Windsor University
School of Medicine
STRUCTURE & BONDING
A SOCIETY GROWS GREAT WHEN OLD MEN PLANT TREES IN WHOSE SHADE
THEY KNOW THEY SHALL NEVER SIT. -- GREEK PROVERB
Ch.5
J.C. Rowe
LEARNING OBJECTIVES

Bohr’s Theory of atomic structure is the beginning of
quantum theory and describes electrons as residing in
very particular "quantized" energy levels within the
atom.
Shells, Sub-shells and Orbitals are the backbone of the
organizational structure of electrons in the atom
 Electron Configurations use shells, sub-shells and orbitals to
describe the relative electronic organization within the atom.


Periodicity
of electron configurations relate to periodicity of physical
and chemical properties.
 of valence shell configurations relate directly to relative
reactivity.

Bohr's Theory of the Atom and its structure
Bohr found that:
 electrons move in particular 'orbits' within the atom-meaning electrons have particular places or levels
within the atom,
 electrons have certain allowed energy values,
 the closer the electron is to the nucleus, the lower the
energy
 when electrons are excited, they gain (or absorb)
energy and move to a higher level (farther away from
the nucleus),
 when electrons give off light, they emit energy and
move to a lower level (closer to the nucleus),
SHELLS, SUBSHELLS, ORBITALS
Visualize the atom as an onion with different levels of structure and sub-structure,
all of which make up electron energy levels – each major energy level is
designated by a number (called the quantum number):
Major Shell



The major shell is the major layer of the onion--the
large, thick peels which make you cry when you chop
them.
The major shells are designated by numbers which have
integral values of 1, 2, 3, etc.
These numbers are designated as the major shell
quantum numbers, n.
n = 1 designates the first major shell, the one closest to the
nucleus and the one with the lowest (most negative) energy;
 n = 2 designates the second major shell, a little further from
the nucleus;
 n = 3 designates the third major shell, etc.

Major Shell Cont’d.

The major shell is the primary influence on the
energy of the electron.
 The
electron "assumes" the energy of the major shell in
which it resides.
 The closer the (negative) electron is to the (positive)
nucleus, the stronger the attraction will be, and the
harder the electron is to remove. So an electron in shell
n=1 will be very hard to remove while an electron in
shell n=4 would be much easier to remove.
Major Shell Cont’d.

Major shells contain subshells which are a
substructure within the major shell. The number of
the subshells is the same as the quantum number of
the major shell,
 For
example, major shell n = 1 has one subshell
(specifically, an 's' subshell); n = 2 has two subshells (an
's' and a 'p' subshell), n = 3 has three subshells (s, p,
and d subshells), and n = 4 has four subshells (s, p, d
and f).
Subshell


The subshells are a substructure within the major shells-imagine each onion skin or peel containing another set
of peels within each major peel
The subshells are designated by letters s, p, d, f (these
letters are terms which originated from the field of
spectroscopy).
s, p, d, and f are different types of subshells which have
different properties which we will learn about soon.
 These subshells contain a further substructure, which are sets
of orbitals which have the same energy.

Orbitals

Each subshell is further structured into orbitals.
s subshells have one orbital
p subshells have three, equal energy orbitals
d subshells have five, equal energy orbitals
f subshells have seven, equal energy orbitals

each orbital can accommodate a maximum of two
electrons, and
Orbitals Cont’d

each orbital takes on the characteristics of its
subshell and major shell.
s
subshell orbitals (there is only one per major shell) are
spherical,
 p subshell orbitals (there are three of them per major
shell) are dumbell shaped and directed along the x, y
and z axis,
 d subshell orbitals (there are five of them per major
shell) have a different shape [I only want you to
recognize s , p and d shapes]
the relative energy levels and structure of the
hydrogen atom
What is the maximum number of electrons that can reside in the first major shell? In
the diagram, the major shells are designated by the black lines coming from the
energy arrow. The different types of subshells are indicated by the differently
colored boxes. The orbitals themselves are the individual boxes. Remember, each
box (orbital) can contain only a maximum of two electrons.
Energies of the subshells

Although the major shell primarily determines the
energy of an electron, there are minor differences
between the energies of the subshells within any
major shell, particularly in an atom with more than
one electron (actually, any atom other than
hydrogen):
s
subshells are the lowest energy subshell within a shell
 p subshells are next lowest
 d subshells are higher than p subshells
 f subshells are higher than d subshells
Orbital/Subshell energy levels
Overlaps of subshells
Once we have many electrons (which means many shells and subshells) in an
atom, we begin to see some overlap of subshells--particularly, the s subshell from
the fourth level (4s) overlaps with the d subshell from the third level (3d) such as
shown above.
Shapes of the subshells:



For many reasons, we cannot define exactly where
an electron is within an atom.
We must, instead, define the volume which will most
likely contain electron density.
Consequently, the shapes of those volumes which
can contain electron density are somewhat "fuzzy".
High electron density vs. lower electron density.
We can depict volumes of high electron density (many dots) or volumes of low
electron density (few dots), as shown in the diagrams below.
Orbitals
An orbital is a region of high probability of
finding an electron.
 There are different types of orbital :
 an (s) orbital with a spherical shape
 a (p) orbital with 2 egg-shaped lobes
 a (d) orbital with 4 egg-shaped lobes
 An (f) orbital with 8 lobes & more complex.

Shape of (s) orbital
Shape of (p) orbital
Shape of (d) Orbital
Types of Orbitals & Contents
(s) orbital
(p) orbital (d) orbital
# orbitals
1
3
5
# electrons
(maximum)
2
6
10
Subshells & shells
Subshells grouped in shells.
 Each shell has a number called the
principal quantum number, n
 The principal quantum of the shell is
the number of the period or row of
the periodic table where that shell
begins

Shells & subshells cont’d
The principal quantum number(n) tells you
how far an electron is from the nucleus
 The secondary quantum number (l) tells
you what subshell (type of orbital) an
electron is in.
 The magnetic quantum number (m or ml)
tells you which orbital an electron is within
a given subshell.

Quantum Numbers

Each orbital is a function of 3 quantum
numbers
n
(major)……………….shell
 l (angular)………………subshell
 m.l (magnetic)……………orientation of
electron within a subshell
Quantum numbers cont’d
symbol
values
description
n (major)
1, 2, 3,………
Orbital size and
Energy
l (angular)
0,1,2,……n-1
Orbital shape or
type
m.l (magnetic) -l,….,0,….+l
Orbital orientation
Electron Configurations
Pattern:
2
2
6
2
6
2
10 4p6 5s2 4d10 5p6 6s2 5d10
 1s 2s 2p 3s 3p 4s 3d
6p6
 lower energy higher energy
 strongest held e less tightly held e
 the letters refer to the type of subshell s, p, d or f
 the coefficients (numbers) to the left of each letter
represent which major shell 1, 2, 3, etc.
 the subscript numbers represent the number of
electrons present in the subshell.
Pauli Principle & Hund’s Rule


Pauli Exclusion Principle: an orbital can contain no
more than two electrons and those two electrons
must be paired, in other words, they must have
opposite spin (usually indicated by one up-arrow
and one down-arrow, as we will see shortly).
Hund’s Rule: in a subshell where there are multiple
orbitals with the same energy, electrons will enter
then each orbital singly until all orbitals are half
filled before pairing with other electrons in the
subshell.
Arrangement of electrons in atoms




Each orbital can be assigned no more than 2
electrons.
This is tied to the existence of a 4th quantum number,
the electron spin quantum number, m.s
Pauli Exclusion Principle states that No two
electrons in the same atom can have the same set
of 4 quantum numbers.
That is, each electron has a unique address.
Electron filling order (Aufbau order)
1
2
3
4
5
6
7
8
Aufbau order
Electrons in Atoms

When n=1, then l =n-1=1-1=0



This shell has a single orbital (1s) to which 2 electrons
can be assigned.
When n=2, then l= n-1= 2-1=1….. 0, 1
This shell has
2s orbital with 2 electrons
 Three 2p orbital with 6 electrons


Thus a total of 8 electrons
Electrons in atoms cont’d.

When n=3, then l = n-1=3-1=2... 0, 1 ,2
 3s
orbital with 2 electrons
 Three 3p orbital with 6 electrons
 Five 3d orbital with 10 electrons

Thus a total of 18 electrons
Electrons in Atoms cont’d.

When n=4, then l= 0, 1, 2, 3
4s
orbital with 2 electrons
Three 4p orbitals with 6 electrons
Five 4d orbitals with 10 electrons
Seven 4f orbital with 14 electrons

Thus a total of 32 electrons
Arrangement of electrons in atoms

Electrons in atoms are arranged as
Shells (n)
Subshells (l)
Orbitals (m.l)
Periodic table vs. Aufbau order
Valence electrons & shells

Valence Electrons and Valence Shells:
 the
most important electrons in an atom are those on
the outer-most edge of the atom (the outer major
shell) – they come into contact with other atoms first and
participate fully in reactions – they are the valence
electrons
 valence shells contain the valence electrons
Periodic Table Pattern




s-block Group IA and IIA
valence electrons are in the ns subshell where 'n'
denotes the major shell or row
p-block Group IIIA - VIIIA
valence electrons are in the ns and np subshell where
'n' denotes the major shell or row
d-block Group IB - VIIIB (Transition Metals)
valence electrons are in the ns and (n-1)d subshell
f-block Inner Transition Metals sometimes called the
lanthanides & the actinides
valence electrons are in the ns and (n-2)d subshell
Electronic Configuration
Bonding
Ionic vs. covalent bonding.
Ionic bonds.
Covalent bonds.
Hydrogen bonds vs. Van der Waals’s forces
Electronegativity
Allotropy.
Ionic vs. Covalent bonds


Ionic bonding results from the transfer of
electrons from a metal to a nonmetal forming
a positively charged metallic ion and a
nonmetallic ion, which are then held together
by ionic bond.
Covalent bonding results from the sharing of
electrons btwn 2 nonmetals, btwn a nonmetal
& a metalloids
Ionic bonds
Ionic bond is a
complete
transfer of one
or more
electrons from
one atom to
another
Covalent bond
A covalent bond is a bond btwn
two atoms in which the electrons
are shared.
1.
Single covalent bond
2.
(one pair of electrons shared.)
3.
Doube covalent bond
4.
(two pairs of electrons shared)
5.
Triple covalent bond
6.
(three pairs of electrons
shared)
Forms of chemical bonds
Ionic bonds

Any compound
made out of a
metal & a nonmetal
is always an ionic
compound and is
held by ionic bonds.
Covalent bonds

1.
2.
3.
Normally forms btwn
A nonmetal & a
nonmetal
A nonmetal & a
metalloid
A metalloid & a
metalloid
Ionic vs. Covalent Compounds
ionic





Crystalline Solid
Conduct electricity when
molten
Variable H2O soluble
High Melting pt, boiling
pt, heat of fusion, heat
of vaporization
Tend to be very reactive
Covalent





Variable state
Do not conduct electricity
when molten
Like dissolve like
Low melting, boiling pt &
low heat of vaporization,
low heat of fusion
Not as reactive
Hydrogen bonds vs. Van der Waals’forces
Hydrogen bonds

Whenever a
hydrogen atom is
sandwiched in space
between 2 oxygen
atoms, hydrogen
bond is made to
hold the 3 atoms
together.
Van der Waals’ forces
Exist btwn atoms
& molecules but
this bond is neither
a covalent nor
ionic bond.
 It is a weak bond

Electronegativity

Electronegativity is the ability of an atom
in a molecule to attract shared electrons
to itself.
 Metals
have a low electronegativity.
 Nonmetals have a high electronegativity.

Fluorine is the most electronegative
element; it pulls very strongly on the
shared electrons in a covalent bond.
Electronegativity Cont’d.


Excluding the noble gases, elements that are closer
to fluorine have a higher electronegativity than
elements that are further away from fluorine.
The greater the difference in electronegativity btwn
two elements that are bonded to each other, the
more polar that bond is.
If the electrons are shared equally, the bond is a
nonpolar covalent bond.
 If the electrons are not shared equally, then the bond
is a polar covalent (meaning that the electrons spend
more of their time around one atom than the other).

Allotropy
Elements whose atoms bond together in more
than one way exist as allotropes.
 Because the properties of a substance are
partly set by its structure, the allotropes of an
element usually have different appearances,
melting point and boiling points, and electrical
properties.
 E.g. Carbon has 3 allotropes : diamond,
graphite and fullerene.

Diamond vs. Graphite
Diamond vs Graphite
Allotropes
Fullerene allotrope



Any carbon constituted
molecule in the form of
a hollow sphere,
ellipsoid, tube etc
shape.
Detected in nature &
outer-space
Important for the future
of nanotechnology &
electronics + cancer +
HIV research
André Gide
One doesn't discover new lands without
consenting to lose sight of the shore for a very
long time.