Electron Configuration and Periodic
Properties

Atomic Radii


The size of an atom is defined by the
edge of its orbital
Since this boundary is fuzzy, the
radius is defined as one-half the
distance between the nuclei of
identical atoms that are bonded
together
Atoms tend to get smaller as you move
across a period due to the increased
positive charge
 They get larger as you move down a
group due to the increasing energy
levels occupied

Ionization Energy
Ionization energy is the energy required
to remove one electron from a neutral
atom
 Made on isolated atoms in the gas phase
 In general, ionization energies of the
main group elements (s&p) increase
across a period
 Generally decrease down a group


With sufficient energy, electrons can be
removed from positive ions as well as from
neutral atoms
 The energies are referred to as the second
ionization energy, third ionization energy, and
so on
 These energies generally increase due to the
stronger effective nuclear charge
 There are large jumps in energies when stable
arrangements are ionized (in particular- the
noble gas configurations)
Electron Affinity
The energy change that occurs when an
electron is acquired by a neutral atom is
called the atom’s electron affinity
 Atoms that release energy have a
negative affinity (they want the electron)
 Atoms that require energy to “force” the
electron on them have a positive affinity
(they will lose the electron
spontaneously)

The halogens gain electrons most readily
 The p group elements generally become
more negative as you move across a
period (again exceptions caused by
stable electron arrangements)
 The trends in groups are not as regular
(competing increased nuclear charge
and atomic radius)


Generally the size predominates

For an isolated ion in the gas phase, it is
always more difficult to add a second
electron to an already negatively
charged ion
Second affinities are therefore always
positive
 Ions like Cl-2 never occur

Ionic Radii

A positive ion is known as a cation
Caused by the loss of electrons
 The remaining electrons are drawn closer to
the nucleus by the unbalanced charge


A negative ion is known as an anion
Formed from the addition of extra electrons
 The electrons are not drawn as tightly as
they were before the addition

The metals on the left tend to form
cations, while the nonmetals on the
upper right tend to form anions
 Cationic radii decrease across a period
due to increasing nuclear charge
 Anionic radii (starting w/ group 15)
decrease across a period
 Ionic radii tend to increase down a group

Valence Electrons
Chemical compounds form because
electrons are lost, gained, or shared
between the outermost energy levels of
atoms (the inner electrons are too tightly
held
 These available electrons are called the
valence electrons
 For the main group elements these are
in the s & p shells

Electronegativity
Valence electrons hold atoms together
 In many compounds, the negative
charge of the valence electrons is
concentrated closer to one atom than to
another
 Electronegativity is a measure of the
ability of an atom in a chemical
compound to attract electrons
 Fluorine is assigned a number of 4.0

Electronegativities tend to increase
across each period
 Electronegativities tend to either
decrease down a group or remain about
the same
 Noble gases do not form many
compounds and may not have values

Properties of the d and f block
elements
The properties of the d block elements
vary less and with less regularity than
those of the main group elements
 Both the outer s and the d electrons are
available to interact with their
surroundings
 The atomic radii of the d block elements
generally decrease across a period

The d electrons shield the outer electrons
 The electrons repel each other

The f block elements behave in a similar
way
 Ionization energies generally increase
across a period for d & f block elements


In contrast, they generally increase down a
group because the electrons available for
ionization in the outer s level are less
shielded (incomplete d shell) from the
increasing nuclear charge

Ion formation in the d & f block elements
follows the reverse order of electron
configuration


For d block, although electrons are being
added to the d, they are removed from the
outer s first (most d block elements
therefore form +2 ions)
The d & f block elements all have similar
electronegativities

Follow general trend
Chemical Bonding
Introduction to Chemical
Bonding
A chemical bond is a mutual electrical
attraction between the nuclei and
valence electrons of different atoms that
binds atoms together
 Atoms bond because it decreases their
potential energy, creating more stable
arrangements of matter

Chemical bonding that results from the
electrical attraction between large
numbers of cations and anions is called
ionic bonding
 Covalent bonding results from the
sharing of electrons pairs between two
atoms


In a purely covalent bond, the shared
electrons are “owned” equally by the two
bonded atoms

Bonding is rarely purely ionic or covalent
 Electronegativity is a measure of an atom’s
ability to attract electrons
 The degree of ionic or covalent character is
determined by calculating the difference in
electronegativity
The d indicates
a partial charge
Covalent Bonding and Molecular
Compounds

A molecule is a neutral group of atoms
that are held together by covalent bonds
Individual unit capable of existing on its own
 May consist of two or more atoms


A chemical compound whose simplest
units are molecules is called a molecular
compound

A chemical formula indicates the relative
numbers of atoms of each kind in a chemical
compound by using atomic symbols and
numerical subscripts
 A diatomic molecule is a molecule containing
only two atoms
A balance is reached between
the attractive forces and the
repulsive forces between the
nuclei and electrons. This results
in the most energetically stable
arrangement.

In a covalent bond, the electrons orbitals can be
pictured as overlapping (the electrons are free to
move in either orbital)
 The distance between two bonded atoms at their
minimum potential energy is the bond length

The atoms will vibrate a bit

The difference between the potential
energy zero level (separate atoms) and
the bottom of the valley (bonded atoms)
is the bond energy that is released when
the bond is formed


It is also the energy required to break a
chemical bond and form neutral isolated
atoms
Atoms tend to acquire noble gas
configurations when bonding

Octet rule: Chemical compounds tend to form so
that each atom, by gaining, losing, or sharing
electrons, has an octet of electrons in its highest
occupied energy level

There are exceptions to the octet rule
Boron: In BF3 , boron will share its three
valence electrons and acquire a total of 6
 When some elements combine with the very
electronegative atoms of F, O, and Cl, an
expanded valence that involves electrons in
the d orbitals occurs


Electron-dot notation
is an electron
configuration notation
in which only the
valence electrons of
an atom of a
particular element are
shown, indicated by
dots placed around
the element’s symbol


Electron dot notations
can also be used to
represent molecules
A shared pair of electrons
is drawn between two
atoms, an unshared pair
is a pair of valence
electrons that belongs
exclusively to one atom
and is not involved in
bonding
H:H

A shared pair of electrons is often
represented with a dash
 The are called Lewis structures
 A structural formula indicates the kind,
number, arrangement, and bonds, but not the
unshared pairs of atoms in a molecule
A single bond is a covalent bond produced
by the sharing of one pair of electrons
between two atoms
 A double covalent bond is produced by the
sharing of two pairs of electrons between
two atoms
 A triple covalent bond is a bond produced
by the sharing of three pairs of electrons
between two atoms
 Double and triple bonds are referred to as
multiple bonds


C, N, and O can have multiple bonds
 H can have only one bond

Resonance structures cannot be correctly
represented by a single Lewis structure
 Ozone


Once thought to split time between two structures
Experiments show that bonds are equivalent (
average of two bonds)
Not all covalent compounds are
molecular
 Some are continuous 3 dimensional
networks of covalently bonded atoms


Called covalent-network bonding
Ionic Bonding and Ionic
Compounds
An ionic compound is composed of
positive and negative ions that are
combined so that the numbers of positive
and negative charges are equal
 Most are crystalline solids
 The formula simply represents the
simplest ratio of ions that give neutrality
of charge – called a formula unit



To compare
bond
strengths in
ionic
compounds,
chemists
compare
lattice
energies
Lattice
energy is the
energy
released
when one
mole of an
ionic
crystalline
solid
compound is
formed from
gaseous ions

The attraction between positive and
negative ions is generally very strong
 In molecular compounds, the
covalent bonds are also very strong,
but the intermolecular attractions are
generally much weaker than ionic
attractions
 The melting point, boiling point, and
hardness of a compound depend on
how strongly these basic units are
attracted to each other

Many molecular compounds melt at low
temperatures, while many ionic compounds
have high MP and BP
 Ionic compounds are brittle (a shift in layers can
cause a strong repulsive force)

As a solid, ions cannot move, so ionic
compounds are not conductors


In the molten or aqueous state, they are free to move
and are conductors
A charged group of covalently bonded atoms is
known as a polyatomic ion
Metallic Bonding
Chemical bonding is different in metals
than in ionic, molecular, or covalentnetwork compounds
 The valence electrons are highly mobile
 For most metals, the highest p orbitals
are vacant (and often some d orbitals as
well)
 In metals, these vacant orbitals overlap


The electrons then can roam freely
throughout the metal (delocalized)
The chemical bonding that results from the
attraction between metal atoms and the
surrounding sea of electrons is called
metallic bonding
 As a result, metals have high electrical
and thermal conductivity
 Since they have many orbitals separated
by extremely small energy differences,
they can absorb a wide range of
frequencies (and radiate them back) –
causes shiny appearance


Metals are malleable (hammered or beaten into
shapes) and ductile (drawn into a thin wire)


Caused by uniformity of bonding throughout the metal
Metallic bond strength varies with the nuclear
charge and the number of electrons in the metals
electron sea
 Reflected in the heat of vaporization
Molecular Geometry

Molecular properties depend not only on the
bonding of atoms but also on molecular
geometry
 The polarity of each bond, along with the
geometry of the molecule, determines the
molecular polarity
 There are two theories to explain geometry


VSEPR
hybridization
VSEPR Theory

Stands for valence-shell,
electron-pair repulsion
 States that repulsion between
the sets of valence-level
electrons surrounding an atom
causes these sets to be
oriented as far apart as possible
 Diatomic atoms are linear
The number of
bonds determines
the bond shapes

If the central atom has both shared and unshared
electrons, the unshared electrons must be
accounted for also

They also take up space around the atom


Double and
triple
bonds are
treated in
the same
way as
single
bonds
Polyatomic
ions are
treated in
the same
way as
molecules
Hybridization
VSEPR is useful for explaining the
shapes of molecules- but it does not
reveal the relationship between a
molecule’s geometry and the orbitals
occupied by its bonding electrons
 Hybridization is the mixing of two or
more atomic orbitals of similar energies
on the same atom to produce new
orbitals of equal energies


Methane (CH4) is a tetrahedron



How does carbon (outer shell 2s22p2) form four
equivalent bonds?
The 2s and three 2p orbitals hybridize to form 4
equivalent hybrid orbitals called sp3
The orbitals all have energy that is greater than the 2s but
less than the 2p

Hybrid orbitals are orbitals of equal energy
produced by the combination of two or more
orbitals on the same atom
 Explains many Group 15 & 16 elements

The linear geometry
of BeF2 can be
explained by the
hybridization of one s
and one p orbital
(called sp hybrid)
 BF3 is trigonal planar


Involves one s and two
p orbitals
Called sp2 hybrid
Intermolecular Forces

As a liquid is heated, the kinetic energy of
its particles increases
At the boiling point the energy is sufficient to
overcome the forces of attraction between the
liquid’s particles
 Boiling point is a good measure of the force of
attraction between particles of a liquid (higher
= stronger)


The forces of attraction between
molecules are known as intermolecular
forces
The strongest intermolecular forces exist
between polar molecules
 Polar molecules act as tiny dipoles
because of their uneven charge
distribution


A dipole is created by equal but opposite
charges that are separated by a short
distance

Dipole


The dipole direction is from the positive to the negative
side
Represented by an arrow with the head pointed toward
the negative pole and a crossed tail at the positive pole
The negative region in one polar molecule
attracts the positive region in adjacent
molecules
 The forces of attraction between polar
molecules are known as dipole-dipole
forces


Short range forces only

Shows in the difference between the boiling
point of BrF (-20oC) and that of F2 (-188oC)

For molecules containing more than two atoms,
molecular polarity depends on both the polarity
and the orientation of each bond

A polar molecule can induce a dipole in a nonpolar
molecule by temporarily attracting its electrons


This results in a short range intermolecular force that is
somewhat weaker than the dipole-dipole force
This accounts for the solubility of nonpolar O2 in water

Some hydrogen containing compounds
have unusually high boiling points
This is explained by the presence of a
particularly strong type of dipole-dipole force
 In compounds containing bonds between
hydrogen and fluorine, oxygen, or nitrogen the
large electronegativity difference makes them
highly polar
 The small size of the hydrogen allows it to
come very close to the unshared pair of
electrons on an adjacent molecule


The intermolecular force in which a hydrogen
atom that is bonded to a highly electronegative
atom is attracted to an unshared pair of
electrons of an electronegative atom in a
nearby molecule is known as hydrogen
bonding
Hydrogen bonds are usually represented
by dotted lines
 H2S boils at -61oC while water boils at
100oC



Since electrons are constantly moving, temporary
dipoles can be created
The intermolecular attractions resulting from the constant
motion of electrons and the creation of instantaneous
dipoles are called London dispersion forces

Fritz London proposed in 1930

London dispersion forces
Act between all atoms and molecules
 They are the only intermolecular forces
acting among noble gas atoms and
nonpolar molecules (low boiling points)
 Since they are dependent on electron
motion, they increase with the number of
electrons (increase with increasing molar
mass)
