Download Chapter 2 Periodic Properties of the Elements

Chapter 2
Periodic Properties of the Elements
•
the periodic table
•
effective nuclear charge
•
atomic and
ionic radii
•
ionization
energy
•
electron
affinity
2.1 Development of the Periodic Table
The periodic table is the most significant tool that chemists use for organizing and
recalling chemical facts.
elements are arranged in
order of increasing
atomic number
arrangement reflects trends in chemical and physical properties
The periodic table arises from the periodic patterns in the electronic configurations of the
elements
- elements in the same column contain the same number of valence electrons
- trends help us make predictions about chemical properties and reactivity
Mendeleev published the first modern PT in 1869, and in it he predicted the
existence of:
eka-silicon (germanium, Ge)
eka-aluminum (gallium, Ga)
predicted
actual
predicted
actual
2.2 Effective Nuclear Charge
Effective nuclear charge (Zeff) is the charge experienced by an electron on a manyelectron atom.
It’s not the same as the charge on the nucleus because of the effect
of the inner electrons.
The electron is attracted to the nucleus, but repelled by electrons that shield or screen
it from the full nuclear charge.
Zeff = Z – S
S is called the screening constant or shielding
constant. Shown at right for sodium
Na is 1s22s22p63s1
represents the portion of the
nuclear charge that is screened
from the valence electron by
other electrons in the atom
S (screening) is usually close to
the number of core electrons in
an atom
Core electrons are effective at screening the outer electrons from the full charge of the
nucleus.
- electrons in the same shell do not screen each other effectively
- Zeff experienced by valence electrons increases as we move left to right
- Zeff increases slightly down a
column because the screening is
always less than perfect
no peak here: 2p electron
stays far from nucleus
Zeff for a given n decreases
with ℓ. E (ns) < E(np) < E(nd)
peak here means 2s
electron spends some time
close to nucleus. Thus
2s electrons screen 2p
electrons
2.3 Sizes of atoms and ions
We can define atomic size in several ways, based
on distances between nuclei in different
situations.
electron distribution
in molecule
non-bonding radius =
half the closest distance
between nuclei during a collision
bonding radius =
half the distance between
nuclei in a compound = ½ d
Knowing atomic radii allows estimation of
bond lengths in molecules
e.g. C-C (diamond) = 1.54 Å = 154 pm
distance
between
nuclei = d
Cl-Cl (Cl2) = 1.99 Å = 199 pm
estimate C-Cl distance in CCl4 as 0.5 × (1.54 + 1.99) = 1.77 Å = 177 pm
Atomic size varies consistently through the periodic table
As we move down a group the atoms become larger.
Increase in principal quantum number, n – electrons are more likely to be further from
the nucleus
As we move across a period atoms become smaller.
Increase in Zeff – valence electrons are drawn closer to the nucleus
increasing Zeff
increasing n
Ionic size also important in predicting lattice energy and in determining the way in which
ions pack in a solid. Like atomic size, it is also periodic. Radii in Angstroms (Å = 10-10 m).
Cations are smaller than
their parent atoms
- electrons removed from
largest orbital
- Zeff has increased
Anions are larger than their
parent atoms
- electrons added to largest
orbital
- total electron-electron
repulsion has increased
For ions with the same
charge, ionic size increases
down a group.
All the members of an isoelectronic series have the same number of electrons
As nuclear charge increases the ions become smaller
O2– > F– > Ne > Na+ > Mg2+ > Al3+
Ionic size important in many applications: lithium ion batteries work because of the small size
and low charge of the Li+ ions allows them to migrate from cathode to anode.
2.4 Ionization Energy
The ionization energy is the minimum energy required to remove an electron from the
ground state of the isolated gaseous atom or ion.
I1 = energy required to remove an electron from a gaseous atom
e.g. Na(g) → Na+(g) + e I2 = energy required to remove an electron from a gaseous 1+ ion, etc
The larger the ionization energy, the more difficult it is to remove the electron.
positive values – energy is required to remove electrons
Ionization energies for an element increase in magnitude as successive electrons
are removed.
Note sharp increase in ionization energy when a core electron is removed.
Ionization energy generally increases across a period.
Zeff increases, making it more difficult to remove an electron
Two exceptions are removing the first p electron and removing the fourth p
electron
s electrons are more effective
at shielding than p electrons
N
O
P
Mg S
Be
Zn
B
Al
Ga
Ionization energy decreases down a
group
Small atoms have high IE
easier to remove an electron
from the most spatially extended
orbital
removal of 4th p electron easier
than expected due to electronelectron repulsion
Electron configurations of ions are derived from the electron configurations of elements
with the required number of electrons added or removed from the most accessible
orbital.
Li: [He] 2s 1
becomes
Li+: [He]
F:
becomes
F-:
[He] 2s 2 2p 5
[He] 2s 2 2p 6 = [Ne]
Transition metals tend to lose the valence ns electrons first and then as many d electrons
as are required to reach the charge on the ion.
Electrons are removed first from the orbital with largest principal
quantum number n, e.g. 4s then 3d
2.5 Electron Affinities
Electron affinity is the energy change when a gaseous atom gains an electron to form a
gaseous ion. Electron affinity and ionization energy measure the energy changes of
opposite processes:
Electron affinity:
Cl(g) + e – → Cl–(g) ΔE = –349 kJ/mol
Ionization energy:
Cl(g) → Cl+(g) + e –
ΔE = 1251 kJ/mol
Energy is usually released when an atom adds an electron. Eea is negative for stable
anions.
Electron affinities do not change much as we move down a group
Going down a group, the attraction of the added electron to the nucleus is
less – but so is the electron-electron repulsion.
Electron affinity becomes more
exothermic as we move left to right
across a row
Attraction of the electron
to the nucleus goes up with
increasing charge on the
nucleus.
Note discontinuities when
entering new subshell
Be -> B, Ne -> Li
or when pairing electrons
in p orbitals
N -> O