Periodicity of Atomic Properties
Elements in the same group have the same number of
valence electrons and related electron configurations;
hence have similar chemical properties.
The ground state electron configuration of the elements vary
periodically with atomic number; all properties that
depend on electron configuration tend to vary periodically
with atomic number
Valence electrons - involved in bonding, determine reactivity
Bi (Atomic number 83): [Xe] 4f14 5d10 6s2 6p3
6 valence electrons
Os (Atomic number 76): [Xe] 4f145d6 6s2
8 valence electrons
The variation of effective nuclear charge through the periodic
table plays an important role in determining periodic trends
Zeff increases from left to right across a period, but drops
every time an outer electron occupies a new shell.
Atomic Radius
Electron clouds do not have sharp boundaries, so do not
really define the radius of an atom.
Atomic radius - half the distance between the nuclei of
neighboring atoms
In solids centers of atoms are found at definite distances
from one another.
For metals: atomic radius is half the distance between the
centers of the atoms in the solid.
For nonmetals the atomic radius is half the distance between
the nuclei of atoms joined by a chemical bond; the
covalent radius.
van der Waals radius (used for noble gases): half the distance
between the centers of the atoms in the solidified gas
Trend: Atomic radius generally decreases from left to right
across a period and increases down a group
Across a period: new electron added to the same shell, and
nuclear charge is increasing
Ionic radius: effective radius of an ion in an ionic solid
Cations are smaller than the neutral atom; anions are larger
Ionic radii generally
increase down a group and
decrease from left to right
across a period.
Ionization Energies
Ionization energy: energy needed to remove an electron from
an atom in the gas phase.
First ionization energy (I1): energy required to remove an
electron from the neutral gas phase atom
Cu(g) -> Cu+(g) + eenergy = I1
For Cu I1 = 785 kJ/mol
Second ionization energy (I2): energy required to remove an
electron from the singly charged gas phase atom
Cu+(g) -> Cu2+ (g) + eenergy = I2
For Cu I2 = 1955 kJ/mol
Trend: Ionization energy generally decreases down a group,
and tends to increase moving across a period from left to
right
Departures from these trends can usually be traced to
repulsion between electrons, particularly electrons
occupying the same orbitals
Elements with low ionization energies can be expected to
form cations readily and to conduct electricity in their solid
form.
Elements in the lower left of the periodic table tend to have
lower ionization energies than those in the upper right.
These are the elements in the s block, d block, f block and the
lower left of the p block - metallic solids
Electron Affinity
Electron attachment energy: energy change when a gas
phase atom in its ground state gains a single electron
X(g) + e- -> X- (g)
DE = electron attachment energy
Electron affinity = - DE(electron attachment)
Electron affinity higher in
the upper right
Main Group Elements (s and p)
An s-block element has low ionization energy; outer
electrons can easily be lost
Group 1 form +1 ions; Group 2 form +2 ions
An s block element is likely to be a reactive metal
Since ionization energies are lowest at the bottom of each
group , these elements are most reactive in the s-block
p- block
Elements on the left have low enough ionization energies to
be metallic, but higher than the s-block elements and so
are less reactive
Elements on the right have high electron affinities (tend to
gain electrons to form closed shell ions)
Transition Metals (d block)
All d-block elements are metals; properties are transitional to
s and p block
Many d-block elements exhibit more than one oxidation state
since the d-electrons have similar energies (inner shell)
and a variable number can be lost
Chemical Bonds
Chemical Bond: arise from sharing or transfer of electrons
between two or more atoms
Covalent bond: when electron density is mainly shared
between two atoms
Ionic bond: electron density mainly transferred from one
atom to another.
Potential energy (kJ/mol)
Electronegativity: ability of an atom to attract electrons
shared in a bond toward itself
Elements to the lower left have low ionization energies and
low electron affinities; tend to act as electron donors
Elements to the upper right have high ionization energies and
high electron affinities; tend to act as electron acceptors
R. Mulliken
c = 0.5 (IE1 + EA)
L. Pauling: quantitative measure of electronegativity, c
| cA - cB| = 0.102{D(A-B) - 0.5 [ D(A-A) + D(B-B)]}0.5
A polar colavent bond is
a bond with partial
electric charges arising
from their difference in
electronegativity.
Pauling Scale for c
The partial charges give
rise to an electric dipole
moment
Bond dissociation energy DEd : energy required to separate
the bonded atoms
H-Cl(g) -> H(g) + Cl(g)
Greater the DEd, stronger the bond
Bond enthalpy: enthalpy change (heat absorbed at constant
pressure) when bond is broken
DEd = DHd - R T
Since bond enthalpy of a bond X-Y between two atoms X and
Y does not vary much between compounds, define an
“average” bond enthalpy or “average” bond energy
Variations in atomic nuclei contribute to trends in bond
energy
Number of bonds between two atoms also influence bond
energy
single < double < triple
Bond length: distance between the centers of two atoms
joined by a covalent bond
Bonds between larger atoms tend to be longer
Multiple bonds are shorter than single bonds
A bond length is approximately the sum of the covalent radii
of the two atoms
Covalent radius: effective radius of an atom in a covalent
bond
Bond Order: number of shared electron pairs
Single bond: bond order = 1
Double bond: bond order = 2
Triple bond: bond order = 3
Bond length decrease with bond order
Bond strength increases with bond order
Chemical Bonding
The Lewis model of the chemical bond assumes that each
bonding electron is located between two bonded atoms localized electron model.
However, location of an electron in an atom cannot be
described in terms of a precise position; same is true of
electrons in a molecule.
Valence Bond Theory - first quantum mechanical description
of distribution of electrons in bonds.
Valence Bond Theory
Minimum requirement for the formation of a covalent bond is
two electrons and two valence orbitals that can overlap.
The overlapping of the two orbitals concentrate the electron
probability between the two nuclei, creating a chemical
bond.
H2: two H atoms, each in the ground state 1s1
In VB theory: the two H atoms come together and their atomic
orbitals merge or overlap
Greater the extent of overlap, greater the strength of the bond
Overlap of the two 1s orbitals form a sigma (s) bond.
In a sigma bond - electron density is distributed along the
bond axis
Note that the electron spin is paired
HF
H: 1s1
F: 1s22s22p5
Overlap between the valence orbital of H (1s) and valence
orbital of F (2p) to form a s bonds
Note: electron spin is paired in the s orbital
By definition: z is the direction along the internuclear axis