UNIT 2 PERIODICITY
Structure
2.1 Introduction
Objectives
2.2 Atomic Radii
Covalent Radius
,Van der Waals Radius
Metallic or Crystal Radius
Ionic Radius
Factors Affecting Atomic Radii
Periodicity in Atomic Radii
2.3 Ionisation Energy
Factors Affecting Ionisation Energy
Periodicity in Ionisation Energy Across Periods
Trends in Ionisation Energy Down the Groups
Trends in Second and Higher Ionisation Energies
2.4 Electron Affinity
Factors Affecting Electron Affinity
Periodicity in Electron Affinity
2.5 Electronegativity
Pauling Electronegativity Scale
Mulliken-Jaffe ~ l e c t r o n i ~ a t i vScale
it~
Allred-Rochow Electmnegativity Scale
Periodicity in Electmnegativity
2.6 Summary
2.7 Terminal Questions
2.8 Answers
2.1 INTRODUCTION
In the preceding unit, you have studied that the properties of elements are a periodic
function of their atomic numbers. As a result of their similar valence-shell electronic
configuration, elements belonging to the same group of periodic table exhibit similarities in
properties like valence and formulae of their highest oxides, hydrides and chlorides. Further,
they also exhibit a gradation in properties with their position in the group because of the
valence electrons occupying different electronic levels. For example, in Group 14 (IVA) of
the periodic table, C is a non-metal, Si and Ge are semi-metals. whereas Sn and Pb are
metallic elements. All these elements exhibit the highest oxidation state of 4, as the valence
shells in all of these belong to s2p2type. However, the stability of tetravalent species
decreases down the group and bivalent species become more stable.
In Group 1 (IA)having highly electropositive alkali metals, the reactivity of elements
towards water increases down the group. At 298 K, Li reacts slowly, Na reacts vigorously, K
inflames whereas Rb and Cs react explosively with water. These differences in the
properties arise due to differences in the atomic properties like the atomic size, ionisation
energy, electron affinity and electronegativity.These atomic properties are directly related to
the total electronic configurations of the elements and form an important link between the
properties of elements and their electronic configuration. In this unit you will study the
periodicity in these atomic properties in general and in the following units, you will learn to
make use of these in explaining the trends in the properties of elements in a particular group.
Objectives
After studying this unit, you should be able to:
8 define atomic radii, ionisation energy, electron affinity and electronegativity,
8 discuss the factors affecting atomic radii,
8 describe the relationship of atomic radii with ionisation energy and electron affinity,
8 describe the periodicity in atomic radii, ionisation energy, electron affinity and
electronegativity.
,Periodicity and s- lock'
Elements 2.2 ATOMIC RADII
Atomic radii are a measure of the size of the atoms. Atomic radii are important, because
other atomic properties like ionisation energy, electron affinity and electronegativity can be
related to them. You have studied the wave mechanical picture of an atom, according to
which an atom is composed of a compact nucleus surrounded by an electron cloud. This
electron cloud does not have a definite boundary surface similar to that of a bal1:There is a
definite but very small probability of finding an electron at an infinite distance from the
nucleus of the atom. Does this mean that atom is infinitely large? This just doesnot make
any sense. Thus, we have to find a way to define the size of an atom. The radius of an atom
can be defined as the distance from the centre of the nucleus to the point where the
electron density is virtually zero.
Aftcr we have defined the size of an atom, the problem arises as to how we are going to
measure it? Thus, if we are measuring the size of an atom when it is occupying a lattice site
in the crystal, the value will he different from the one when it is colliding with another atom
in the gaseous state. Furthermore, the si7e of the neutral atom will be different from the one
when it is present as a cation or anion. Consequ~ntly.we cannot have one set of atomic radii
applicable under all conditions. It, therefore, becomes n e c e s s q to specify the bonding
conditions under which the size is being measured. Pertaining to four ~ r , ~ j types
o r of
bonding, atomic radii are of following four types:
i) Covalent radius iii) van der Waals radius
ii) Crystal or Metallic radius iv) Ionic radius.
2.2.1 Covalent Radius
Covalent radius is defined as one half of the distapce oetween the nuclei of two like
atoms bonded together by a single covalent bond. If in a homonuclear diatomic molecule
of A, type (e.g., F2,CI,, Br2, I?), rA-A
is bond length or intemuclear distance and r, is the
covalent radius of atom A, then r, = 'I2rA-,.
The intemuclear distance r,, between two carbon atoms in diamond is 154 pm, so the
covalent radius of carbon, r,, is equal to 77 pm. Similarly, the r,,~,, for solid C1, is 198 pm,
r,, is, therefore, 99 pm. In a heteronuclear, diatomic molecule of AB type, if the bonding is
purely covalent, then the bond length r,., is equal to the sum of covalent radii of A
and B, i.e.,
rA.B= 'A +B
'
Thus covalent radii are additive. It is possible to calculate the radius of one of the atoms in a
heteronuclear diatomic molecule of AB type if we know the intemuclear distance r,~, and
radius of the other atom. For example, the Si-C bond length in carborundum is 193 pm and
covalent radius of C is 77 pm; so you can calculate the covalent radius of Si as follows:
rS,-C= rSi+ rc or rSi= r S i ~-C
o r r s i = 1 9 3 - 7 7 = 116pm
As stated earlier, the above relation holds good only if the bond between the atoms A and B
is purely covalent. If there is a difference in the electronegativities of the bonded atoms, it
causes shortening of the bonds. Schomaker and Stevenson have proposed the following
relationship between the shortening of the bond and the electronegativity difference of the
atoms:
r,., = r, + r, - 0.07 (X, - X,)'
Here X, and X, are the electronegativities of A and B, respectively, about which you will
study in Section 2.5 of this unit.
Multiplicity of the bond also causes a shortening of the bond. Usually a double bond is about
0.86 times and a triple bond about 0.78 times the single bond length for the second period
elements. Covalent radii of the elements are listed in Table 2.1.
, 'Fable 2.1: Covalent and ban der Waals radii of elements
I 7 3 4 5 6 7 8 9 10 11 12 13 14. 15 16 17 18
IA IIA lllB IVB VB VIB VllB VlllB 1B 11B lIlA IVA VA VIA VllA VllIA
H Be
37<- Covalent radius in pm 120
12k- van der Waals rad~us~npm
Li Be B C -N 0, F Ne
123 89 82 77 70 66 64
150 140 135 131
Na Mg A1 Si P S C1 Ar
156 136 125 117 110 104 99
180 190 I85 174
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge' As Se Br Kr
203 174 144 132 122 118 117 117 116 115 117 125 125 122 121 117 114
200 200 195 189
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
216 191 162 145 134 130 127 125 125 128 134 144 144 140 141 137 133
220 220 215 210
Cs Ba La Hf Ta W Re 0 s Ir Pt Au Hg TI Pb Bi Po At Rn
235 198 169 144 134 130 128 126 127 130 134 147 155 154 148 146 - 215
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
165 165 164 - 166 185 161 159 159 158 157 156 170 156
between two fluorine atoms in F,?
..........................................................................
..........................................................................
- <
2.2.2 van der Waals Radius
In the solid state, the non-metallic elements usually exist as aggregates of molecules. The
bonding within a non-metal molecule is largely covalent, yet individual molecules are held
to each other by what is called van der \nTaalsforce. Half of the distance between the
nuclei of two atoms belonging to two adjacent molecules in the crystal lattice is called
van der Waals radius. Table 2.1 also lists the values of van der Waals radii of some
elements.
Figure 2.1 illustrates the difference belween the covalent and van der Waals radii of
chlorine. It is evident from the figure that half of the distance between the nwlei X and X'
of the two non-bonded neighbouring chlorine atoms of adjacent molecules A and B is the
van der Waals radius of chlorine atom, whereas half of the distance between the two nuclei
X and Y in the same molecule is the covalent radius of chlorine atom. Thus van der Waals
radii represent the distance of the closest approach of an atom to another atom it is in contact
with, but not covalently bound to it. Values of van der Waals radii are larger than those of
covalent radii because the van der Waals forces are much weaker than the forces operating
betwee,; ?toms in a covalently bonded molecule.
Fig. 2.1: Covalent and van der Waals radii of solid chlorine
The benefits of buying summaries with Stuvia:
Guaranteed quality through customer reviews
Stuvia customers have reviewed more than 700,000 summaries. This how you know that you are buying the best documents.
Quick and easy check-out
You can quickly pay through credit card or Stuvia-credit for the summaries. There is no membership needed.
Focus on what matters
Your fellow students write the study notes themselves, which is why the documents are always reliable and up-to-date. This ensures you quickly get to the core!
Frequently asked questions
What do I get when I buy this document?
You get a PDF, available immediately after your purchase. The purchased document is accessible anytime, anywhere and indefinitely through your profile.
Satisfaction guarantee: how does it work?
Our satisfaction guarantee ensures that you always find a study document that suits you well. You fill out a form, and our customer service team takes care of the rest.
Who am I buying these notes from?
Stuvia is a marketplace, so you are not buying this document from us, but from seller eyyanshyamu. Stuvia facilitates payment to the seller.
Will I be stuck with a subscription?
No, you only buy these notes for $20.49. You're not tied to anything after your purchase.
