Download Chapter 7 Periodic Properties of the Elements

Lecture 3
CHEM102
Periodic Properties
of the Elements
Dr. Noha Osman
Learning Outcomes
By the end of this session, the student should be able to:
1. Identify the trends of properties of elements in the
periodic table.
2. Arrange different elements according to the size of
their atomic and ionic radii.
3. Arrange elements according to their ionization
energies.
4. Arrange elements according to their electronegativity.
Development of Periodic Table
Dmitri Mendeleev and
Lothar Meyer
independently came to
the same conclusion
about how elements
should be grouped.
Development of Periodic Table
Mendeleev, for instance, proposed the existence of an unknown
element (which he called eka-silicon) and predicted a number of
its properties. When Germanium was discovered later, it matched
the predicted properties of eka-silicon remarkably well.
Periodic Trends
• Periodic trends among the elements will be
rationalized with respect to:
 Atomic/ Ionic radii.
 Ionization energy.
 Electron affinity.
Effective Nuclear Charge
Zeff is the nuclear charge felt by an electron when both the
actual nuclear charge and the repulsive effects (shielding) of
other electrons are taken into account.
• In a many-electron atom,
electrons
are
both
attracted to the nucleus
and repelled by other
electrons.
• The nuclear charge that an
electron experiences
depends on both factors.
Effective Nuclear Charge
• The effective nuclear
charge, Zeff, is found this
way:
Zeff = Z − S
where;
- Z : is the atomic number (or
actual nuclear charge)
- S : is the shielding constant,
usually close to the number
of inner (core) electrons.
Atomic Radius
• The size of an atomic orbital cannot
be specified exactly, neither can the
size of an atom.
• The atomic radius is defined as
one-half of the distance between
the nuclei in a molecule consisting
of identical atoms.
Atomic Radius
The bonding atomic radius tends to:
— Decrease from left to right across a row
(due to increasing Zeff).
— Increase from top to bottom of a column
(due to the increasing value of n).
Ionic Radius
Cation
Anion
Loss of 1 or more electrons:
Gain of 1 or more electrons:
- Reduces electron-electron
repulsion
- Increases electron-electron
repulsion
- No change in the nuclear
charge (p+)
- No change in the nuclear
charge (p+)
- Electron cloud shrinks
- Domain of electron cloud
enlarges
The cation is smaller than its
corresponding atom.
Na+ (0.095 nm) vs Nao (0.186
nm)
 The anion is larger than its
corresponding atom
Cl- (0.181 nm) vs Clo (0.099nm)
Practice Excercise
Arrange the following sets of atoms and ions in order of
increasing size using the periodic table
1. Mg, Al, Ca.
2. S, Cl, S23. Fe, Fe2+, Fe3+
1.
2.
3.
Comparison with reference to Mg: Al is to the right to Mg, thus smaller
than Mg. Ca is below Mg, thus larger. The order is Al < Mg < Ca.
Comparison with reference to S atom: Cl is to the right to S, thus is
smaller. The S2- anion is larger than the S atom.  The order is Cl <S < S2-.
Comparison with reference to Fe2+ ion: Feo atom is neutral and thus is
larger than its cations. The Fe3+ ion is charged by loss of 3 e- while Fe2+ is
charged due to loss of 2 e-.  The order is Fe3+< Fe2+< Fe.
THE ELECTRONIC CONFIGURATION OF
ELEMENTS
(only outer shell level of electrons is represented)
THE CHEMICAL PROPERTIES OF ANY ATOM
ARE DETERMINED BY THE CONFIGURATION
OF THE ATOM’S VALENCE ELECTRONS.
Ionisation Energy
• Ionization energy is the minimum energy (in kJ/mol) required
to remove an electron from a gaseous atom in its ground state.
X + energy  X+ + 1 e• The 1st ionization energy (endothermic process = requires E) is
the E change for the removal of the outermost e- from a gaseous
atom to form a +1 ion. The more difficult it is to remove e-, the
larger the ionization energy.
• 2nd and 3rd IE (to remove a 2nd and a 3rd e-) are certainly more
difficult due to the increasing attraction exerted by the nucleus
at every loss of 1 e- after the other. Moreover, there is less
shielding effect due to electron loss, therefore a higher Zeff.
Trends in First Ionization Energies
• As one goes down a
group, less energy is
required to remove the
first electron since the
valence electrons are
farther
from
the
nucleus.
Trends in First Ionization Energies
• Generally, as one goes
from left to right across
a period, more energy
is required to remove
an electron.
– As you go from left
to right, Zeff
increases (electrons
are more tightly held
by the nucleus)
Exceptions to trends of first IE (1st exception)
Between Group 2A and 3A elements in the same period
E.g: Be vs B and Mg vs Al.
IE of 3A < 2A in the same period:
In this case the electron is
removed from a p orbital rather
than an s orbital.
• The electron removed is farther
from the nucleus.
• There is also a small amount of
repulsion by the s electrons
(which decreases the Zeff)
Exceptions to trends of first IE (2nd exception)
Between Groups 5A and 6A elements in the same period
E.g: N vs. O and P vs. S
IE of 6A < 5A in the same period.
In 5A ( ns2np3 ), the p e- are in 3 separate
orbitals while in 6A (ns2np4), the
additional e- is paired with 1 of the 3 p
e-  The proximity of two e- in the
same orbital leads to high electrostatic
repulsion, which makes it easier to
ionize an atom of the Group 6A element,
even though the nuclear charge has
increased by one unit.
Test yourself
Complete:
1. Ionization energy (IE) …………a………….across the periodic table
from left to right and ………b……….. moving down the periodic
table.
2. Comparing the trends of ionic radii and ionization energy it is
clear that there is an ………c………. correlation between them.
3. The ………d…….. the atom, the more tightly its electrons are
held to the positively charged nucleus and the more difficult
they are to remove thus the ………e…….. the ionization energy.
Answers:
a) increases, b) decreases, c) inverse, d) smaller, e) larger, f) large, g) ionization.
Test yourself
Which atom with following electron configurations has the
largest first ionization energy, explain your choice.
1s22s22p6
1s22s22p63s1
1s22s22p63s2
The atom with the largest value of IE is 1s22s22p6 (this is a Ne),
because it is found at the right end of period 2, and the IE
increases from left to right in the same period of the periodic
table.
As for the other 2 configurations including 3s electrons they will
be of lower IE because it is easier to loose e- from a sublevel that
is farther away from the nucleus.
Electron Affinity (Electronegativity)
Electron affinity is the energy change that occurs when an e- is
accepted by an atom in the gaseous state to form an anion.
X + e−  X−
Electronegativity increases from left to right across a period.
This is because the Zeff increases from left to right across a
period.
N.B The halogens (Group 7A) have the highest electronegativity,
however irregularities to this trend exists.
Exceptions to trends in electron affinity
Between Groups 1A and 2A elements in the same period
E.g Li (2s1) vs Be (2s2)
EA of 2A (ns2) < 1A (ns1) in the same period:
• The added electron must go in a higher-energy np orbital,
not an s orbital, which is farther from the nucleus.
• Moreover, the added electron is effectively shielded by the
ns2 electrons, therefore experiences a weaker attraction to
the nucleus.
Exceptions to trends in electron affinity
Between Groups 4A and 5A elements in the same period
E.g C (2s2 2p2) vs N (2s2 2p3)
EA of 5A (ns2np3) < 4A (ns2np2) in the same period:
• The added electron must go into an np orbital that already
contains 1 electron, creating greater electrostatic repulsion.
Exceptions to trends in electron affinity
Group 8A (Noble Gases; ns2np6)
E.g Ne (2s2 2p6)
• An electron added to an atom with ns2np6 configuration has
to enter an (n+1)s orbital, where it is well shielded by the
core electrons, is farther away from the nucleus and
therefore will only be very weakly attracted by the nucleus
(i.e Zeff is v. low).
• This also explains why species with complete valence shells
are chemically stable.
THANK YOU