Transcript Periodic properties of the elements

Periodic properties of the
elements
A journey across the table is a
journey through chemistry
The shape of the table is explained
by the orbital model
2
groups
6
groups
10
groups
14
groups
Number of elements per row
matches the capacity of the
orbitals
Relating chemical and physical
properties to the table
• Mendeleev’s table was constructed 50
years before the discovery of the nucleus
• The arrangement of elements was by
atomic weight
• Atomic number (Moseley 1915) provides a
more systematic ordering
• Note: X-ray photon energy varies smoothly
with atomic number, providing a method of
identifying elements
Explaining periodicity: effective
nuclear charge and shielding
• Inner shell (core) electrons screen the
outer (valence) electrons from the nuclear
charge
– Shielding is more effective for electrons with higher
angular momentum (higher l value: 3d > 3p > 3s)
– Decreases as the valence shell fills
• Increased shielding means lower effective
nuclear charge
Effective nuclear charge
• Effective nuclear charge
Zeff = Z – S
– Z = atomic number
– S = screening constant (roughly number of
core electrons)
• Model is approximate because the valence
electrons penetrate the inner shells
– Actual Zeff > Z - S
Zeff and angular momentum
• Electron in 2s orbital spends more time
close to nucleus than electron in 2p
• 2s orbital lower in energy than 2p
• Shielding of 2s lower than 2p
• Shielding follows order: 3d>3p>3s
Shielding and effective nuclear charge
• The “shell” picture helps to explain these
observations
• Electrons in same shell experience stronger
attraction to nucleus as shell fills
Correlate atomic radius and
effective nuclear charge (Zeff)
Atomic radii and periodic variation
Isoelectronic ionic series: same
electrons, different charge
• In a period the ions will tend to be
isoelectronic (all have filled shells)
– Na+, Mg2+, Al3+
– P3-, S2-, Cl-
• For the cations, Zeff ↑, radius ↓ as Z ↑
• For the anions, Zeff ↑, radius ↓ as Z ↑
Ionization energy
• Ionization energy (I):
The energy required to remove an electron
from an isolated gaseous atom
M(g) = M(g)+ + e
• As I increases, atoms are harder to ionize
• Successive ionizations for electrons are
represented by Ii (i = 1,2,3,…)
I1 < I2 < I3 <…
Ionization energy and shells
• There is always a huge increase in Ii when i
represents an electron in an inner shell
Kink and ionization energy
• Unlike atomic radius
the trend is not as
smooth
• Note the dip at Be – B
and N – O
• Random or rational?
– B adds electron to 2p
– O pairs electron (2p4)
Electron affinity
• Electron affinity:
The electron affinity is the energy released
when an electron is added to a gaseous
neutral atom
X(g) + e = X(g)• Unlike ionization, addition of an electron is
usually favourable – even for
electropositive elements
Periodicity and electron affinity
• Smooth trends are
not in evidence
• Affinity tends to be
zero or positive when
electron is added to
new (Be, Mg, He, Ne
etc) or half-filled (N)
shell
Electronegativity and periodicity
Density and periodicity
Trend in melting point not as easily
explained
Periodicity and chemical properties
• Most obvious trend is from metallic to nonmetallic
• Acid-base properties of oxides:
– Metal oxides tend to be basic
• (Na2O = NaOH in H2O)
– Nonmetal oxides tend to be acidic
• (SO2 = H2SO3 in H2O)