Notes on Periodic Classification of Elements
1.
Position
of hydrogen: It has been confirmed that
hydrogen has many characters which match with halogen elements too. But it is
only placed in Group IA. Thus position of hydrogen is not clear.
2. Position
of isotopes: Since in Mendeleev’s
periodic table elements have been arranged according to their atomic masses,
isotopes (more than one element) are placed at a one place which is confusing.
3.
Position
of the isobar: Since in Mendeleev’s
periodic table elements have been arranged according to their atomic masses,
different elements having different atomic numbers but same mass number
(isobars) are placed at one place which is again confusing.
4.
No
attempts were made to separate metals from non-metals.
5.
Dissimilar
elements placed together in the same group:
Elements showing different chemical properties have been placed together
in the same group. For example Li, Na, k and Cu, Ag, Au have been placed in
Group IA.
Modern Periodic Law:
Moseley’s
experiment confirmed that atomic number of an element and square root of the
frequency of X-ray of that element (which is obtained when various elements
were bombarded with cathode rays) are directly proportional to each other. This
triggered the chemist to modify the old periodic law as:
“The physical
and chemical properties of elements are the periodic function of their atomic
numbers. Clearly when elements are arranged in increasing order of their atomic
numbers similar elements are repeated after regular intervals.”
Periodicity of Elements (Cause of Periodicity):
The
recurrence of elements with similar properties after certain regular intervals
when these are arranged in the increasing order of their atomic numbers is
called the periodicity.
Looking
at the elements of alkali metals in the modern periodic table, it’s not
difficult to realize that every element here has a similar outer electronic
configuration. This repetition in similar outer electronic configurations after
certain regular interval (as the atomic number increases while arranging the
elements in the periodic table) is the cause of periodicity in properties of
elements. This is because the chemical behavior of elements is due to the
electrons in the outer most shells. Since all the elements in a particular
group have similar outer electronic configuration their chemical behavior are almost
same.
3Li = 1s2 2s1 or
[He]2s1
11Na = 1s2 2s2 2p6
3s1 or [Ne]3s1
19K = 1s2 2s2 2p6
3s2 3p6 4s1 or [Ar] 4s1
37Rb = 1s2 2s2 2p6
3s2 3p6 3d10 4s2 4p6 5s1 or [Kr]
5s1
Note: The numbers 2, 8, 8, 18, 18, 32 are called Magic
Numbers because the properties of elements get repeated in such intervals.
For example atomic number of hydrogen is 1 (1H) , the next element in alkali metals in the periodic
table is Li whose atomic number is 3 (1+2=3),
the next one is sodium of atomic number 11 (3+8=11), next is potassium of atomic number 19 (11+8=19), next 37Rb (19+18=37)
and so on.
For elements in any
group of p-block (except inert gas elements for which it is 8, 8, 18, 18, 32, 32) this is 8, 18, 18, 32 and 32. For elements in
the first group of d-block (Sc, Y, La, Ac) this is 18, 18, and 32 and for all other groups in this block this is 18, 32, and 32. For f-block this is 32.
Long Form of the Periodic table
(Present Form or Bohr’s table):
Features of Long Form of the Periodic
Table:
1. All elements have been arranged in
the increasing order of their atomic numbers.
2.
Elements of equivalent properties due to their similar electronic
configuration have been placed together at one place.
3.
For a better classification of elements, the periodic table has been
classified into the following sections and sub-sections:
i. Four
Blocks (s, p, d, f)
ii. Seven
Horizontal Rows called periods, each represented by a unique principal quantum
number, n (i.e., n=1 for first period, n=2 for second period, and so on….).
iii. Eighteen vertical columns called groups, numbered as 1 to 1
Classification of elements into
blocks and general electronic configuration:
On the basis of electronic
configuration of atoms of various elements (more specifically based upon the
name of the orbital which receives the last electron), elements have been
divided mainly into three types:
i. Representative
or s and p-block elements
ii. Transition
or d-block elements
iii. Inner
transition or f-block elements.
s-Block or the first (left) block of
representative elements:
The elements (except
Helium) in which the last electron enters
into the s-subshell of the outer most main energy level are called s-block
elements.
Maximum number of
electrons that can
be accommodated in s-subshell = 2 (by
the formula 2(2l+1)). Therefore there are
only two groups in this block. The first one being consists of the “hydrogen and alkali metals’’ and
the second one is of the ‘alkaline earth
metals’.
Consider an element
sodium, which is present in the first group (called The Alkali metals) of this
block, 11Na = 1s2
2s2 2p6 3s1
or [Ne]3s1. It has been previously
mentioned that all the elements in a group have equivalent outer electronic
configuration. The only difference is that, in hydrogen the last electron
enters into 1s1 (because hydrogen is present in first
period), in lithium it is 2s1
(because hydrogen is present in second
period) and in sodium it is 3s1
(because hydrogen is present in third
period) and so on. Thus the general electronic configuration of
the alkali metals is ns1 (‘n’
representing the number of period in which the element is present).
Again consider an element from the
second group of this block (The alkaline
earth metals) calcium, 20Ca =
1s2 2s2 2p6 3s2 3p6 4s2
or [Ar] 4s2. It can easily be observed that all the elements in
the group of alkaline earth metals have the following outer electronic
configurations, beryllium (2s2,
because in 2nd period), magnesium (3s2, because in 3rd
period), calcium (4s2,
because in 4th period)
and so on. Thus the general electronic configuration of the alkaline earth
metals is ns2.
Combining the general electronic configuration of alkali
metals and alkaline earth metals, we can write the general electronic
configuration of s-block as ns1-2.
General Characteristics of s-block elements:
1. Soft metals and low melting points
2. Low ionization energies
3. Highly reactive and form univalent (alkali metals) and divalent (alkaline
earth metals) ions readily.
4. 4. Compounds of these elements (except beryllium) are predominantly ionic.
5. 5. Most of the metals (except beryllium and magnesium) impart characteristic
color to the flame which can be used to identify them.
6. 6. Strong reducing agent.
7.
Good conductor of heat and electricity.
8.
Metal hydroxides are strong bases.
9.
The last elements in both group-1 and group-2 in this block namely
francium and radium are radioactive.
1 Very low electron affinity.
p-block elements or the second (right)
block of representative elements:
The
elements (except helium) in which the last electron enters into the p-subshell
of the outermost main energy level, are called the p-block elements. e.g 17Cl = 1s2 2s2 2p6
3s2 3p5 or [Ne] 3s2 3p5
Maximum number of electrons that can be
accommodated in p-subshell = 6 (l=1
for p-subshell). Therefore there are six
groups in this block. They are boron family, carbon family, nitrogen
family, oxygen family, halogen family, and 18th group elements.
Consider an element from boron family say 13Al = 1s2 2s2
2p6 3s2 3p1
or [Ne] 3s2 3p1. Now we can easily predict the outer
electronic configurations of all the elements present in boron family; 5B = 2s2 2p1 (in 2nd period),
13Al = 3s2 3p1
(in 3rd period), 31Ga = 4s2 4p1 (in 4th period) and so on. Thus the general electronic
configuration of boron family is ns2
np1 (‘n’ representing the number of period in which the
element is present). Similarly we
can write for the carbon, nitrogen, oxygen and halogen family and also for
inert gas elements as, ns2 np2,
ns2 np3, ns2 np4, ns2 np5
and ns2 np6 respectively. Thus the general
electronic configuration for p-block is ns2 np1-6.
General characteristics of p-block elements:
1.
Contains
both metals and nonmetals (nonmetallic character increases towards right)
2.
High
ionization energies compared to s-block elements
3.
Form
mostly covalent compounds
4.
Some
of the elements show variable valency
5.
Oxidizing
character increases towards right in a period and reducing character increases
top to bottom in a group.
6.
Generally
bad conductor of heat and electricity
Note: Diagonal relationship:
Elements of 2nd period are
called bridging elements. They resemble in certain properties with the elements
of the third period diagonally placed,
From
the above figure, Lithium has similarities in properties with Magnesium;
beryllium has similarities in properties with aluminum, boron with silicon and
so on.
d-block elements or transition elements:
The
block in between s and p-block is called d-block. Since the properties of these
elements are intermediate between s and p-block elements these are called
transition elements. Maximum number of electrons in a d-subshell is equal to
ten. Hence there are 10 groups in this block.
Elements
in which the last electron enters to a penultimate (n-1) d-subshell are called
d-block elements. Strictly speaking, elements which contain incompletely filled
d-orbital are included in this category.
This
block consists of four periods, n=4, n=5, n=6 and n=7, (i.e., 4th, 5th,
6th, and 7th periods of periodic table) each representing one
transition series in which the last electron enters into 3d, 4d, 5d and 6d
subshell respectively. They are called “First transition series (n=4), Second
transition series (n=5), Third transition series (n=6), Fourth transition
series (n=7)”
Consider
an element from first transition series, 21Sc
= 1s2 2s2 2p6 3s2 3p6 3d1
4s2 or [Ar] 3d1
4s2, we see that, after the filling of 18 electrons in the argon
core, the 19th and the 20th electron normally enter into
the 4s (ns) and the rest into the 3d (n-1d) subshell. This trend is followed by
all other elements in all the transition series of this block. Thus the general electronic configuration of d-block
elements is (n-1) d1-10 ns1-2.
Due to the stabilities of half filled
and full filled orbital, in some elements like Cr and copper the (n-1)d and ns
subshell show irregular filling of electrons
compared to other elements. e.g., 24Cr=
[Ar] 3 d5 4s1 and
29Cu = [Ar] 3 d10 4s1
There are some exceptions to the
above rule of assigning an element as d-block element, e.g., Lutetium, 71Lu = [Xe] 54 4f14 5d1 6s2, but this element due to
its properties resemble more with inner transition elements (f-block elements)
is included in f-block elements.
General characteristics of d-block
elements:
1.
Hard,
malleable and ductile
2.
Good
conductor of heat and electricity
3.
Show
variable oxidation states
4.
Magnitude
ionization enthalpy is in between those of s and p-block elements
5.
Form
both ionic and covalent compounds
6.
Compounds
are generally colored and paramagnetic
7.
Strong
tendency to form complexes which are coloured
8.
Elements
like Cr, V, Mn, Fe, Cu, Ni, Pd, Pt are used as catalysts.
9.
Most
of the elements are used in the preparation of alloys.
10. Complexes are colored.
We can see from the periodic table
that after *57La comes *72Hf. More clearly *58Ce to *71Lu (from 6th
period), these fourteen elements have been brought down the main part of the
modern periodic table as “lanthanoids” as all these elements follow Lanthanide,
*57La. Similarly *90Th to *103Lr (from 7th
period) these fourteen elements have
been brought down the main part of the modern periodic table as “Actinoids” as
all these elements follow Lanthanide, *89Ac.
In all these elements, the s-orbital of the last shell (n) is completely
filled, the penultimate (n-1) d-orbitals invariably contains zero or one
electron and the ante-penultimate (n=2) f-orbitals (being lower in energy than
the said d-subshell) gets progressively filled in.
Thus the general
electronic configuration for f-block elements is (n-2) f 1-14 (n-1)
d 0-1 ns2. (n=6 for lanthanoids and n=7 for actinoids).
58Ce = [Xe] 54 4f 1 5d1 6s2, 90Th = [Rn] 86 5f 1 6d 1 7s2
1.
These
are Heavy metals
2.
High
melting and boiling points
3.
Show
variable oxidation states
4.
Form
generally colored compounds
5.
High
tendency to form complexes
6.
Most
of the elements are radioactive
Thus we found that total number of groups in modern periodic table = 2 (s-block) + 10 (d-block)+ 6(p-block.) = 18.
Number of elements in a period:
First
period: This corresponds
to n=1, filling of electrons in 1st shell (K), K shell contains only
one subshell,1s, which can accommodate only 2 electrons, therefore there are
only two elements in the first period. They are hydrogen and helium.
Second period: This corresponds to n=2, , filling of electrons in 2nd
shell (L), L shell contains two subshell,2s and 2p, which can accommodate 2+6=8
electrons, therefore there are eight elements in the second period (Li-Ne).
Third period: This corresponds to n=3, (M shell), the subshells for
n=3 are 3s, 3p, 3d but we know, energy of 3d is greater than 4s. thus 3d is
excluded, thus total electrons to be accommodated are due to 3s (2 electrons) + 3p (6
electrons) = 8 electrons. Therefore there are eight elements in the third period
(Na-Ar) again.
Fourth period: n=4, N shell contains four subshells, 4s, 4p, 4d,
4f. Since energies of 4d and 4f are greater than 5s hence they are excluded.
3d, excluded from third period is included in the fourth period. Thus total
electrons to be accommodated are due to 3d (10 electrons) + 4s (2 electrons) +
4p (6 electrons) = 18. Hence number of elements in the fourth period is 18.
In a similar way we can
show how fifth, sixth, and seventh period contain 18, 32 and 32 elements.
General trends in periodic properties:
1. Electronic configuration: So far we have already discussed the
general electronic configuration of elements present in different groups, while
discussing about various blocks. The general electronic configurations of all the
elements in the same group are similar. The chemical properties of elements depend on the valence electronic configuration.
2. Atomic radii: It is defined as the distance from
the centre of the nucleus to the outer most electron(s).
It is measured by electron diffraction method in angstrom (A°) or picometer (pm) unit.
The
difficulties in measuring the exact atomic radius of an atom are that, except
inert gas elements no other atoms are generally found in isolated or uncombined
state. The exact position of the outer most electrons is uncertain according to
Heisenberg’s uncertainty principle. Furthermore the electron density of an atom
is affected by the presence of neighboring atoms.
Despite the
above limitations, we still need some practical approach to estimate the size
of atoms. The hope is alive as the atoms pack up at certain definite distances
in solids. This gives the idea about the approximate atomic size or radius of
the atom. In this concern, depending upon whether the element is a metal or a
nonmetal, three different types of atomic radii can be discussed.
i. Covalent radius:
a. Covalent Single Bond Radius For Homonuclear Molecules: It is defined as one half of the distance between the nuclei of two covalently bonded atoms (bonded with a single bond) of the same element in a homonuclear molecule such as H2, F2, Cl2 etc.
r = 1/2 (Internuclear distance between two bonded atoms)
r
= 1/2 (bond length) , e.g. the internuclear distance between two chlorine atoms in
Cl2 molecule is 1.96 A°.
Thus the covalent atomic radius of a chlorine atom is 1.96/2 = 0.98 A°.
b. Covalent Single Bond Radius For Heteronuclear Molecules where the electronegativity difference of corresponding bonded atoms is not so high: The covalent single bond radius of an atom X in a heteronuclear molecule is defined as the difference between the single bond length X - Y and the covalent radius of other single bonded atom, Y.
Covalent radius of X in the single bond X - Y = Bond length of X -Y - Covalent radius of Y in X - Y
For example, the bond length of Si - C bond (d Si - C) is 193 pm and covalent radius of carbon atom (r C) is 77 pm. Thus the covalent single bond radius of Si = (d Si - C) - r C = 193 - 77 = 116 pm
Let us now calculate the internuclear distances of elements by adding the tetrahedral radii and compare them with corresponding observed value in cubic or hexagonal crystals.
AlP (Aluminium phosphide): The sum total of tetrahedral radii of Al and P taken from the above table = 126 + 110 = 236 pm and the observed value is also 236 pm.
In some other elements the sum total of the tetrahedral radii and the observed radii deviate by 1 pm. e.g., HgTe , The sum is = 148 + 132 = 280 pm but the observed internuclear distance between Hg and Te is 279 pm.
In some other elements there exist a large deviation from the calculated values. The addition of tetrahedral radii in SiC gives 194 pm whereas the observed value is 189 pm.
e. Octahedral Covalent Radii: Crystals having arrangement of lattice points similar to pyrite (FeS2) structure fall into this category.
Arrangement of lattice points where one atom is surrounded by 6 other atoms (similar or dissimilar atoms) octahedrally give rise to octahedral radii.
ii. Van der Waals radius:
It is defined
as one half the distance between the nuclei of two identical non bonded
isolated atoms or distance between two adjacent identical atoms belonging to two
neighboring molecules of an element in solid state.
To make it
clear we can take the example of chlorine in solid state. The internuclear
distance between adjacent chlorine atoms of the neighboring molecules is 3.6 A°.
Thus the Van der Waal's radius (r van
der Waal's) is equal to half of 3.6 A°, i.e., 1.8 A°.
Note: since
inert gas elements normally do not form chemical compounds, their atomic radii
are usually expressed in terms of Van der Waal's radii which are sometimes
referred to as Inert Gas Radii.
Note: the magnitude of various atomic radii is in the order as
follows:
Van der Waal's >
metallic > covalent
Factors on which atomic radii depend:
a.
Hybridization: This can be demonstrated by taking
suitable examples.
Radius of carbon atom is 77pm in methane. Here ''C’’ is sp3 hybridised
(s character 25%). Radius of carbon atom is 67pm in ethane. Here “C” is sp2 hybridised (s character 33%). Radius
of carbon atom is 60pm in ethyne. Here “C”
is
sp hybridised
(s character 50%).
Greater is the
s-character of an atom, smaller the atomic radius of that atom.
b.
Nuclear charge: Greater is the nuclear charge
(increase in atomic number means increase in number of protons and hence
increase in the nuclear charge) of the nucleus, it attracts the electrons more
towards it and the atomic size decreases.
c.
Number of orbits (shielding effect): With increase in the number of orbits,
distance between the nucleus and last orbit increases and the electrons present
in between nucleus and the last orbit tend to decrease the attractive force of
the nucleus by screening the nucleus (shielding or screening effect). Thus the
atomic radius increases.
d. Bond order: Greater is the bond order between the bonded atoms, shorter will be bond length due to greater extent of overlapping. This is because their p orbitals overlap to a greater extent and the bonded atoms come closer.
e. Other factors which affect the atomic radius are the nature of bond (whether ionic or covalent or metallic) and the oxidation states of the bonded neighbouring atoms.
Variation of atomic radii:
Variation along a period: As we move from left to right across a period, there is regular decrease in atomic radii of the representative (s and p-block) elements. This is due to the fact that number of energy shells remains the same in a period (electrons are added in the same shell) but nuclear charge increases gradually as the atomic number increases. Thus the force of attraction of the nucleus on the electrons increases which brings contraction in size.
This can also be explained on the basis of effective nuclear charge which increases gradually in a period. (Click and zoom the figure)
|
Li(1.23 A0) |
Be(0.89 A0) |
B(0.80 A0) |
C(0.77 A0) |
N(0.75 A0) |
O(0.73 |
F(0.72 |
Ne(1.2A0) |
|
Na(1.54A0) |
Mg(1.36A0) |
Al(1.20 A0) |
Si(1.17 A0) |
P(1.10 A0) |
S(1.04 |
Cl(0.99 |
Ar(1.91A0) |
|
K(2.03 A0) |
Ca(1.74 A0) |
Ga(1.26A0) |
Ge(1.22A0) |
As(1.20A0) |
Se(1.16 |
Br(1.14 |
Kr(2.0 A0) |
Variation in a group: atomic radii in a group increases as
the atomic number increases. The increase in size is due to extra energy shells
(a new shell is introduced in elements each time when we move down the group)
which dominates over the increased nuclear charge.
The atoms of 18th group elements do not form chemical bonds. Hence,
their van der Waaals’ radii are considered which are always greater than the
covalent radii of halogens. The following table illustrates the above three
points:
Variation in transition elements: Though, in vertical columns of
transition elements, there is an increase in the size from first member to
second member as expected, there is an abnormal change when we move from second
to third member of a column. Either a very small change or no change is
observed. This is due to the Lanthanide
contraction. In the lanthanide elements the last electron (differentiating
electrons) enters into 4f subshells which do not effectively screen the
nucleus. This result in the increase of effective nuclear charge and the size
gradually decrease in size. The covalent radii in angstrom units are given
below:
|
Sc(1.44) |
Ti(1.32) |
V(1.22) |
……………………..Cu(1.17) |
|
Y(1.62) |
Zr(1.45) |
Nb(1.34) |
……………………..Ag(1.34) |
|
La(1.69) |
Hf(1.45) |
Ta(1.34) |
……………………..Au(1.34) |
3. Ionic Radii: It is defined as the distance between the nucleus and the outer most electron of that ion in an ionic bond or it is the distance from the nucleus of an ion up to which it has an influence on its electron cloud.
The internuclear distance between the two oppositely charged ions touching each other in a crystal can be determined using X-ray technique. But measuring ionic radii is not easy because it is nearly impossible to define a clear boundary between the cation and anion and up to which the electron clouds of the oppositely charged ions are spread.
Still the hope remains alive as X ray crystallography can measure the electron density (ED in electrons per angstrom cube) in the region surrounding the nucleus which is known as the electron density map (the ED map).
Similarly the size of
an anion is always greater than the corresponding atom.
Note: in a set of species having same number of electrons (isoelectronic species), the size decreases as the charge on the nucleus increases. Ex: O2-> Ne > Mg2+ > Al3+.
4. Ionization enthalpy or ionization potential or ionization energy (I.E/I.P/Δi H): it is defined as the minimum potential difference maintained in a discharge tube to remove the most loosely bound electrons from an isolated gaseous atom to form gaseous cation or the minimum amount of energy required to remove the most loosely bound electrons from an isolated gaseous atom to form a monovalent positive ion. It is also known as the first ionization enthalpy or energy.
1 ev = 1.602 x 10-19 j per
atom = 1.602 x 10-19 x 6.02 x 1023 x 10-3 kj
per mole = 96.49 kj mol-1 = 23.06 kcal per mole
Like the removal of
first electron from neutral gaseous atom, we can remove second, third and
successive electrons from positive ions one after another. The amount of
energies required to do so are termed as second, third, fourth….ionization
energy respectively.
Factors affecting ionization
enthalpy:
a. Size of the atom: ionization
energy decreases as the atomic radius increases. This is because the distance
between the nucleus and the outer most electron increases and the attractive
force on the outer most electron decreases and that electron can now be removed
with a lesser amount of energy.
b. Nuclear charge, shielding or
screening effect and effective nuclear charge: the
ionization energy increases with increase in nuclear charge (increased numbers
of protons in the nucleus) because the attractive force on the outer most
electron increases.
In
multi-electron atoms, the attractive force exerted by the nucleus on the
valence shell is partly reduced by the presence of inner shell electrons. Thus
the valence electrons do not feel the full charge of the nucleus. The actual charge felt by the valence shell electrons is
called the effective nuclear charge and the partial reduction of the nuclear
charge by the inner shell electrons is called the shielding or screening
effect. They are related as follows:
1.
For the ns or np orbital electrons:
i.
Write the electronic configuration in the
following order and group as:
(1s) (2s, 2p)
(3s, 3p) (3d) (4s, 4p)
(4d, 4f) (5s, 5p) (5d, 5f)
(6s, 6p) ….
ii. ii. Electrons present in group which are right to the
(ns, np) group do not contribute to the shielding effect constant.
iii. The electrons in (ns, np) group contribute to
an extent of 0.35 each to the screening effect constant (except the electrons
in 1s which contribute to an extent of 0.30 each).
iv. The electrons in (n-1) shells contribute to an
extent of 0.85 each.
v. All other electrons in the (n-2) and lower shells contribute to an extent of 1.0 each.
2.
For the d- or f- orbital electrons:
Rules (I)
to (iii) remain the same and (iv) and (v) get replaced by the rule (vi) as:
vi. The
electrons present in the groups lying left to the (nd, nf) group contribute 1.0
each to the screening effect constant.
Ex -1: Calculate
the effective nuclear charge of the nucleus of zinc atom on the electrons in
(i) 4s and (ii) 3d orbital.
Solution:
(i) for 4s- electron:
(1s)2 (2s, 2p)8 (3s, 3p)8 (3d)10 (4s)2
σ * = 1 x
0.35 + 18 x 0.85 + 10 x 1.0 = 25.65
Zeff = Z - σ * = 30 - 25.65 = 4.35
(ii) for d- electron:
σ * = 9 x
0.35 + 18 x 1 = 21.15
Zeff = Z - σ * = 30 - 21.15 = 8.85
c.
Penetration effect of the electrons
(shape of orbitals): since s-orbital is spherical , electrons in s-orbital are more penetrated towards the nucleus. The
order of penetration of various orbital is:
S > p > d > f.
If
penetration of an electron in an orbital is more, it is closer to the nucleus
and attracted more by the nucleus. Thus an electron from a more penetrated
orbital requires more energy to be removed. Ex: comparing Mg and Al,
since last electron is being removed from the s orbital in Mg, it has higher
ionization energy.
d. Electronic configuration (half filled and full filled electronic
configuration) : since
exactly half and full filled electronic configuration are stable, higher energy
is required to remove an electron from
such electronic configurations. For example Be (1s2 2s2)
has higher ionization energy than B (1s2 2s22p1),
also nitrogen has higher ionization energy than oxygen. All inert gaas elements
have very high ionization energy.
Variation of ionization energy:
Along a period: in general, ionization energy
increases as we go from left to right in a period. This is due to the increase
in nuclear charge and decrease in atomic radius. The irregularities observed
can be explained from any one of the factors influencing ionization energy
(discussed above). E.g., Be and b; N and O; Mg and Al etc.
Note: in case of transition
elements, some irregularities are observed along a period as the inner
d-electrons screen the outer most s-electron. Increased screening effect is due
to the increased number of (n-10 d-electrons. As a result the effective nuclear
charge decreases and outer most electrons can be removed with comparatively
smaller ionization energy. Ex: Fe, Co, Ni
Down a group: A regular decrease in the ionization
energy is observed in case of representative elements as we go down a group.
This is due to the increase in atomic size, increase in screening effect and
decrease in effective nuclear charge.
Note: exceptionally, when we move from
second to the third transition series, the corresponding member of the third
transition series has a higher value than the member in the second in the same
group. This is due to the lanthanide contraction, i.e., the atomic radii of
corresponding elements in any group in these series (second and third transition
series) are nearly the same but atomic number differs by 32. Thus the outer
electrons are held firmly and higher energy required in removing such
electrons.
The following table of ionization energy of elements in kj mol-1
confirms all the above mentioned points:
|
Li 520 |
Be 899 |
B 801 |
C 1086 |
N 1402 |
O 1314 |
F 1681 |
Ne 2080 |
|
Na 49 |
Mg 737.6 |
Al 577 |
Si 786 |
P 1011 |
S 999 |
Cl 1255 |
Ar 1520 |
|
Fe 762 |
Co 758 |
Ni 737 |
|
Ru 711 |
Rh 720 |
Pd 804 |
|
Os 840 |
Ir 900 |
Pt 870 |
Ex -2: The first (Δi H1)
and second (Δi H2) ionization energies
(kj mol-1) of some elements are given below:
|
Element |
Δi H1 |
Δi H2 |
|
I |
2370 |
5250 |
|
II |
520 |
7300 |
|
III |
900 |
1760 |
|
IV |
1680 |
3380 |
Predict that who among them is likely
a reactive metal, a metal that can form a stable binary compound with formula
AX2, a reactive non-metal and a noble gas.
5.
Electron affinity or electron gain enthalpy (Δeg H or EA): it is defined as the energy change
that occurs for the process of adding electron to an isolated neutral gaseous
atom to convert it into a negative ion (to form a monovalent anion).
Note: since electron affinity is defined at
absolute zero temperature, therefore, at any other temperature, heat capacity
instead of electron affinity is to be considered. For this reason the two terms
electron gain enthalpy and electron affinity are only same at absolute zero and
at any other temperature the two terms are related as:
The
value of 5/2RT at 298k is just 2.477 kj mol-1 which is very small
and can be ignored. For this reason both the terms are considered to be same
While
for majority of the elements, energy is released (exothermic) when electron is added
to the outer shell of atoms like halogens (attain stability by gaining electrons),
for some other type of elements like noble gas elements energy is gained
(endothermic). In other words noble gas elements have positive electron gain
enthalpies because they attain a highly unstable state by adding electrons.
Like
second and higher ionization enthalpies, second and higher electron gain
enthalpies are also possible. However addition of a second electron to a
monovalent anion so as to make it X2- is difficult as the new
electron now experiences a repulsion from the electron cloud of the atom. Thus
energy must be supplied (or gained by the atom) to add one more electron
against the repulsion, and the second electron gain enthalpy becomes positive. the
following example clarifies the above said concept: [Single click on android to see image clearly]
Ex-3: The electron gain enthalpy of bromine
is 3.36 ev. How much energy in kcal is released when 8 gram of bromine is
completely converted into Bromide ion in the gaseous state?
Solution: number of moles of bromine = 8/80 =
0.1 mole
Required energy = 0.1 x 3.36 x 23.06
kcal mol-1 = 7.748 kcal mol-1
Factors influencing electron gain
enthalpy:
a.
Atomic Size: As
previously explained under ionization energy, the distance between the nucleus
and the last shell receiving the electron increases as atomic size increases.
As a result the force of attraction between the nucleus and the incoming
electron decreases. The electron is added less easily and electron gain
enthalpy becomes less negative.
b.
Nuclear charge or effective nuclear charge: as the nuclear charge increases, the
force of attraction between the nucleus and the incoming electron increases and
hence the electron gain enthalpy increases, i.e., becomes more negative.
c.
Electronic configuration: elements having exactly half filled and fulfilled electronic
configuration are very stable. Adding electron to such elements require energy since
they do not accept electrons easily and on the addition of the electron they
attain an unstable state. Thus electron gain enthalpies of such elements are
positive. for example nitrogen (1s2
2s2 2p3) has +31 kj mol-1 of electron gain
enthalpy.
Variation of electron gain enthalpy: Both in period and group the values of electron gain enthalpies are not regular, still a generalization is being done as follows:
a.
Along a period: in general, the electron gain enthalpy becomes more and more negative as
we move from left to right in a period. This is due to the decreasing atomic
size and increasing nuclear charge. Thus the electron gain enthalpies of the
halogen elements are the highest in a period and those of the noble gas has
positive values.
b.
Down a group: In
general, the electron gain enthalpies of elements in a group become less
negative as we down. This is due to the increasing atomic size and decreasing
nuclear charge.
Note: The electron gain enthalpies of some elements
in the second period (O, F) are less negative than the corresponding elements
in the third period (S, Cl).
Due to the compact size of the O and
F atoms, considerable electron-electron repulsion is already there inside such
atoms. Thus further addition of electron is not as easier as it is the case
with corresponding elements of the group such as S and Cl. As a result electron
gain enthalpies of O and F atoms are less negative than S and chlorine
respectively.
The following table explains the above said points (electron affinities in kj mol-1):
|
Li -60 |
Be +66 |
B -83 |
C -122 |
N +31 |
O -141 |
F -328 |
Ne +116 |
|
Na -53 |
Mg +67 |
Al -50 |
Si -119 |
P -74 |
S -200 |
Cl -349 |
Ar +96 |
|
K -48 |
Ca - |
Ga -36 |
Ge -116 |
As -77 |
Se -195 |
Br -325 |
Kr +96 |
6.
Electronegativity: This is the property of a bonded atom. The relative tendency of an atom
to attract the shared pair of electron towards itself is called the
electronegativity of that atom.
Consider covalently bonded molecules like H2 or F2
. Since the two bonded
atoms are identical, the shared pair of electrons is equally attracted by
nuclei of the two atoms and the electron distribution around the two nuclei is
similar. But when the bonded atoms in a molecule are different like HF, they
attract the shared pair of electron to different extent and the electron
distribution around the two nuclei does not remain symmetrical. The atom
attracting the shared pair more gets a partial negative charge and the other
gats a partial positive charge, as shown below:
a.
Size of the atom: The smaller the size of the atom, the greater the attraction of bonding
electrons, greater is the electroneggativity.
b.
Types of ion: Cations
are more electronegative then the corresponding atom from which it is formed as
the cation has smaller size (for example Li+ has E.N of 2.5 and Li
has 1.0). Similarly anion has less electronegativity than the corresponding
atom from which it is formed as anion larger in size (for example F-
has E.N of 0.78 and F has 4).
c.
Hybridization: The
greater the s-character of the hybrid orbital of an atom greater is the electronegativity
of that atom. Let us compare the basicity of three molecules:
Since greater the s-character greater is the
electronegativity the nitrogen atom, it tends to attract the lone pair of
electron more towards it and availability of lone pair becomes less and
basicity becomes less.
d.
Electron affinity and ionization energy: Mulliken says, the electronegativity
of an atom is one half of the sum total of its E.A and I.E. Thus an element
having higher E.A and I.E should have higher electronegativity.
e.
Effective nuclear charge: As shown by Allred and Rochow, smaller the effective nuclear
charge (or greater the screening effect) smaller is the electronegativity. For
example the electronegativities of the halogen atoms are F(4.0), Cl(3.0),
Br(2.8), I(2.5).
Measurement of electronegativity:
A.
Mullikens’ Scale (from electron affinities and ionization energy): According to Mulliken, the
electronegativity of an atom is one half of the sum total of its electron
affinity, E.A and ionization energy, I.E.
The following example makes it clear:
Ex: 4: calculate
the electronegativity of carbon in C-H bond if EC-H, EH-H
and EC-C are 98.8, 104 and 83 kcal mol‑1 respectively.
E.N of hydrogen = 2.05 [Single Click to see image solution clearly]
C. Allred – Rochow scale: The force of attraction on the electron in the valence shell by the nucleus can be calculated using coulombs’ law:
Where
Z*e = effective nuclear charge felt by the electron
Z*
can be calculated by the Slaters’ rule.
e =
charge on the electron, r = mean radius of the orbital (covalent radius).
Accordingly, the force of attraction
of the atom on the shared pair of electron in a molecule which is known as
electronegativity can be calculated by the formula:
Ex: 5: Calculate the electronegativity of arsenic atom (z=33) having covalent radius 1.21 A° .
Solution: the effective nuclear charge can be calculated from Slaters’ rule and then the formula of Allred Rochow scale can be used to calculate the E.N of the said atom taking r = 1.21 A° .
Trends in electeronegativity:
Along a period: Since the effective nuclear charge
increases from left to right, the E.N also increases as we move from left to right in a period.
For example if we will consider the second period then the lowest value of E.N
is of Li (0.98) where as the highest is that fluorine (3.98).
Down a group: Since the effective nuclear charge
decreases from top to bottom in a group, the E.N also decreases accordingly. For
example if we will consider the alkali metals then the highest value of E.N are
of Lithium (0.98) where as the lowest is that cesium (0.79).
Application of electronegativities:
1. Calculation of percent ionic
character: A
homonuclear diatomic molecule due to its similar electronegativities of atoms
is nonpolar. Its bond can be considered as purely covalent. Whereas a partial
ionic character is developed in a covalent bond joining any two different atoms
differing in their electronegativities. This can be calculated by using
Pauling method:
Percent ionic character = 18 (XA
- XB) 1.4 ; XA - XB = E.N
difference between two atoms A and B.
Note: It has been observed, that two atoms having an electronegativity difference greater than 1.7 (% ionic character = 50%) develop an ionic character in their bond predominantly. In a similar way E.N diff. less than 1.7 gives rise to a covalent bond predominantly.
2.
Calculation of Bond length:
In case a molecule (AB) is
purely covalent (non polar) then Bond length (A-B) =
Covalent radius (rA) of atom A
in A° + cov. Rad. (rB) of
atom B in A°
And in case atoms A and B differ in their electronegativities, bond A-B
becomes polar, increased polarity in the bond results in the shortening of the
bond.
Bond length of a polar bond A-B can be calculated using Shoemaker and Stevenson’s’
formula:
Bond length of a polar bond A-B = rA + rB – 0.09 (XA - XB)
3.
Reactivity: Greater
electronegativity difference also indicates greater reactivity. Let us compare
the reactivity of Cl2 and ICl molecule. Even a small electronegative
difference between I and Cl atom in ICl molecule makes it more reactive than Cl2
which has a covalent bond.
4.
Prediction of acidic hydrogen and bond cleavage of a particular type: Consider an example of acetic acid.
It contains four hydrogen atoms, yet it releases only one hydrogen ion (H+)
to the solution (mono basic acid). This is because the electronegative
difference between the carbon and hydrogen is only 0.35 (2.55 – 2.2 = 0.35),
whereas this difference between hydrogen and oxygen is 1.24 (3.44 – 2.2 = 1.24).
5.
Explanation of why HCl (g) is polar covalent and where as HCl (aq) is
ionic. In HCl (g)
the E.N difference is 0.96 (3.16 – 2.2 = 0.96), which is small and hence it is
polar covalent. But in HCl (aq) water plays a vital role in weakening the force
of attraction between H and Cl atom by a factor 80, (dielectric const. of
water, D = 80) given by the Coulombs’ formula; thus H+ and
Cl- get separated and HCl is ionic by nature.
---------0---------
Expected Long Questions (7-10 marks):
1.What is electronegativity? Factors
affecting it. Trends. Applications.
2.Ionisation energy, Electron affinity,
Slater’s Rule.
Expected
Short Questions (2-3 marks):
Questions may be asked from Atomic radii. Lanthanide contraction. Slater’s Rule. Effective nuclear charge. General electronic configurations of various blocks. Diagonal relationship. Applications of E.N. isotons, isosters, isoelectronic species. Comparision between E.A of F and Cl.
· Which among 3d and 4s subshell should be closer to the nucleus of scandium?
· The first and second ionisaton energies of magnesium atom are 7.64 and 15.03 eV respectively. Calculate the amount of energies required in kj to convert all the atoms in 12g of magnesium.
· The five successive ionization energies of an element are 800, 2427, 3658, 25024 and 32824 kj mol-1 respectively. What is the number of valence electrons in this element?
· The second electron gain enthalpy of oxygen is positive. Explain.
· The values of E.N of atoms A and B are 1.2 and 4.0 respectively. What is the percent ionic character in A—B bond?
· Calculate the effective nuclear charge on a 3d electron of Zn atom.
·
What is the formula for determining the ionic
radii of cation and anion in an ionic compound according to Pauli?
Hints: r+ (radius of cation) = (Effective
nuclear charge of anion in the compound / sum of eff. Nuclear charges of anion
and cation) x internuclear distance in the compound,
Similarly r --
(radius of anion) = (Effective
nuclear charge of cation in the compound / sum of eff. Nuclear charges of anion
and cation) x internuclear distance in the compound
·
Calculate the E.N of arsenic (Z = 33) having
covalent radius 1.20 A0.
Expected Very Short
Questions (1 mark):
1.Diagonal relationship is not shown by
the pair: Mg,Ba; Mg,Na; Mg,Cu; Mg,Cl
2.What is the characteristic electronic
configuration of transition elements?
3.What is the number of elements in the
fifth period of modern periodic table?
4.Ce (Z=58) belongs to which block?
5.Write the electronic configuration of
Eu (Z= 63).
6.Which among Ag+ , K+
, Fe2+and Mg2+ is paramagnetic?
7.Which among oxygen, nitrogen, carbon
and boron has highest ionization potential?
8.The most predominantly electrovalent
compounds are formed between which groups?
9.What are isosters?
10.
What
happens to the basic nature of oxides of elements in going from left to right
in a period?
11.
Write four characteristics of transition
metals.
12.
An
element ‘E’ forming its highest oxide as E2O5 should
belong to which group?
13.
Two
elements A and B have 3 and 6 valence electrons respectively. What should be
the formula of the compound made from them?
14.
Which
among HF, HCl, HBr and HI is a strong acid and why?
15.
What
is the name and atomic number of the last element of the modern periodic table?
16.
What is the relationship between pauling and
Mulliken’s scale of E.N?
17.
What
are the atomic numbers of the 14 elements of lanthanides?
18.
First
member of each group shows anomalous behavior. Explain.
19.
The
first ionization enthalpy of sodium is lower than that of magnesium but the
second ionization energy of sodium is greater than that of magnesium. Explain.
20.
How
can you summarise the nature of acidic strength of oxides of elements along a
period and down the group? Hints: Acidic strength increases left to right in
a period and decreases down a group.
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