high temperature and pressure (450oC and 300 atmospheres) in a stainless steel bomb.
Na2O2 + O2 → 2 NaO2
Sodium Bicarbonate, NaHCO3
NaHCO3 can be used on its own to make cakes or bread 'rise' since it decomposes between
50oC and 100oC giving bubbles of CO2.
2NaHCO3
Na2CO3 + H2O + CO2
Baking powder is more commonly used, and contains NaHCO3, Ca(H2PO4)2 and starch. The
Ca(H2PO4)2 is acidic and when moistened it reacts with NaHCO, giving CO2. The starch is a filler. An
improved combination baking powder contains about 40% starch, 30% NaHCO3, 20% NaAl(So4)2
and 10% Ca(H2PO4)2. The NaAl(So4)2 slows the reaction down so the CO2 is given off more slowly.
Oxosalts-Carbonates, Bicarbonates, Nitrates and Nitrites
Group 1 metals are highly electropositive and thus form very strong bases, and have quite
stable oxosalts.
[15]
Chemistry
The carbonates are remarkably stable. And will melt before they eventually decompose into
oxides at temperatures above 1000oC. Li2CO3 is considerably less stable and decomposes more
readily.
Because Group 1 metals are so strongly basic, they also form solid bicarbonates (also called
hydrogen carbonates). No other metals form solid bicarbonates, though NH4HCO3 also exists as a
solid. Bicarbonates evolve carbon dioxide and turn into carbonates on gentle warming. This is one test
for bicarbonates in qualitative analysis. The crystal structures of NaHCO3 and KHCO3 both show
hydrogen bonding, but are different. In NaHCO3 the HCO-3 a ions are linked into an infinite chain,
whilst in KHCO3 a dimeric anion is formed.
Lithium is exceptional in that it does not form a solid bicarbonate, though LiHCO3 can exist
in solution. All the carbonates and bicarbonates are soluble in water.
Over 50000 tonnes of Li2CO3 are produced annually. Most of it is added as an impurity to
Al2O3 to lower its melting point in the extraction of aluminium by electrolysis. Some is used to
toughen glass (sodium in the glass is replaced by lithium). Na2CO3 is used as washing soda to soften
water in hard water areas, and NaHCO3 is used as baking powder.
The nitrates can all be prepared by the action of HNO3 on the corresponding carbonate or
hydroxide, and they are all very soluble in water. LiNO3 is used for fireworks and red-coloured
distress flares. Large deposits of NaNO3 are found in Chile, and are used a nitrogenous fertilizer.
Solid LiNO3 and NaNO3 are deliquescent, and because of this KNO3 is used in preference to NaNO3
is gunpowder (gunpowder is a mixture of KNO3, sulphur and charcoal). KNO3 is usually obtained
from synthetic nitric acid and K2CO3, but at one time it was made from NaNO3.
2HNO3 + K2CO3 → 2 KNO3 + CO2 + H2O
NaNO3 + KCl → KNO3 + NaCl
Group 1 nitrates are fairly low melting solids, and are amongst the most stable nitrates known.
However, on strong heating they decompose into nitrites, and at higher temperatures to the oxide.
LiNO3 decomposes more readily than the others, forming the oxide.
2NaNO3 2NaNO2 + O2
4NaNO3 2Na2O + 5O2 + 2N2
Halides and Polyhalids
Since Li+ is the smallest ion in the group, it would be expected to form hydrated salts more
readily than the other metals. LiCl, LiBr and Lil form trihydrates LiX.3H2O, but the other alkali metal
halides form anhydrous crystals.
All the halides adopt a NaCl type of structure with a coordination number of 6 except for
CsCl, CsBr and CsI. The latter have a CsCl type of structure with a coordination number of 8. Rather
more compounds adopt the NaCl type of structure than would be expected from the radius ratios of
the ions r+ / r-, and the reason for this structure being adopted is that it gives the highest lattice energy.
The alkali metal halides react with the halogens and interhalogen compounds forming ionic
polyhalide compounds.
Kl + l2 → K [l3]
KBr + |C| → K [BrlCl]
KF + BrF3 → K [BrF4]
Hydrides
Group 1 metals all react with hydrogen, forming ionic or salt-like hydrides M+H-. However,
the ease with which they do so decreases from lithium to caesium. These hydrides contain the H- ion.
It can be proved that H- ions exist because on electrolysis hydrogen is liberated at the anode.
The hydrides react with water, liberating hydrogen, and lithium hydride is used as a source of
hydrogen for military purposes and for filling meteorological balloons.
LiH + H2O → LiOH + H2
[16]
Chemistry
Lithium also forms a complex hydride Li[AlH4], called lithium aluminium hydride, which is a
useful reducing agent. It is made from lithium hydride in dry ether solution.
4LiH + AlCl3 → Li[AlH4] + 3LiCl
Lithium aluminium hydride is ionic, and the [AlH4]- ion is tetrahedral. Li[AlH4] is a powerful
reducing agent and is widely used in organic chemistry, as it reduces carbonyl compounds to alcohols.
It reacts violently with water, so it is necessary to use absolutely dry organic solvents, for example
ether which has been dried over sodium. Li[AlH4] will also reduce a number of inorganic compounds.
BCL3 + Li[AlH4] → B2H6
diborane
PCL3 + Li[AlH4] → PH3
phosphine
SiCl4 + Li[AlH4] → SiH4
silane
Sodium tetrahydridoborate (sodium borohydride) Na[BH4] is another hydride complex. It is
ironic, comprising tetrahedral [BH4]- ions. It is best obtained by heating sodium hydride with
trimethyl borate.
4NaH + B(OCH3)3
Na[BH4] + 3NaOCH3
Solubilty and Hydration
All the simple salts dissolve in water, producing ions, and consequently the solutions conduct
electricity. Since Li+ ions are small, it might be expected that solutions of lithium salts would conduct
electricity better than solutions of the same concentration of sodium, potassium, rubidium or caesium
salts. The small ions should migrate more easily towards the cathode, and thus conduct more than the
larger ions. However, ionic mobility or conductivity measurements in aqueous solution give results in
the opposite order Cs+ > Rb+ > K+ > Na+ > Li+. The reason for this apparent anomaly is that the ions
are hydrated in solution. Since Li+ is very small, it is heavily hydrated. This makes the radius of the
hydrated ion large, and hence it moves only slowly. In contrast, Cs + is the least hydrated, and the
radius of the hydrated Cs+ ion is smaller than the radius of hydrated Li+, and hence hydrated Cs+
moves faster, and conducts electricity more readily.
The size of the hydrated ions is an important factor affecting the passage of these ions through
cell walls, It also explains their behaviour on cation exchange columns, where hydrate Li + ions are
attached less strongly and hence eluted first.
The decreases in hydration from Li+ to Cs+ is also shown in the crystalline salts, for nearly all
lithium salts are hydrated, commonly as trihydrates. In these hydrated Li salts Li+ is coordinated to
6H2O, and the octahedra share faces, forming chains. Many sodium salts are hydrated, e.g. Na 2CO3.
10H2O, Na2CO. 7H2O and Na2CO3.H2O. Few potassium salts and no rubidium or caesium salts are
hydrted.
The simple salts are all soluble in water, and so in qualitative analysis these metals need to be
precipitated as less common salts. Thus Na+ is precipitated by adding zinc (or copper) uranyl acetate
solution and precipitating NaZn(UO2)(Ac)9. H2O sodium zinc uranyl acetate. K+ is precipitated by
adding a solution of sodium cobaltinitrite and precipitating potassium cobaltinitrite K 3[Co(NO2)6] or
by adding perchloric acid and precipitating potassium perchlorate KCIO4. Group 1 metals can be
estimated gravimetrically, sodium as the uranylacetate and potassium, rubidium and caesium as
tetraphenylborates. However, modern instrumental methods such as flame photometry and atomic
absorption spectrometry are much quicker and easier to use and are now used in preference to
gravimetric analysis.
K   Na3 CONO2 6   Na   K3 CONO2 6 


K  NaClO4  Na  KClO4
K  NaBC6 H 5 4   Na   K BC6 H 5 4 
potassium cobaltinitrite
potassium perchlorate

potassium
tetraphenylborate
quantitative precipitate
If a salt is insoluble its lattice energy is greater than the hydration energy. K[B(C 6H5)4] is
insoluble because the hydration energy is very small as a result of the large size of its ions.
[17]
Chemistry
The solubility of most of the salts of Group 1 elements in water decreases on descending the
group. For a substance to dissolve the energy evolved when the ions are hydrated (hydration energy)
must be larger than the energy required to break the crystal lattice (lattice energy). Conversely, if the
solid is insoluble, the hydration energy is less than the lattice energy.
The reason why the solubility of most Group 1 metals decreases on descending the group is
that the lattice energy only changes slightly, but the free energy of hydration changes rather more. For
example, the difference in lattice energy between NaCl and KCl is 67 kJ mol -1, and yet the difference
in G (hydration) for Na+ and K+ is 76 kJ mol-1. Thus 1 is less soluble than NaCl.
The Group 1 fluorides and carbonates are exceptional in that their solubilities increase rapidly
on descending the group. The reason for this is that their lattice energies change more than the
hydration energies on descending the group. The lattice energy depends on electrostatic attraction
between ions, and is proportional to the distance between the ions, that is proportional, to I/(r +/r-). It
follows that the lattice energy will vary most when r- is small, that is with F-, and will vary least when
r+ is large (with I-). The weight if solute dissolving does not provide a very useful comparison of the
solubilities, because the molecular weights differ. The easiest way to compare the number of ions is to
compare the solubilities as molar quantities.
Reactions with NH3 (gaseous): When ammonia gas is passed over alkali metals, the
corresponding amides are formed.
2M  2 NH 3  2MNH 2  H 2
The alkali amides are also formed when the solutions of alkali metals in liquid NH3 are
evaporated.
Solubility in Mercury: The alkali metals dissolve readily in mercury forming what are called
amalgams. The process is highly exothermic.
Solubility in Liquid Ammonia: All the alkali metals dissolve in anhydrous liquid ammonia
giving deep blue solutions. The colour depends with increase in concentration of the solution. The
concentrated solutions are bronze in colour and posses metallic luster. The solutions are good
conductors of electricity and are paramagnetic in nature. The paramagnetic character decreases with
increase in concentration.
The dissolution of the metal is accompanied by its ionization.
M → M+ + eWhen the alkali metals are dissolved in liquid ammonia, the ammonia molecules solvate the
cation as well as the electron, as represented below.
M + NH3 (in excess) → M+ (NH3)x + e- (NH3)y
The values of x and y depend upon the extent of salvation of the alkali metal ion and the
electron, respectively. The exact mechanism for the salvation of electrons is not clear. It is argued,
however, that the electron released by the alkali metal causes polarization of the charge of the
electrons of the surrounding ammonia molecules causes a 'hole' in which the released electron is
trapped. It is the ammoniated electron which is responsible for the blue colour of the solution. The
electrical conductivity is due to the ammoniated cation as well as the ammoniated electron. The dilute
solutions are paramagnetic because they contain free electrons. The Decrease in paramagnetic
character with increase in concentration suggests that in concentrated solutions the ammoniated
electrons associate to form electron pairs.
2e- (NH3)y → [e- (NH3)y]2
If blue colour is allowed to stand, the colour slowly fades until it disappears owing to
formation of metalamide. At concentration above 3M, solutions are copper-bronze coloured and have
a metallic luster because metal clusters are formed. In the presence of impurities or catalysts such as
Fe, the alkalie metals react with NH3 to form metal amide and hydrogen.
M + NH3 → MNH2 + H2
If all impurities or catalysts are absent, the group 1 metals dissolve in liquid ammonia.
The alkali metal solutions in liquid ammonia are stable and becompose only slowly liberating
hydrogen. The presence of transition metals (such as Fe, Pt, Zn, etc.) even in traces, catalyses the
reaction.
[18]
Chemistry
Due to the presence of the free or ammoniated electrons, these solutions behave as strong
reducing agents as is evident from the following reactions.
These solutions reduce aerial oxygen to superoxide and peroxide ions without causing
cleavage of O–O bond.
These solutions reduce metal ions to their unusually low oxidation states.
K2[Ni(CN)4] + 2K
Mn2(CO)10 + 2K
Fe(CO)5 + 2Na
K4[Ni(CN)4]
2K[Mn(CO)5]
Na2[Fe(CO)4] +CO
(Ni is in 0 oxidation state)
(Mn is in -1 oxidation state)
(Fe is in –2 oxidation state)
Ionic Conductance: The alkali metal ions in aqueous solution have high electrical
conductance. The electrical conductance of Li+ ion, which is the smallest of alkali metal ions, is
expected to be the highest. Since ionic size increases on moving from Li+ to Cs. the ionic conductance
is expected to decrease in the order.
Li+ > Na+ > K+ > Rb+ > Cs+
However, we find by experiment that ionic conductance of Cs+ is the highest and that of Li+ is
the lowest. Thus, ionic conductance actually decreases in the order.
Cs+ > Rb+ > K+ > Na+ > Li+
This is evidently the reverse of what we expect. This anomaly arises from the fact that alkali
metal ions are hydrated to different extents. The Li+ ion is hydrated to the maximum extent due to its
high charge/size ratio compared to other alkali metal ions. The extent of hydration decreases on
moving from Li+ to Cs+. Thus, Cs+ ion is the least hydrated. As a result of the variation in the extent of
hydration, Li+ ion due to heavy hydration is the biggest while Cs+ ion due to much smaller hydration
is the smallest. Thus, the ionic size in aqueous solutions, instead of increasing, actually decreases on
moving from Li+ to Cs+ as a result of which the ionic mobility increases as we move from Li+ to Cs+.
Hence, the ionic conductance, instead of decreasing, actually increases on moving from Li+ to Cs+.
The ionic radii, radii of hydrated ions and ionic conductances of different alkali metal ions at
infinite dilution are given in following table.
Role of Na+ and K+ ions in Biological Systems: Sodium cation is a major cation of
extracellular fluids of animals including human beings which is known to activate certain enzymes in
the animal body. Sodium ions are relatively harmless except when present in excess amounts in which
case they may cause hypertension. The saline water containing excessive amounts of NaCl is harmful
to plant and aquatic life because of the toxicity of Na+ ions.
The potassium ion is essential to all organisms with the possible exception of blue green
algae. It is a major cation of intracellular fluid of animal cells. It is moderately toxic to mammals but
only when injected intravenously.
The sodium and potassium ions, taken together, constitute what is known as Na+ - K+ pump or
simply as 'sodium pump'. The functioning of this pump is actually a biological process which goes on
occurring in each and every cell of all animals. It transfers Na+ ions from the intracellular fluids to the
extra cellular fluids with the help of 'carrier proteins'. Simultaneously it transfers K+ ions from the
extra cellular fluids to the intracellular fluids. Since each operation of the pump (i.e., each cycle of
this biological process) pumps out larger number of sodium ions from the cell than the number of K +
ions that it pumps into the cell, the interior of the cell acquires an excess negative charge and the
exterior of the cell acquires an excess positive charge. This results in the development of electrical
potential gradient across the cell membrane which is responsible for the transmission of nerve signals
in animals. The Na+ – K+ pump also maintains the volume of the cell. Without the existence of the
Na+ – K+ pump in the cell, the latter would have swelled in volume and ultimately bursted.
[19]
Chemistry