Electronic structure of atoms, modern periodic table and periodic law, variation of periodic properties within periods and groups, ionization potential, electron affinity, electronegativity, usefulness and limitation of periodic table.

1. Written By
Department of Pharmacy
University of Rajshahi
Rajshahi-6205, Bangladesh

2. t ea a 1 meta s iave t ieir sum ar va ence shell con iguration an hence have similar properties.
Element Atomic no. Electronic configuration
.. -
Valence shell
electronic configuration
LI 3 Isi 2s1
•.
2s
I
Na II ts· 2sl 2p6 3s' I 1.·
3s
K 19 ls2 2si 2p6 3s2 3pf> 4s1
4s
I ! •.
Rb 37 ls2 2s2 2p6 3s2 3p6 3d11' 4s· 4pb5s1
5s
t
. '
Merits of long form of periodic table [Mosley's law] over Mendeleef's periodic table '"• · .,·, t!···i ... :•.
The long form of periodic table has a lot of merits over the Mendeleef's periodic table in the following respects;'
I. . The classification of the elements is based on a more fundamental property like atomic number.x'. 1r:1r-:·
2. Each group contains elements with similar electronic configuration and hence similar.properties, Fot example. all
lk r I I I . ·1 I fi dh
1., ",
Mosley's Modern Periodic Law: Mosley's modem periodic law may be! stated as follows: . :! · '. .,,
The properties of ~he elements are a periodic function of their atomic numbers that is if the 'Clements are arranged
in the increasing order of their atomic numbers, the properties of the. similar elements are repeated after deflulte
regular intervals or periods. · ·,,
,; 1. .:·
Law of octaves: In I 864. John Newland» arranged the clements in the increasing order of their atomic wcights-ilnd
observed that. every cighthelements in the list beginning from any given element showed a rcpitition of the physical and
chemical .properties of the first element. Because of its resemblance with the eight notes in musical scale, this concept of the
classification was known as the law ofoctaves. ·
::>: • Limitations: Unfortunately this concept was not well received by scientist of Newland's time. It may be pointed out
that. the inert gases were not then known and the discovery of the inert gases completely shelved 1~~ "fl'!I the idea of the
law of octaves.
.,,
5. Position of isotopes: If the clements are arranged in the order or their increasing atomic weights, it is not possible to
accommodate large number of isotopes in the periodic table.
6. Group dolls not represent valency: Excepting osmium. elements placed in-group eight. do not show a valency of
x Also the clements lying in the middle of long periods show two or more valencies. for example: Cr, Mn etc.. ' . . ~ .
Defects of Mendeleef s periodic table: • ,
The defects of Mendeleef's periodic tablearc given below:
I. Position of hydrogen: Hydrogen resembles 1~ ~-.cm1 both the alkali metals and the halogens. Its position in the
periodic table is anomalous (~1 .
.2. Position of lanthanides and actinides: A group or 15 elements (At. no. 57 to 71 ). which is called rare earths or
lanthanides docs not find its proper place in the table and has been placed at one place in group Ill and period 6. Similarly.
, another group of 15 elements (At. no. 89 to I 03 ). called actinides does not find its proper place and has been placed at one
place in group Ill and period 7.
3. Similar elements are separated while dissimilar elements are placed in the same group: Similar elements arc
placed in different groups in the periodic table. As for example. Cu and Hg; Ba and Pb: Ag and Tl (Thallium). while many
dissimilar elements has been grouped together. As for example. Cu. Ag and Au arc grouped along with alkali metals.
4. Existence of four anomalous pairs of elements: The order of increasing atomic weight has been ignored in c,..,l. or
four pairs of elements in order to place them in a position justi lied by their properties. Thus elements of higher atomic 11 eight
proceeds those of lower atomic weight atfour places as shown below: _
· (a) Ar (Z= 18. at. wt.=40) proceeds K (Z·' 19, at. wt.=39)
(h) Co (Z=.27, at. wt.=5lJ.9) proceeds Ni (7'"° 28. at. Wt."'58.60)
(c) Te (h=52, at. wt."'127.6) proceeds I (Z--53. at. wt.r- 126.9)
(d) Th (Z=90. at. wt.=2'32.12) proceeds Pa (Z-c9). at. wt. -231)
'.· 1.j.
Working on this law. Mendelecf set up the clements in till· incrcasi ng order or their ah unic weights' in the form of a table.
which is known as Mcndclecf'x periodic table. In this table, the clements arc arranged in groups and periods'.
Mendeleef's Periodic law:
Many attempts were made to classify the clements. Mendelcef gave a law known as Mcndeleef's periodic law. which 1
stated as: , · ·''' , , .
The properties of the elements are a periodic function of their atomic weights that is if the elements are arranged
in the increasing order of their atomic weights, the properties of the similar elements arr repeated after dcfluite
regular intervals or periods. ' ' ·· ·· .: ·
Periodic Table
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3. a s o one an wit 1 respect to oxygen increases rem to ass iown c ow:
Elements of second period Li Be B c N 0 F
Hvdrides of the elements LiH BeH2 BH-' CH~ NH-' H20 HF
Val ency of the elements with respect to H2
I 2 3 4 3 2 I
6eneral character1st1cs ot: Per1oas
I. Number of valency electrons: Number of valency electrons increases from I to 8 when .we proceed from left to
right in a period.
2. Valency: The valency of the elements with respect to hydrogen in each short period increases from I to 4 and then
f; II t d . I f I 7 I b I
i?roup
Elements Electronic configuration Valence shell .elee~11i l11figt1r111iaR
Li z, I 2 - .,
Na 2, 8, I 3
K 2, 8, 8, I 4
Rh 2, 8, 18, 8, I 5
Cs 2, 8, 18, 18, 8, I 6
Fr 2, 8, 18, 32, 18, 8, I 7
.. •---a-!~1"-1 . -
4. Size of atoms: Size of atoms increases on descending a group. for example in-group IA. atomic size increase- from
Li to Cs. ·
5. Metallic character: The metallic character of the clements increases in moving from top to bottom in a group. This
is particularly apparent in groups IVA. VA and VIA. which begin with non-metals (namely C. N and 0 respectively). and
end with metals (namely Pb, Bi and Po respectively). -
6. Number of electron shells: In going down a group the number of electron shells increases by one at each step and
ultimately becomes equal to the number of the period to which the elements belongs as shown below for the elements of
IA
2. Valency: The valencies of all the elements of the same group are the same. ; r · , , .
3. Properties of elements: All the elements 'of a given group possess very similar physical and chemical properties,
There is a regular grad/adon 1ar11111f"l-of'41 in their properties when we move from top to bottomin a group. For example.
(a) The alkali metals resemble each other and their base-forming tendency increases from 1.i to Cs.
(h) The reactivity of the halogens decreases as we pass from F to I.
or cxamp e, t e e ernents o J?roup
Element Atomic 110.1 Electronic configuration Valence shell
electronic confiuuraiion
Li 3 ·ls' 2s' 2s
I
l'ia II Is· 2s" 2p6 3s1
3s
I
K 19 ls1 2s1 2p6 3s' 3p6 4s1
4s
I
Rb, 37 Is' ls' 2pb 3s' 3p" 3~ "' 4s' 4p0 5s' Ss
I
•
Defects of loni: form of periodic table or Modern periodic table or Mosley's periodic tahl«
Although the long form of periodic table is superior to Mcndclccf's periodic table in 111ai1y respects. it retains some of the
detects such as:
I. The problem of the position of the hydrogen still remains unsolved.
2. It fails to accommodate the lanthanides and actinides in the main body of the table.
3. The arrangement is unable to reflect the electronic configuration of many elements.
General charactel'is.ti£mgrougs ~ lAA_ltf 1l 1-
1. Number of valency electrons: On moving down a given group the number of valency electrons does not change.
F I h I f IA .
3. It explains the similarities and variations in the properties of the clements in terms of their electronic configurations.
4. The inert gases having completely ti lied electron shells have been placed at the end ofeach period.
5. In this form of the periodic table. the elements of the two sub-groups have been placed separately and thus dissimilar
elements do not fall together. ·
6. It provides a clear dernarcation pttst C!'!~1 of different types of the elements like active metals, transition metals. non-
metals. metalloids, inert gases. lanthanidcs and actinides.
7. It is easier to remember. understand and reproduce.
7s1Is 2s 2p 3s 3p 3dFr
2
6s1Is" 2s 2p 3s 3p 3d
87
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4. .~ The examination of the properties of these elements will show that the elements belonging to the same group have
similar properties. ·
IA (Alkali metals). zero (Inert gases) and VllA (Halogens) niven below:
Periodic number Group IA Group zero GrQUP vllA
I H (I) He (2)
2 Li (3) Ne(IO) F (9)
3 Na(ll) Ar ( 18) Cl (17) ..4 K (19) Kr (36) Br (35) '
5 Rb (37) x- (54) I (53)
6 Cs (55) Rn (86) At (85)
7 Fr (87)
3rd period=-s
fadicit~·of the properties . •
The repetition of the elements with similar properties in the order of increasing atomic number in the periodic table is
called periodicity of properties.
In order to understand the concept of periodicity of properties we may consider the properties of the elements of group
8. Diagonal relationship: It is one of the important features 1~1 of the earlier elements of 2nd period of the long
form of the periodic table that they resemble more with the elements lying at their right hand side in the 3rd period. Thus Li
resembles more in its properties with Mg. which lies at its lower right in the 3rd period. Similarly Be shows similarities with
Al and B resembles Si. The two elements showing similarities in properties are called diagonal pair and the similarity in
properties existing between two elements of a diagonal pair is known as diagonal relationship.
Three diagonal pairs like Li - Mg. Be - Al and B - Si is shown below:
Group+-e IA llA IIIA Iv A
2nd period ~ Li Be B C
N~M~A~Si
7. Number ofshells: In going from lc::ft to right in a perio the number of ~n-a shell remains same:
Elements of 2nd period Li Be B c N '
0 ·-- F
... Ne
Atomic number 3 4 5 6 7 8 9 IO
Electronic configuration 2, l 2, 2 2,3 2,4 2, 5 2,6 2, 7 2, 8
No. of shells 2 2 2 2 2 2 2 2
d
C cmcnts ecome exx an css asic in 11c same c rrecuon. -or cxarnp c:
Oxides of the Na20 MgO Al20.1 Si02 P20s S0.1 Cl201
-clemcnts
of 3rd period
Basic character . Strongly basic Basic Amphoteric Feebly Acidic More acidic Most acidic
acidic
..,J
,.,y..o
Acidity and alkalinity: The gradual decrease of the metal I ic character from ldl to right ~hat the oxides or the
h I di b .. I r r I
6.
Metals Non-metals
-Metallic character decrcasing-»Metallic character:
3. Size of atoms: Size of atoms decreases from left to right in a period. Thus alkali metals have the largest size while
halogens have the smallest size.
4. Properties of the elements: The properties of the clements of a given period differ considerably hut the elements in
the two adjacent 1~i1 periods show marked similarity between them. For example, when we consider the clements or 2nd
and 3rd periods. we find that Na resembles Li. Mg resembles Be. Al resembles I3 etc. ·
5. Metallic character: On moving from left to right in a period the metallic character of the elements decreases. For
example in 3rd period. Na. Mg and Al arc metals while Si. P. Sand Cl arc non-metals.
Elements of 3rd period: Na, Mg, Al Si, P, S, Cl
Elements of 3rd period Na Mg Al Si p s· Cl I'•
Oxides of the elements Na20 MgO AhO, Si02 P20, so, Ch01
Valency of the elements with I 2 3 4 5 6 7
respect to 0 ~
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5. l
I
•
- ~ Q.7o R·D - - .'
An anion is bigger in size than its parent atom: We know that an anion is formed by adding one or more elcclrons to
the neutral atom. T.Jll,llian anion has more electrons than its parent atom. With lhe jnrreas" gt:, the--nYf»~uulectrons Ilic
magnitude of screening c"OilStant:<Jalso increases. Thei~~re~~c in th;~iagnitude of o ~1s~s tl1e m(!gnitude ~ctive
~· The decreased effective nuclear charg~ulls the electron cloud of the anion outward 1~ ~' awav from
the nucleus and thus makes the anion larger in size tharilt0?.arcnt atom.
Thus halide ions are bigger in size than the halo!!en atoms as shown below
/
Alkali metals Atomic radius (A) Cations · Radius of cation (A)
Li (3) I. 23 Li+ (2) 0.60
Na (11) ' 1. 54 Na+ (10) 0.95
K (19) 2.03 K + (18) I. 33
"
Value of n:
Value ofl:
._;f'"··
,.! ' ~{~~~1'~f ;he o~bital o~~~pied (lf'IWf TIT) by the electrons: When the penetraiio!1 l~I p6~~f:;,d~cre~~c~ 1li·'ihe
orbitals. the: shielding effect is also decreases. The penetration power of up electron is greater than that of an nd electron
(np>nd). Again penetration power of nd electron is greater than that of an nf electron (ndxnf), Thus the penetration power of
an electron inns. np, nd andnf orbitals of a given nth is in the following order: · .
---Penetration power decreasing---4
ns >np >nd >nf
~;-(~n( 2< 3
---:Shielding effect decreasing _ ,,! 1, .i~.1
Size of cation, atom and anion ,....__, p._,D ~ <26~ , _ · r , - .
A cation is s~aller in size than its parent atom: We know tlJ<~ F.,_atl~~j~i.mncd by the removal of one or more
electrons from the atom. Tllw:;. a cation has lesser' number of electrons than its parent atom. With the decrcai;e,.of the-number
Qf electrons, the magnitude of screening constant. o also decreases. The decrease in the niagnitudeiof CJ ,increases the
ma1;mitudc of eff~-0.Utlearcharge. The increased effective' nuclear charge p_u!!s the .:.!_cctron cloud ofjhe cation i~~
1~ ~1 nearer to the nucleus and thus makes the cation smaller in size than its parent .neutral atom.. -, .: -
------- - . -- . '
Thus alkali metal cations arc smaller in size tham the alkali metal atoms as shown below:
Elements Electronic confignratidt No. of electrons in the inner shel No. of inner shell! Order of shielding effect
Li 2, I
'
2 I
Na 2, 8, I
-,
-:10 2 I •<'
K 2, 8, 8, I 18 3 Increasing
Rb 2, 8, 18, 8, I 36 4 ,, J.
Cs ,__ 2. 8. 18. 18. 8 I 54 5,. - ...~ ( . .~,
Fig: Shellcllng e-r-rec:::1:
;:e'
-__ / .
/ ~·
·:. , ". . 'L.umbe" Theatomic number intervalsat which the elements with similar properties reappear are 2, 8, 8, 18. 18.4
32 and 32, that is we have passed 2. 8, 8. 18, 18. 32 and 32 elements before_ we come across an elemen_t_ with s.i~nilar
.,pr~oe-rti. The number 2. 8. 8. 18. 18. 32 and 32 are called magic number. · ·. · · · · ·
;-. S . elding or Screening effect: The decrease in the attractive force exerted by the nucleus on the valence-shell electron.
. whic , is obviously due to the presence of the electrons lying between the nucleus and valence-shell electrons is nilled
. shielding effect or screening effect.
-- 4==- - -=-,.-_....-·- ~--:> --- --;:-::~--?>"- <, ....,, 1.-.n~r- ==-~t-.~~I ~•e-•=='tr·•:=-r1:~~
( C
~£~~). -.,.~.. ~---F.'.oeF=··-~•'s'IC••·-~-----------~~
~) . ·<--'';>
&--=:::- £2:; c~__...-" ~""."-Attr·ac1:ion------~=~::.:-.-::~ shell
----~~ 1_.,.1:::1ei_.1.s -=31ec;:tron
,/
•
Factors affecting the magnitude(~.~> of shielding effect
ft~ing are the important factors on which the magnitude of shielding effect depends.
y. Number of inner shell electrons or inner shells: Greater is the number of inner shell electrons or inner shells.
greater is the magnitude of shielding effect caused by the inner electrons on the valence shell electron. Thus as we move
down a group. the number of inner shell electrons or inner shells increases and hence l<!I~ ~'IJ the shielding effect also
increases. For example, in the elements of group IA. · ·-· · ·_,.. ;,,,-;;:,·; ·
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6. t---- ----··--.·-
II'•,
(b) In a period: In gP111g from left 10 right in a period atomic size or radius decreases with the increase of atomic
number. This fact can be illustrated 1~ ~01 ~<1111by considering the atomic size of the clements of 2nd period as shown
below:
2.352. 16
CsRbAlkali metals Li Na K
Atomic size or Atomic radii: The atomic size is an arbitrary 1~J concept and is influenced by the nature of
neighbouring atoms. However. it is usually considered as the distance 'from the center of the nucleus to the point where the
electron density is effectively zero.
Variation of atomic size in a group and a period
(a) In a group: In going down a group the atomic size increases with the increase of atomic number. For example, the
atomic size of the alkali metals increases on going from Li to Cs as shown below:
ie varrauon o atomic vo umc in a peno cnn e seen 111 case o r pcno s e ernents,
Elements Na Me: Al Si p s Cl Ar
Atomic no. 11 12 13 14 15 16 17 18
Atomic volume (c.c.) 23. 7 14 10 12. I l7 15. 5 18. 7 24. .t
~eriod: In going from lefto.right in a period. atomic volume ·~ecreases at first for spme-elefuents, becomes
minimum in the middle and then increases. The variation of atomic volume in going from left tp "right in a period is
influenced by the following two factors: ·
I. Nuclear charge, . l
11. Numberof valence electrons
Tl f I . d b f J d d'
number of shells. The larger the number of shells. the biucer is the volume of atom. For example. the elements of group IA.
Elements Atomic no. Atomic volumes (c.c.)
Li 3 13. I
Na l 1 23. 7
I( 19 45. 3.
Rb 37 55.9
Cs 55 70
Var~ of atomic volume in a group and period ·. .. ·
~In a group: Atomic volume increases more or less regularly in going down a group is due to the increase in the
Atomic/Atomic volume 0~
Dcnsitv
weight .
In other words atomic volume is the volume in c.c, Occupied by 6. 023x I 02·' atoms of an element.
state and hence is commonly called gram atomic volume. It is obtained by dividing the atomic weight of the clement by i_ts ,
density.
if
·'
~ Question: Why atomic volume in a period decreases at first, becomes minimum in the middl~ and then Increases?
Answer: The atomic volume of the elements in a period decreases at first. becomes minimum in the middle and then
increases for the following two causes: ' ·· ; .,. 'f,·1 '
(a) Nuclear charge: We know that the nuclear charge that is atomic number increases by one, as we move from left to
right in a period. The increased nuclear charge attracts each electron more strongly towards the nucleus, resulting in
decreases in the volume of the atom.
(b) Number of valence electrons: Towards the close of a period. due to an increase in the number of valence electrons
the volume of the atom increases. so that it may accommodate all the electrons.
These two factors, one causing an increase and other causing a decrease. combine two results that in a period atomic
volume decreases at fir~t fix some clements. becomes minimum in the middle and then increase,
Halogens Atomic radius (A) Halide ions Radius of halide ion (A)
F (9) 0. 72 F - (I 0) I. 36 . (
Cl ( 17) 0.99 Cl. (18) I. 81 :..
Br (35) I. 14 Br - (36) I. 95
I(53) I. 33 I.. (54) 2. 16
/1
~ volume: Atomic volume is defined as the volume in c.c. Occupied by one gram atom of the clement in the solid
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7. Q-vf)-(b) In a period: On moving from left .to right in a period. the size of the atoms decreases and the effective nuclear
charge increases. Both these factors favour an increase in the force of attraction exerted r~ ~J by the nucleus on the
Electron afflnitv KJ/mole) -60 -53 -49 -47
Variation of Electron affinity in a group and a period .
#'. >(a) In a group_: Electron affinity goes on decreasing as we 111ov1.·_!·ro111 top to botto1~1 in a group. This decrease is due to.
the size of the atom mcrease. For exam le. the values of lst electron atf1111t ot crou IA 1s u1ven below:
Elements Li Na K Rb
atom anion _.#
Electron affinity defined above is called first electron affinity and it is represented as ~E1• The above process. which
depicts 1<!~ <Jlf!J the addition of an electron to a gaseous atom. is an exothermic process. ·
~Second electron affinity of an element is defined as the amount of energy required to add one more electron to its mono-
negative anion. M·1~, to convert it into di-negative anion. M2·1~1
M"c•> + e·w + Energy supplied (+E2) -+ M2·.~,
Thus the process of adding a second electron to M-1"1 anion is an endothermic process.
Example: The formation of 02•1µ1 from 01~1 can be shown by the following equations:
o(R) + e'(jt) -+ o·IR) + E1 (=-141 K.J/mole) Exothermic process
o·,.) +e·IR) + E2 (=+921 KJ/mole) -+ 02·,~, Endothermic process
( ;;asronslsolatcd •••rm"
205"'"--Electron affinity: Electron affinity of an element is defined as the amount of energy released in adding an extra electron
from outside lo an isolated neutral gaseous atom in its lowest energy state to convert it i1~10 a gaseous anion. Thus electron ·
affinity of an atom. M1µ1 can be defined by the following process:
M 12, + e <a>-+ M·12> + Energy released (-E1)
I086.4800.6Ionization otential KJ/mole 520.3 899. 5
cBElements Li Be
Q..01-.-ib) In a period: In general, as we move from left lo right in a period. the ionization potential of the elements increases '
due to the successive 1~~ increases in the nuclear charge and decrease in atomic size. For example, the values of
ionization eneru of the elements of 2nd eriod are iven below:
Ionization otential 'KJ/mole 520. 3 495. 8 .tJ8. 9 403. 4
Elements Li Na K Rb
~Variation of Ionization energy in a group and a period ·
. .z.0l~ ......;(a) In a group: When we proceed from top to bottom in a group the ionization energy-values of the elements go on
decreasin '· For exam le. the values of ionization ener 1 of the elements of grou TA are given below: .·,
Thus we can see that, 2nd ionization energy is greater than lst ionization energy and Jrd ionization energy is greater than
. , those of I st and 2nd ionization energy:
4578
Ionization Ener v 11 12
e: Ionization energy or Ionization potential: Ionization energy of an element is detined as the amount of energy required
to remove the most loosely hound electron from isolated neutral gaseous atom in its lowest energy state that is ground state to
convert it into a unipositivc gaseous cation. Thus the ionization energy of an atom. M,~1 .can be defined by a process
represented by the following equation:
Mru> - .e· +Epergy supplied (+11)-+ M'ru>
I
. L_ __ ' .
lso atcd gaseous ~~J ( rascous < 'ntion
Ionization energy defi ncd as above is cal led tirst ionization energy and it is represented as 11 or -1 11 _ • ~ 1
0 Second ionization energy ( 12) is the energy required to remove one more electron from the gaseous cation. M ·1~1 to get
the doubly positive charged gaseous cation. M2',ei·
M'12, + 12-+ M2•1~, + e· . ·
? 82. ~ Similarly. third ionization potential (It) is the energy required to remove still one more electron from M2' 1~, cation to get
M t •
1~1
cation. ··
M2'1i> + 1.,-+ M·1'<R> + e'
For exam rle, the 11• I~ and I.1 of Nitroucn atom are as follows:
0.90I. 23 0. 77Atomic radius 0. 730. 750.82 0. 71o. 72
Li NBElements c NeF0Be
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8. 7
Polarisation: When a cation (C.) of an ionic molecule (CA) approaches 1f~i ~·3l111 closely to an anion ;-. il
withdraws the electrons of the anion Inwards itself' and the electron cloud or the anion gets distorted I~ ~'3IDI from its
Electrone ativit I I. S 2 2. S 3 3. 5 4
Elements Li Be B C N 0 F
(b) In a period: In going from left to right in a period, the electronegativity values increase. This increase can be
explained by the following facts.
I. On moving from left to tight in a period. there is a decrease in the size of the at0111s. Smaller atoms have greater
tendency to attract the electrons towards themselves that is smaller atoms have higher clectronegativity values.
II. On moving from left to right in a period there is an increase of ionization _energy and electron affinity of the
elements. The atoms of the elements. which have higher values of ionization energies and electron affinities. also have higher
electronegativities.
For exam le, the values of electroneuativit 1 of 2nd eriod arc uiven below:
0. 700.80Electrone ativit I 0. 9 0. 89 0. 86
Cs FrElements Li Na K Rb
Variation of electroncgativity in a group and period
(a) In a'group: In going down in a group. the electroncgativity values decrease. This decrease can also be explained by
the following facts.
I. As we move down a group, there is an increase in the size of the atoms. With the increase in size of the atoms. their
electronegativity values decrease. . ; : -· .,~..,.,,
II. Ionisation energy and electron affinity on which electronegativity depends decrease as the group is descended. With
the decrease of these quantities the electronegativity values also decrease. ,/
For exam le. the values of electrone ativit of rrou IA are ziven below:
0.86 0. 700.800.89
K FrRb Cs
IA are riven below:
Electroneganvlry; When two different atoms in a molecule are bonded together by a covalent bond, the electron ,rair
forming the covalent bond is not shared equally hy both the atoms. The electron pair lies nearer to one atom than the other.
The relative tendency of a bonded atom in a molecule to attract the shared electron pair towards itself is termed as its
electronegativity, -
Example: The formation of o~·,~, from 01~1 can he shown by the following equations:
o(RI + e"(al -4 o-(R) + E, (=-141 KJ/mole) Exothermic process
0-1a) +t'1~1 .+ E2 (=+921 K.J/mole) ~ 02·12, Endothermic process.;
Mom ;111io11
Electron affinity defined above is called lirst electron affinity. The above process. which depicts the addition of an
electron to a gaseous atom. is an exothermic process since in the addition of an electron energy is released.
Second electron affinity of an element is defined as the amount of energy required to add one more electron to "its mono-
negative anion, M-1~1 to convert it into di-negative anion. M1·1~1
M·,21 + e·121 + Energy supplied (+E2) ~ M2-121
In the process of adding one more electron to M'1µ1 anion against the electrostatic repulsion 1~"11 between the extra
electron being added to M-1~1 and the negative charge on rvr1~1 • n. energy instead of being released. is supplied to M·,"' ion
to convert it into M~.1~1 anion. Thus the procc~-; of ad · ~ second electron to M·1~1 anion is an endothermic process.
Question· E'q'.!lain the term, "Second electron affinity always enoothennic P.rncess."
Electron affinity: Electron affinity of an element is deli ncd as the amount of energy released in adding an extra electron
from outside to an isolated neutral gaseous atom in its lowest energy state to convert it into a gaseous anion. Thus electron
· affinity of an atom. M1µ1 can he de tined hy the following process:
M lal + e ,a,~ M-,2, + Energy released (-Ei)
-0+20-122 -141-23-60Electron affinity (Kd/mole)
Be()NB cLiElements
electron. Consequently 1~1 the atom has a greater tendency to attract an extra electron Irom outside towards itself and
hence its electron affinity increases from left to right.
For example. the values of I st electron affinity of 2nd period is given below:
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9. 11. Representative or normal elements: In the atoms of these elements the outermost shell is partly filled while the
inner shells arc completely filled. The electronic configuration of the outermost shell of the atoms of these elements varies
I , '
from ns to nsp . The elements of groups IA. llA. I IlA, IV A. YA. VIA and YI IA belong to this type.
The chemical properties of these elements depend on the valence shell electron and these are both metals and non-
metals. Thus, the alkali metals are s-block elements. The valence electron of all the elements from group IJIA to VllA
vertically occupies p-orbital, Hence these elements are called p-block elements. They are generally form colorless compound.
I. Noble gases: In the atoms of these elements the outermost shell which has nsp" configuration (He which has ls!
configuration is an exception) is completely filled. ns!p6 configuration is stable and hence -these elements generally do not
enter into ordinary chemical reactions. These elements are present in zero group of the periodic table.
< Classificat_ion of elements according to their electronic configuratio~1 .
On the basis of electronic configuration the clements may be classi fied into tour types:
I. Noble gases III. Transition elements
II. Representative or normal elements IV. Inner-transition elements
IV. f-Biock elements (Irmer-transirlon.elemeurs). In these elements .the last electron e'nters into the f-ortlital of ante-
penultimate 1c-fll' ~ ~1 shell. Hence these elements arc /call'ed, f-block elements. The valence shell electronic
configuration of the atoms of these clements is (n-2)f-14( n-1 )d11•1 ns2• these elements are located in group 1118 and have been
given a separate place at the bottom of the periodic table. f-block elements are also called inner-transition elements.
Ill. d-Block clements (Transition elements): In these elements the last electron enters intp the d-orbital of the
penultimate 1 >1<1c•111fGJ!l ~1 shell. Hence these elements are called d-block elements; In the atoms of these elements thens-
orbital is completely filled and the valence shell configuration of these clements varies from (n-1 )d1ns2 to (n-I )dlOns!. The
clements of groups lB. IlB. IIIB. IVB, VB. VfB, VflB; and Vijl belong to' this block. The elements of this block are .
collectively known as transition elements. since thei r properties are intermediate between those of s and p-block elements.
These elements are located in the middle of the periodic table and consist of metals only.
~':
II. p-Block elements: In the atoms of these elements the last electron enters into the p-orbital' of the outermost shell.
Hence these elements are called p-block elements. In the atoms of these elements thens-orbital is completely filled and the
valence shell configuration of these clements varies from ns~p1 to ns2p''. The clements of groups lllA. IVA, VA. VIA. VllA
and zero (Ne to Rn) belong to th'is block. The elements of this block. like s-orbital elements. are also called normal elements
or representative elements. . . ,
p-block elementsare localed at the extreme right of the periodic table and consist of metals. non-metals. metalloid- and
inert gases. The properties of p-block elements arc determined by the number of electrons present in np orbitals.
I. s-Block elements: In the atoms of these elements the last electron enters into the s-orbital of the outermost shell.
Hence these elements ~re called s-block elements. The valence shell electronic configuration of these elements varies from
ns' to ns!. The clements of group IA. group llA and He belong to this group. The elements of this block arc collectively
known as normal elements or representative clements.
s-block elements are located at the extreme left of the periodic table and consist of active metals. The properties of s-
block elements depend on the number of electrons present inns orbital.
•...'
Classification of elements according to outermost shell configuration
On the basis of outermost shell configuration. the elements may be classi lied into four types:
I. s-Block elements Ill. d-Block clements
II. p-Block clements IV. f-Block clements
Fig: Polarisation.
Polarization power or ability: The ability or a cation to polarize a nearby anion is called its polarizing power or
polarizing ability.
Fajan's rule: Polarisation depends on the following statements:
I. The charge either on cation or anion should be large. r--..2~ The size of the cation should be xmal I. '-..(..)
3. The size of the anion should be large.
/
.
'
'
'
. .
8
symmetrical shape. Thus the electron cloud or anion no longer remains symmetrical but 1s elongated PF~ <Jinl towards
the cation. This phenomenon is called distortion or deformation or polarisation.
..
4. The presence of electron in d and f orbital of the clement.
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10. lliii!Pounds of alkali metals Na20, Na2S, NaX J
L;o
f<Ov
A' Valency: Like alkali metals valency of hydrogen is one. because it has only one electron on· its outermost shell and
forms unipositive cation.
' iii. Formation of M' ions: Like alkali metals hydrogen is a strong electropositive clement. It also has a tendency trr '
lose its only one electron to form the unipositive ion. 7
H ' H' ....
. Li ·c Li' J ~~ducing agent: Just like alkali metals hydrogen is a very good reducing agent.
/Forms stable compound: Like alkali metals hydrogen forms stable compound with oxygen. halogens and other
elements. pc;~1'0..,, o-J 0,UMI) • ,
Elements Atomic no. Electronic conflguration
..
H I Is'
Li 3 ls22s1
Na 11 ls22s22p63s1
K 19 I s22s22p63s23p64s 1
•~ Question: Explain the position ofhydrogen ·n · geriodic table
Hydrogen rescmoles both the alkali metals and halogens. Its position in periodic table 1s therefore anomalous. The·
position of hydrogen in the periodic table is explained below:
RJM(sons for placing in group IA -. · ·
...,(. Electronic configuration: Like alkali metals hydrogen also has only one electron in its outermost shell as shown
below· -rs'
Question; Hat is lantnanides contr8ctio '!
Lanthanides contraction (~~): A look at the values of the radii of tri-positive lant,ranides cations (M'J) express
that. these values decrease as we move from La'' to Lu'·' in lanthanide series.This steady 1'tif) decrease in the values of the
radii ofM'' cations irt the lanthanide series is called lanthanides contraction.
,71 Q estion: ust1fytlie position of actinides in the periodic table .
...,,.- I Answer: All the 15 actinides IAc(89) - Lr( I 03)] have atomic weight between those of Radium(88) and Dubnium( I 04l.
and therefore must be placed between these two clements.
Radium has exactly the same outermost electronic configuration as Strontium(38)and Barium(56) and resembles them
very closely. Similarly. Dubnium is similar to Zirconium(~) and Hafnium(72). Therefore Radium must be placed below
Barium and Dubnium below Hafnium.
since all thefifteen actinides resemble one another in many ways. so it becomes necessary to accommodate all the
fifteen actinides at one place. This has been done by placing the first element Actinium below Lanthanum(57) and remaining
14 elements have been placed separately in the lower part of the periodic table.
Question: ustify the position of lanthanides in the period c talile,
~x<G,, J} /1' Answer: All the 15 lanthanides [La(57J - LLi(71 )] have atomic weight between those of Barium(56) and Hafnium(72)
/and therefore must be placed between these two elements. ·
Barium has exactly the same outermost electronic configuration as Calcium(20) and Strontium(38) and resembles them
very closely. Similarly, Hafnium is similar to Titanium(1,'!) and Zirconium('fe). Therefore Barium must be placed below
Strontium and Hafnium below Zirconium. ·
Since all the fi fteen lanthanides resemble one another in much way. so it becomes necessary to accommodate all the
fifteen lanthanidcs at one place. This has been done by placing the first element Lanthanum below Yttrium(39) and
remaining 14 elements have been placed separately in the lower part of the periodic table.
IV. Inner-transition elements: In the atoms of these clements three outermost shells arc partly filled while the
remaining inner shells are completely filled. l.anthanidcs and actinides belong to this type. These clements lie in 6th and 7th
periods respectively. The orbital in which the electron is added on increasing the atomic number is t-orbital. The series of 14
elements in which 4f orbital is being build up follow lanthanum and are called lanthanides. The series of elements in which 5f
orbital is being build up follow actinium and are called actinides. The inner transition elements are all metals and show
variable oxidation state. Their compounds are highly coloured.
Ill. Transition elements: In the atoms of these clements two outermost shells arc partly filled while the remaining
inner shells are completely filled. These clements lie in 4th. 5th. 6th and 7th periods of the long form of periodic table, I hcsc
clements arc generally heavy metals of sub-group B and contain two incomplete energy level. because of the building up of
the inner d-electron. The chemical properties of these elements depend on the electrons from the two outermost shells. I hcsc
elements general Iy form coInured compound.
[,
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11. ·----........:.._
Question: Explain. why ionization potential of Beryllium is higher than that of Boron?
In ease of N! it is more difficult to remove an electron trom the half lilied Zporhiral. while in case of O. it is easier to
remove the same electron from partially filled 2p orbital. Thus the ionization potential ofN, is.higher than that of 02.
l-lcctron partially filled
itaifliikd ~p orhilai
II '' 11·
~hysical appearance(~@~): Hydrogen is also gas ~kc gaseous halogens.
~uestion: ExQlain. w 1y 10nization potentral of Nz is higher than that of O,'!
. ~ Answer: According.to Hunds rule. fialf ~Ilea or completely filled ~rbitals ~re comparatively 1~01~1~01_4~1c<11 mo~c st~hlc
and hence more energy rs needed to remove an electron from such orbitals. This means that. the 10111zat1on potential of an
U p-rri atom having half tilled or completely filled orbitals in its electronic configuration is relatively higher than that expected
. normally from its position in the periodic table. The equation. which represents ionization potential of N2 and 01• can be . ~1
written as:
F
viii~ectronegative character: Like halides 'when hydrides an; electro.lyzed, hydro.genis liberated at the anode.
W Formation of negative ion: Like halogens hydrogen also gain one electron to form negative ion.
11. 92
BrElements H · ' Cl.
1........__
i'a.J·crrru'a. pnri~fct:Ji'rfYl'. Hr>.., ~ eJ~
vlf."° Non-metal: Like halogens hydrogen is also non-metal.
iii. Atomic stat~ Hydrogen is diatomic molecule like Halogens. . ·
~ombination with QQn;fifJ'~ Halogens form various types of non-metallic halides such as CCl4, SiCl4• GeCl4.
Similarly hydrogen al~o forms CH4, SiH4, GeH4.
v, Exchange of hydrogen and halogenatoms: Hydrogen atom can be exchanged with halogen atoms in compounds
and vic~ersa. ·
''.Y." Formation of hydride: Halogens react /•ith metal to form halides such as NaCl. KBr etc. Similarly hydrogen forms -----
hydrides such as NaH. CaH, etc. · '
vii Ionisation otential: Ionisation otential of h dro en is of the same order as that or halo tens,
Reasons for placing in group VllA
i. Electronic configuration: Like every halogen atoms hydrogen has short of om: electron in its outermost shl"ll as
com arcd to the next inert ras Helium.
. I Compounds of hydrogen I H20, ll!S, HX I r-
,-. uoi» EJ-ffl11} 0J--J e Yl'i- ,l 'tr
. / · ~. Atomic state: Hydrogen is diatomic that is hydrogen exists as H2 (H - H). similarly ellteli metl.tffl such. asy~
I
~1rtd&;:tif(l.i Li);>
~opositive Qberlie:tl't: When hydrogen halides or oxide arc electrolyzed. hydrogen is liberated at the cathode.·..
In the ~me way alkali metals also do so. f-1.-,.... HD
· ~· Affinity for non-metals: Both hydrogen and alkali metals have a strong affinity for non-metals and little affinity for
metals.
... '
Answer: Generally filled and half tilled orbitals are more stable. The electronic configuration of Beryllium and Boron is
given below: ·
Be (4) ~ ls22s2
B (5) ~ ls22s22p1 .
In case of Boron, less energy will be required to remove 2p1 electron. But the electronic configuration of Beryllium is
more stable and removal of one electron from 2s! will cause breakdown of this stable form and for removing this electron
more energy will be required. Thus we can say that. ionization potential or Beryllium is higher than that of Boron.
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12. Again in a group the reactivity of halogens decreases if we move from fluorine to iodine.
Example: Fluorine' decomposes water vigorously even in dark. chlorine decomposes H~O by sunlight, bromine also
decomposes by sunlight but very slowly while iodine does not decompose water at all. · _ ~- .J "-t
F2 + H20 ";jJ>F2 + H2;(Li~i.l-·~ Cl2 + H20-+ HOCI + HCI (L'ghf dfl,~l.OX'ilV
Br2 + H20 -!!+ HO Br + HBr ; <YrrJ•w:I) 12 + H20 -+ No reaction.
Thus. we can say ihat fluorine is the most reactive element of the modem periodic table ..
~ Question: ExP.laiil the te1111. "Flugi~n~nm~t be obtained by oxidation of Fluo1iae:. . . _ _ .
A,.i~rer: Smee fluorine has the nsnp configuration 111 its outermost shcll.J.t uin e·1>il)' obt21e ntt elesirea 10_[1jlhll ll!i
~ ~ 4*Ut I Ilic lBs:# high electronegativity, it can easily obtain one electron tor:·ormnfluoride.ln which electronic
configuration 1s the same to its next element Neon.
1
F (9) ~ ~ rtP~~h:lftl- -r,.J.ylf ,·./-!) · oc.-fa}.I}_. cu.-J
Fc9> + e -+ F iH1i ~ ls"2s"2p ~ Ne (IO)
_5,'.,.,t.D.._:kHhe configuration~is more stal~le o~~ ia~ no tendency t? loss an el~ctron from fluoride (such as Nan
~the most strong oxidizing agent a~IH J· ·, · ti.·.· · • e- · · t. ~ ~ ou•d-a.J, OrL._ R
Nature of oxides Strongly Basic Amphoteric Feebly Acidic More Most
basic acidic acidic acidic
Oxides of elements Na 0 M 0 . Al 01 SiO P O. SO, CI,O
From the above discussion. we can say that ionization potential is a periodic function.
~ Question: Explain tfiat... Fluorin~'is ilie most reactive element." . < '' 'lua hJA.t-
Fluorine is the most reactive element: Fluorine is the most reactive element because of its .mgtlelectronegativity and
very low dissociation energy. Generally in a period, the acidity increases as we move from left to right. As the halogens have
the highest elcctronegativity they form strong acids. J_
Exam le: acidit increases ifwe niove from left to riuht in~-<>criod.
otential will be increased. For exam le, the elements of 2nd eriod.
495.8 418.9 403.4
Nn K RbElements cA1 Li
( iascous Ca1ion
otential will be decreased. For exam le, the elementsof rou l IA.
Isolated gaseous cation
If we move from ro· to bottom in a zrou
~===1;}>~~1Question:Ex~lai·n that. ..Ionization potential is a pcnoaic function ..
.._ Ionization potential: lomsat1on energy of an clement is defined as the amount of energy required to remove the most
loosely hound electron from isolated neutral gaseous atom in its lowest energy state that is ground state to convert it into a
unipositive gaseous cation. Thus the ionization energy of an atom. M1~1 can be defined by a process represented by the
following equation:
M1~1 e + Energy supplied (+Ii)~ M'1~1
lsotones: The atoms of different elements. which haw the same number of neutrons but different atomic number and
mass number. are called isotones.
E I 14 C 15 N d 1110 . • I I - .xamp e: 1, , 7 an x are isotones of eac i otner. LI . • 1 .. ~ C!-~aJ._±L ~f.> uJNc.n I 5 peJUOC1--' '
. "Th»: P'1'ff':n.Ji·~ c:f "" di
~--·r>- Question: hat ·s Rcriodic function'! _}_a_/;-1-::lo '1i(1-f J rUf1A:I- :Ir> ./.Qfl a'"(/ -:/tJp.:lo boffo'rn; laoH-orn-lo for oVl- ca1J.a.d f'.eJL1'
1
Periodic function: If we go left to richt or right to left in a period and if we go top to bottom or bouom to lop in a group
some properties of the clements change periodically and this phenomenon is called periodic function.
Isobars: The atoms of different clements. which have same mass number but different atomic number. are railed
isobars.
40 40 40 . .
Example: ix Ar, 1,1
Kand :o Ca arc isobars of each other.
..
number but different mass number arc
. '
Question: efine Isotopes lsob:ars ano lsotonei;:.
Isotopes: Different atoms of the same element. which have the same atomic
called isotopes of that element.
E I "c 1.1c d14C 1 · rr hxamp e: 11 , 11 an " . are l tree isotopes o ar on.
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13. 12
I
l!i
I
j,
?:
~l
~~"-·'"':;~.~..:.-::-
''; l
: l
I
·1
llsefulne~s orthe eriodfc table
Ttie following important applications of the periodic table may be mentioned 1~ <TiTI ~ ~1:
I. Classiflcatlon of the elements: The classification of elements of~mil:ir properties into groups simplified their
study. For example, sodium. a member of alkali metals group reacts With water vigorously 1~1 giving hydrogen gas and
forming sodium hydroxide. which isa strong base. The other alkali metals also react with water ina similar manner 1"f'tff'51.
II. Prediction (~'fl.ft) of u1tdlscovrl-ed elcAfei,1ts: At presen: all the elements frotrt atomic number I . lo I 09 lmw
' . ,.i~ ,. ~ - '-~ . .~: - . -
been discovered and their properties are mofe or lessknown. t3tit a very remarkable 1c=rQ;<frn1 use or the. periodic table was
~ . . . . '
made by Mendeleef in predictinga number of undiscovered elements, which were shown byd number of gaps in the periodic
table. Mendeleef's tableicont~ined only 65 ele111ents with a lar e number o v c n Jaccs. Jy{endeleef predicted the.mstcncc
~-<i..~n'e~ . C(!rr~sQ~n?!.n&,)~1'.('11 .. tn. tl!e..,ga[>s. These .,;l;ments ha~e~red an~
scand1um.£.t1Jl!urn 1e.rma.nwrn...._l~ue.._t.u1tU_~~P..Qlonrnm. · · •_,.,,-_"._ ....-~ . -- ,. I a=-~
III. Correction of atomic weight: Atomic weights of Some of the elements at the time of Mendeleef gave a wrong
position ~ ..e6MtM::eie.1.ttuti i11 the periodic table. The properties ot these clements required their placement somewhere
else. For instance t~<fl, the element indium was placed in a vacant place in the periodic table between Cd ( 112.4) and Sn
(I IR. 7) and indium with atomic weight of about I 14 titted very well in between Cd and Sn. , .
IV. Periodic table in industrial research: The periodic table has been foundIn be quite l~"'.!~I useful in industrial
researches. Several of the light metals and their alloys1~ 'T~~<ft used in modem mechanical equipments 1~1'f!f51. jet
engines and aircrafts were first studied in detail because of their positio~ in the periodic table.
Question:' Explain that. "Fhmrine is rhc mos! clcclrorn:gali~,.. .
Answer: Ir we go left lo righl in a period the atomic size of the elements willbe decreased and if we go bottom lo lop in
. a group the atomic size of the elements will be decreased. Because fluorineis placed in mos! right side and lop of the periodic
' table, so ii is smalles: among all the elements of the periodic table. Again we know that. smaller atoms have greater tendency
to attract the electrons towards themselves thatis smaller atoms have higher electronegativity values. So fluorine is the most
electronegative in the periodic table. ·
i .
;· rr:..;pl Question; Explain that, "C:itlm1m is the most electrepo);itjve." tJ · -. . .
Answer: {r we go left to right in a period the atomic size of the elements will bedecreased and if we go lop to bottom 111
a group the atJmic size of the elements will b~@re1ued. Because lithium is placedin most left side and lop of the periodic
table. so it is largest among all the elements of the periodic table. Again we know that. larger atoms have lower tendencylo
attract the clcktrons towards themselves that is larger atoms have lower elcctronegativity values. So lithium is the most
clcctrnpnsitivt.~ in the periodic table. -
·~·
Quesnon] Explain the term ··· he=c.t!l~Gtr-on affinity Hf fluorine is less-than ehlonnc."
Answer: t'he electronic configuration of fluorine and chlorine is given below:
F (9) ~ I s22s22p~ .
:. Cl~ -4 ls22s22p63s2.3p~
In case of fluorine Jhe oute1p10sl shell is 2 and in case of Cl is 3. In fluorine atom. the attraction force of nucleus on
corning electrons is les~ than·~h¢ force between outermost shell's electrons and coming electrons becausein case of shell 2
the electron cloud remains more' densed, So the overall electron affinity of fluorineis less than chlorine.
r
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