Download app for all tests and notes
Click here for NEET/IIT-JEE and other competitive exam books
Periodic Classification
(a) Dobereiner grouped elements in
triads e.g. Li, Na, K or Cl, Br, I etc.
(b) Newland
found that similar elements are repeated at 8th place. It happened when
elements were arranged in the increasing order of their atomic weight. It is
applicable for lighter element only.
(c) Lother Mayer plotted a graph between atomic volume of
elements against their atomic weight. He found that similar elements occupied
similar positions on the curve.
(d) Mendeleev’s Periodic Law and Table: Mendeleev arranged all
the elements in order of their increasing atomic weights. A table which has
been formed with the help of classification of elements is called periodic
table. The method of arranging similar elements in one group and separating
them from dissimilar element is called classification of elements. Mendeleev
prepared the table on the basis of his periodic law called Mendeleev’s periodic
law.
Mendeleev’s periodic law : The physical
and chemical properties of elements are the periodic function of their
atomic weights.
Mendeleev’s periodic table
consists of seven horizontal rows known as periods and nine vertical
columns known as groups.
Period
: Out
of the seven periods in all, the first three periods are known as short
periods while the fourth, fifth and sixth periods are called long
periods.
Periods
Period Total No. of Starts with Ends
with Elements Remark
Elements Elements
1. 2 Hydrogen(1) Helium(2)
Very short
period
2. 8 Lithium(3) Neon(10) Short period
3. 8 Sodium(11) Argon(18) Short period
4. 18 Potassium(19) Krypton(36) Long period
5. 18
Rubidium(37) Xenon(54)
Long period
6. 32
Cesium(55) Radon(86) Very long
period
7. 26
Francium(87) (Not named
yet)(112) Incomplete period
Groups
(i) There are nine groups in all including 8th
group of transition elements and zero group of inert gases.
(ii) All the group from I to VII (except zero and VIII) are divided into
sub-groups A and B.
(iii) The group number of an element represent its valency.
(iv) The elements of same sub-group resemble one another more closely and
generally differ to some extent from the elements of the other subgroups.
Defects
in Mendeleev’s Periodic Table
(i) Position of Hydrogen : Hydrogen resembles both the alkali
metal (I group) and the halogen (VII). Hence its position in periodic table is
undecided.
(ii) Position of Isotopes :According to Mendeleev’s periodic
law isotope of an element should occupy different position in the periodic
table, but this is not so.
(iii) Position of VIII group elements : Nine elements in the VIII
group do not fit into the system.
(iv) Position of Lanthanides and Actinides :
Their position can not be justified according to the periodic law and
cannot be arranged in the order of their increasing atomic weight.
(v) Dissimilar elements placed in the same group : Alkali metals
(Li, Na, K, etc.) are placed with coinage metals (Cu, Ag, Au).
(vi) Similar elements are placed apart : Chemically similar elements
like Cu, Hg and Ag and Ti, Au and Pt have been placed in different groups.
(vii) Anomalous pair of elements : Some elements of higher
atomic weight have been placed before elements of lower atomic weight. For
example, argon (At. wt. = 39.9) has been placed before potassium (At.wt =
39.1); cobalt (At. wt. = 58.94) is placed before nickel (At. wt. = 58.69) ;
tellurium (127.5) has been placed before iodine (126.9).
(e) Modern periodic
table : It is also known as long form or Bohr’s table as it is based on
Bohr’s scheme of the arrangement of elements into four types according to their
electronic configuration. Recent work has established that the fundamental
property of an atom is atomic number and not the atomic weight. Therefore,
atomic number is taken as the basis of the classification of elements.
The modern periodic law
may be stated as “The properties of elements are periodic function of their
atomic number.”
IUPAC nomenclature for the
superheavy elements
Atomic Name Symbol Atomic Name Symbol
Number Number
110 un-un-nilium Uun
101 un-nil-unium Unu 111 un-un-unium Uuu
102 un-nil-bium Unb 112 un-un-bium Uub
103 un-nil-trium Unt 113 un-un-trium Uut
104 un-nil-quadium Unq 114 un-un-quadium Uuq
105 un-nil-pentium Unp 115 un-un-pentium Uup
106 un-nil-hexium Unh 116 un-un-hexium Uuh
107 un-nil-septium Uns 117 un-un-septium Uus
108 un-nil-octium Uno 118 un-un-octium Uuo
109 un-nil-ennium Une 119 un-un-ennium Uue
120 un-bi-nilium Ubn
130 un-tri-nilium Utn
140 un-quad-nilium Uqn
150 un-pent-nilium Upn
Note : Hyphens have been
put in the name for clarity. They should be omitted.
Classification
of Elements
On the basis of electronic configuration the elements can
be classified into the following four types :
(a) s-block elements : These elements contain 1 or 2
electrons in s-subshell of outermost shell. Elements of 1 and 2 group belong to this class. These elements
enter into chemical reaction by losing valency electrons so as to acquire noble
gas configuration in the outermost orbit.
ns1 (group 1) ns2 (group 2)
(alkali metals) (alkaline earth metals)
These elements generally
form electrovalent compounds and basic oxides.
(b) p–block
elements : These elements contain 1 to 6 electrons in the p–subshell of the
outermost orbit
(ns2 np1–6). The elements belonging to 13th to 18th group except He are p-block elements. In these last electron enters to the p-sub-shell. For example.
(ns2 np1–6). The elements belonging to 13th to 18th group except He are p-block elements. In these last electron enters to the p-sub-shell. For example.
13 Boron (B) Z
= 5 1s2 2s2 2p1
14 Carbon (C) Z = 6 1s2 2s2 2p2
15 Nitrogen (N) Z
= 7 1s2 2s2 sp3
16 Oxygen (O) Z
= 8 1s2 2s2 sp4
17 Fluorine (F) Z
= 9 1s2 2s2 2p5
The main characteristics of these
elements are :
(a) The non-metallic character increases along a period from 13 to 17.
(b) They form covalent compounds among themselves but electrovalent
compounds with s-block elements.
(c) Their oxides are generally acidic, few
are amphoteric also. For example Al2O3, Ga2O3.
(c) d–block
elements : These are called transition elements or ‘d’ block elements. The
elements
of group 3 to 12 belong to this class. Their general configuration can be represented as :
(n–1)d1–10 ns1–2
of group 3 to 12 belong to this class. Their general configuration can be represented as :
(n–1)d1–10 ns1–2
General characteristics of
transition (d–block) elements:
(i) They are metals, hard, malleable, ductile and possess high tensile
strength.
(ii) They are good conductors of heat and electricity.
(iii) These elements exhibit variable valency.
(iv) They generally form coloured compounds. This is due to the presence
of incomplete d–subshell.
(v) These metals, their alloys and compounds possess marked catalytic
activity.
(vi) They are generally paramagnetic, i.e., attract magnetic lines
of force.
(d) f–block
elements : They are inner transition or f-block elements. These elements
are arranged in the two row at the bottom of the periodic table. In the first
row 14 elements from atomic number
58 to 71, known as Lanthanides or rare earth elements. The second row of elements from atomic number 90 to 103, known as actinides. Their general electronic configuration can be represented as
58 to 71, known as Lanthanides or rare earth elements. The second row of elements from atomic number 90 to 103, known as actinides. Their general electronic configuration can be represented as
(n – 2) f 1–14 (n–1)d 0–1 ns2
They show most of the properties
similar to each other since outermost and penultimate orbits are similar.
Their properties are similar to ‘d’ block elements.
Periodic Trends in properties
A. Atomic Size
(or atomic radius)
Atomic radius is the size of the atom of
an element. Atomic radius is defined as “the distance from the centre of the
nucleus upto the centre of outermost electron.” It is measured in Angstrom unit
(Å). It is not possible to measure exact atomic radius as an atom is unstable
and it cannot be isolated to get its radius. Moreover, the exact position of
the outermost electron is uncertain. The values for radii are obtained
from x-ray measurements. Following points are to be noted in this reference :
(a) The
size of an atom or ion decreases in a horizontal period as we move from left to
right.
(b) The
atomic radius increases in a group with the rise in atomic number.
(c) A
positive ion (cation) is smaller than the corresponding atom : A positive
ion or cation is formed by the loss of one or more electrons from an atom and
the number of protons remains the same in the nucleus. Thus the ratio of the
positive charge in the nucleus to the number of electrons i.e.,
effective nuclear charge increases. Hence the force of attraction of nucleus to
the outer electrons increases thus decreasing the size of cation. In case of
alkali metals, the removal of an electron removes the entire outermost shell.
(d) A
negative ion (anion) is bigger than the corresponding atom : In the
formation of negative ion (anion) one or more electrons(s) are added to the
atom. Thus results in the expansion of the size of the nuclear charge, which in
turn decreases the force of attraction and increases the size of an anion or
the pull exercised by the nucleus on the electron become less i.e., they
move a little farther resulting in an increase in the ionic size.
Note
:
Size of Iso-electronic ions
: These are
such cations or anions which carry the same number of electrons. The size of
such ions depends upon the effective nuclear charge. Greater the nucelar charge
of an ion, greater will be the force of attraction for same number of
electrons. As a result, the size of the ion decreases. For example :
N3–, O2–, F–, Na+, Mg2+ and Al3+ are isoelectronic ions, among these N3– is largest (1.71 A) and Al3+ is smallest (0.50 A).
B. Ionization Enthalpy
It
is the amount of energy required to remove most loosely held electron from the
ground state of an Avogadro number of the isolated atoms, ions or molecules in
the gaseous state. The ion formed by loss of first electron may lose further
electrons and thus we may have successive ionization energies for removal of
2nd, 3rd and 4th electrons in the gaseous state.
Ionization
is always an endothermic process and ionization energies are therefore always assigned
positive values.
Factors influencing
Ionization energy
(1) Successive Ionization - Generally ionization energy increases
for successive ionizations.
(2) Atomic size - Ionization energy
decreases as the size of atom increases.
(3) Value of Z - Higher the value of Z,
higher is the I.E.
(4) Distance
of electron from the nucleus - Smaller the distance of the electron from
the nucleus larger is the ionization energy
(5) Sheilding
effect - Higher is the sheilding of the electron to be removed lower is the
I.E. Sheilding effect of the electrons of different orbitals follows the order s
> p > d > f.
(6) Penetration
effect - Higher the penetration power of the electron to be removed higher
is the I.E. The penetration power of electrons of various orbitals follow the
order s > p > d > f.
(7) Nature
of shell - Ionization energy increases if the electron to be ionized from
the species belongs to a half filled shell or a completely filled shell. The
relative stability of the these configurations follows the order d5 < p3 < d10 <p6.
(8) Changes
in the quantum shell - During the successive ionization, the electron to be
ionized belongs to the lower quantum shell the I.E. therefore increases many
folds. It is the combined effect of (a) effective nuclear charge (b) stability
of completely filled shells (c) closer proximity of the lower shell to the
positively charged nucleus.
Note : I.E1 : Li < B < Be < C < O < N <
F < Ar
I.E1 : Na < Al < Mg < Si < S < P
< Cl < Ar
I.E2 : O > F > N > C.
C. Electron affinity and Electron gain Enthalpy
(DHeg)
It
is defined as the energy released when electron is added to the valence shell
of the one mole of isolated gaseous atoms or ions and the enthalpy change
accompanying the process is defined as the electron gain enthalpy (DegH). Electron
gain enthalpy provides a measure of the ease with which an atom adds an
electron to form anion, X–(g) + e– ® X–2(g) . It may
be considered to be same as I.E. of corresponding anion. The first E.A. of
active non metals is negative. But the addition of a second electron to an
already formed anion makes the reaction , (X– + e– ® X2–)
endothermic. At the time of formation of oxide or sulphide ion, the effect of
2nd E.A. is so much that the overall E.A. for the formation of oxide and
sulphide ions is endothermic.
Variation of electron
affinity
(1) Elements which have
higher I.E. have higher E.A. also.
(2) The E.A. values of the second period elements
are lower in comparison with the values of third row elements. This is due to
increase in the interelectronic repulsions which are more for the smaller
elements because of higher electron densities.
(3) Effective nuclear sheilding by the 's'
electrons and the necessity of using higher energy orbitals to accept electrons
turn the E.A. of group IIA elements negative.
(4) Au because of very high effective nuclear
charge has higher electron affinity.
(5) The
electron affinity of the elements having d10s2 configuration are positive as electron is to
be accomodated into the higher energy p orbital.
Note : EA1 :
Cl > F > Br > I
EA1 : S
> Se > Te > Po > O
EA1 : C
> B > Li > Be
EA1 : Si
> Al > Na > Mg
EA1 : F
> O > N > Ne.
D. Electronegativity
It
may be defined as : “The tendency of an atom to attract shared electron pair
towards itself in a molecule”. The small atoms attract electrons more strongly
than larger ones hence they are more electronegative.
The
numerical value of electronegativity depends upon the ionisation potential and
electron affinity. Higher ionisation potential and higher electron affinity
both imply higher electronegativity. To measure electronegativity an arbitrary
scale was developed by Linus Pauling which is known as electronegativity scale.
On this scale flourine has maximum electronegativity of 4.0 and Li has a
value of 1.0 while inert gas have no value of electronegativity.
The
value of electronegativity show periodic variations as given below
Generally
in a group electronegativity generally decreases from top to bottom due to
increase in size of atom.
F Cl Br I At
4.0 3.0 2.8 2.5. 2.2
In a period electronegtativity increases from left to
right
Li
Be C N O F
1.0 1.5 2.5 3.0 3.5 4.0
Important
Points
1. Melting point of alkali metal halides
follow following order
M – F > M – Cl > M – Br > M – I
NaF > NaCl > NaBr > NaI
2. The order of
melting point of chlorides of alkali metals is as follows :
LiCl < CsCl < RbCl < KCl < NaCl
3. The melting point
of LiCl is lowest because it is with highest covalent character.
4. The solubility of
alkali metal carbonates in water at 298 K increases down the group from Lithium
to Cesium.
5. The basic
character of oxides and hydroxides of group 1 and group 2 increases down the
group because metallic character increases down the group e.g., LiOH is
least basic whereas CsOH is most basic. Be(OH)2 is amphoteric, Mg(OH)2 is a weak base, Ca(OH)2 and Sr(OH)2 are moderately strong bases, Ba(OH)2 is strong base.
6. The solubility of
hydroxides of Group 1 and Group 2 in water increases down the group.
7. The solubility of
sulphates, carbonates and phosphates decreases down the Group 2 because lattice
energy dominates over hydration energy in Group 2, for example MgSO4 is soluble in water wherease BaSO4 is insoluble in water.
8. Li2CO3 is thermally unstable whereas other alkali
metal carbonates are thermally stable.
9. Thermal stability
of carbonates of Group 2 increases down the group. All are thermally unstable.
10. Properties
of Li almost similar to that of Mg, Be are almost similar to that of Al and B
are almost similar to that of C due to diagonal relationship.
Problems for practice
Download app for all tests and notes
Click here for NEET/IIT-JEE and other competitive exam books
Problems for practice
Q1. Choose the correct
statement
(1) Na
is heavier than K
(2) Thalium
have higher 1st ionisation energy than indium
(3) Electron
affinity of Be is higher than B
(4) electronegativity
order is F > O > N > Cl
Q2. The correct arrangement regarding size
(1) Tl
> In > Al > Ga
(2) Tl
> In > Ga > Al
(3) Al
> Ga > In > Tl
(4) Ga
> In > Tl > Al
Q3. Which of the following is associated with
biggest jump between 1st and 2nd ionisation energy?
(1) 1s2, 2s1
(2) 1s2, 2s2,
2p6, 3s1
(3) 1s2, 2s2,
2p6, 3s2, 3p6,
3d10, 4s1
(4) 1s2,
2s2, 2p6, 3s2,
3p5
Q4. Which process is exothermic out of following?
(1) A
®
A+ + e– (2) B + e– ® B–
(3) B– + e– ® B–2 (4) A+ ® A+2 + e–
Q5. Which is incorrect statement?
(1) Electron
affinity of carbon is higher than silicon
(2) Electron
affinity of fluorine is higher than chlorine
(3) Electron
affinity of oxygen is smaller than sulphur
(4) Electron
affinity of phosphorus is higher than nitrogen
Q6. Which of the following ion have highest
polarising power?
(1) Na+ (2) Te+
(3) Cu+ (4) K+
Q7. Which of the following pair of element have
zero electron affinity?
(1) K,
B (2) N and oxygen
(3) He
and C (4) Ar and Na
Q8. An element
is having 118 proton and 180 neutron in nucleus. It belongs to
(1) 7th period, 18 group
(2) 6th period, 18 group
(3) 8th period, 1 group
(4) 7th period, 17 group
Q9. Choose the correct isoelectronic pair
(1) CH3+ and NH3 (2) CH3– and NH3
(3) SO2 and CO2 (4) CH4 and NH2OH
Q10. Correct statement about valency and oxidation
state is
(1) In
H2O2 valency of oxygen is 2 but oxidation state is
–2
(2) In
H2O valency of oxygen
is 2 but oxidation state is –2
(3) In
N2H4 valency of nitrogen is 3 but oxidation state
is –3
(4) In
PF5 valency of phosphorus is 5 but oxidation state
is +3
Q11. The correct order of electronegativity
(1) F
> Cl > O > N
(2) Mg
< Be < C < B
(3) Mg
< Al < S < Cl
(4) B
< C < O < N
Q12. If element (A) have oxidation state + 6 and
element (B) have oxidation state –2 and (C) have oxidation
state –1. The formula of compound will be
state –1. The formula of compound will be
(1) ABC4 (2) AB3C
(3) ABC5 (4) A2B5C
Q13. Choose the correct statement
(1) Mendeleev Periodic table have 18 group and
7 period
7 period
(2) Cu,
Ag, Au are the member of I B group in Mendeleef periodic table
(3) He,
Ne, Ar are the member of zero group in modern periodic table
(4) III A and III B are continuous in
Modern periodic table
Q14. An element is having
electronegativity greater than two and it is a non-metal, to which block of
modern periodic table it belongs?
(1) s (2) p
(3) d (4) f block
Q15. An element having electronic configuration –1s2, 2s2, 2p6,
3s2, 3p6, 3d10,
4s1 to which group and period it belong?
(1) 11th
group, 3 period
(2) 11th
group, 4 period
(3) 12th
group, IInd period
(4) 12th
group 4 period
Q16. Which is amphoteric oxide?
(1) MgO (2) Al2O3
(3) SiO2 (4) Cl2O7
Q17. The incorrect order of ionic radii is
(1) Fe+3 <
Fe+2 < Fe (2) O–2 < N–3 < F–
(3) Li
< Mg < K (4) Be < Mg < Na
Q18. Element of IA group give flame colour due to
(1) Low
IE
(2) Low
boiling point
(3) Low
melting point
(4) All
of these
Q19. An element having elctronic configuration 1s2, 2s2,
2p6, 3s2, 3p1 forms
(1) Acidic
oxide (2) Amphoteric oxide
(3) Basic
oxide (4) Nautral oxide
Q20. An element have atomic number 24. It belong to
(1) 2nd group (2) 4th group
(3) 6th group (4) 8th groupDownload app for all tests and notes
Click here for NEET/IIT-JEE and other competitive exam books
No comments:
Post a Comment