Material Science and Engr. GROUP2 (2).pdf

ATOMIC STRUCTURE
The smallest
particle of a
chemical element.
An atom is usually
made of protons,
neutrons &
electrons.
Positively
charged centre
of an atom.
Holds most of an
atom's mass.
ATOM NUCLEUS PROTON NEUTRON ELECTRON
A subatomic
particle with a
positive charge.
Found in the
nucleus of an atom.
A subatomic
particle with a
neutral charge.
Found in the
nucleus of an
atom.
A subatomic
particle with a
negative charge.
Orbits around the
nucleus of an
atom.

CHARGE AND MASS
PROTONS
NEUTRONS
ELECTRONS
CHARGE MASS
(-)1.602 × 10^−19 C
(+)1.602 × 10^−19 C
NEUTRAL
1.67 × 10^−27 kg
1.67 × 10^−27 kg
9.11 × 10−31 kg

58.933
Co
Co
Cobalt
27
ISOTOPES
atomic # (Z)
no. of protons
average
atomic mass
27
59 60
Co
27 27
Co
61
mass #
atomic #

CALCULATIONS
PROTONS
NEUTRONS
ELECTRONS
FORMULA
atomic #
mass # - atomic #
atomic # - charge

23
Na
11
EXAMPLE
mass #
atomic #
PROTONS
NEUTRONS
ELECTRONS
FORMULA
atomic # = 11
mass # - atomic #
= 23 - 11
= 12
atomic # - charge
= 11 - 0
= 11

Each chemical element is characterized by the number of protons in the
nucleus, or the atomic number (Z).
For an electrically neutral or complete atom, the atomic number also
equals the number of electrons.
This atomic number ranges in integral units from 1 for hydrogen to 92 for
uranium.
The atomic mass (A) of a specific atom may be expressed as the sum of
the masses of protons and neutrons within the nucleus.
The number of protons is the same for all atoms of a given element.
NOTE

The atomic weight of an element corresponds to the weighted average
of the atomic masses of the atom’s naturally occurring isotopes.
A scale has been established whereby 1 amu is defined as 1 12 of the
atomic mass of the most common isotope of carbon, carbon 12 (12C)
(A = 12.00000).
Within this scheme, the masses of protons and neutrons are slightly
greater than unity,
A ≅Z + N
NOTE

The atomic weight of an element or the molecular weight of a compound
may be specified on the basis of amu per atom (molecule) or mass per
mole of material
In one mole of a substance, there are 6.022 × 1023 (Avogadro’s number)
atoms or molecules.
These two atomic weight schemes are related through the following
equation:
1 amu/atom (or molecule) = 1 g/mol
NOTE

HOW DID IT ALL HAPPEN?
QUANTUM
MECHANICS
(19TH Century)
Bohr atomic model Wave mechanical
model

During the latter part of the
nineteenth century, it was realized
that many phenomena involving
electrons in solids could not be
explained in terms of classical
mechanics.
Quantum Mechanics was
established
Quantum Mechanics is a set of
principles and laws that govern
systems of atomic and subatomic
entities
QUANTUM
MECHANICS
(19TH Century)

COMPARISONS ON BOHR AND
WAVE MECHANIC MODEL
Bohr atomic model was
proposed by Niels Bohr in 1915
to explain the structure of an
atom.
Describes the particle behavior
of an atom.
Does not give information about
quantum number.
Applicable to one-electron
systems
Wave mechanical model is
considered as the modern
atomic model to explain the
structure of an atom
accurately.
Describes the wave-particle
duality of an electron.
Explains quantum numbers.
Applicable to any atom

QUANTUM NUMBERS
Quantum no. Represents Value Meaning
Principal Quantum
number (n)
Energy level of
the electron
Any integer
value
Larger Value of n the
further away from the
nucleus
Second (azimuthal
quantum number
l)
Shape of
orbital
0 - (n-1)
example
n=3
l = 0, 1, 2
l = 0 (s orbitals)
l = 1 (p orbitals)
l = 2 (d orbitals)
l = 3 (f orbitals)

QUANTUM NUMBERS
Quantum no. Represents Value Meaning
magnetic quantum
number Energy level of
the electron
Determines the number
of electron orbitals for
each subshell
Orientation (
up or down) 1/2 or -1/2
1/2 = spin up
-1/2 = spin down
ml
-l to l
example
l=2
= -2, -1, 0, 1, 2
ml
spin moment
s
m

ORBITALS
l=0
ml = 0
thus only 1 s orbital
l=1
ml = -1, 0, 1
thus, 3 p orbitals
l=2
ml = -2,-1, 0, 1, 2
thus, 4 d orbitals
l=3
ml =-3, -2,-1, 0, 1,
2, 3
thus, 7 f orbitals

ELECTRON CONFIGURATION
ELECTRON STATES. Are values of energy that are permitted for
electrons
PAULI EXCLUSION PRINCIPLE. This stipulates that each electron
state can hold no more than two electrons that must have
opposite spins. Thus, s, p, d, and f subshells may each
accommodate, respectively, a total of 2, 6, 10, and 14 electrons.
GROUND STATE. State of an atom when all the electrons occupy
the lowest possible energies in accord with the foregoing
restrictions.

ELECTRON CONFIGURATION. Is distribution of electrons of an atom
or molecule in atomic or molecular orbitals. In the conventional
notation, the number of electrons in each subshell is indicated by a
superscript after the shell– subshell designation. for example, the
electron configurations for hydrogen, helium, and sodium are,
respectively, 1s 1 , 1s 2 , and 1s 2 2s 2 2p6 3s 1 .
VALENCE ELECTRONS. Are electrons that occupy the outermost
shell. These electrons are extremely important; as will be seen, they
participate in the bonding between atoms to form atomic and
molecular aggregates. Furthermore, many of the physical and
chemical properties of solids are based on these valence electrons.

STABLE ELECTRON CONFIGURATIONS. The states within the
outermost or valence electron shell are completely filled. Normally,
this corresponds to the occupation of just the s and p states for the
outermost shell by a total of eight electrons, as in neon, argon, and
krypton; one exception is helium, which contains only two 1s
electrons. These elements (Ne, Ar, Kr, and He) are the inert, or noble,
gases, which are virtually unreactive chemically.

THE PERIODIC
TABLE OF
ELEMENTS
Topic 2

A LITTLE BIT OF HISTORY
In 1869, Russian chemist Dmitri Mendeleev created the
framework that became the modern periodic table, leaving gaps
for elements that were yet to be discovered. While arranging the
elements according to their atomic weight, if he found that they
did not fit into the group he would rearrange them.
Mendeleev predicted the properties of some undiscovered
elements and gave them names such as "eka-aluminium" for an
element with properties similar to aluminium. Later eka-
aluminium was discovered as gallium. Some discrepancies
remained; the position of certain elements, such as iodine and
tellurium, could not be explained.

A LITTLE BIT OF HISTORY
Dmitri Ivanovich Mendeleev
(February 8, 1834 - February 2, 1907)

The numbers in parentheses are the atomic weights of the most stable or
common isotopes.
The elements positioned in Group 0, the rightmost group, are the inert
gases, which have filled electron shells and stable electron configurations.
Group VIIA and VIA elements are one and two electrons deficient,
respectively, from having stable structures.
The group VIIA elements (F, Cl, Br, I, and At) are sometimes termed the
halogens.
The alkali and the alkaline earth metals (Li, Na, K, Be, Mg, Ca, etc.) are
labeled as Groups IA and IIA, having, respectively, one and two electrons in
excess of stable structures.
NOTE

The elements in the three long periods, Groups IIIB through IIB, are termed
the transition metals, which have partially filled d electron states and in
some cases one or two electrons in the next higher energy shell.
Groups IIIA, IVA, and VA (B, Si, Ge, As, etc.) display characteristics that are
intermediate between the metals and nonmetals by virtue of their
valence electron structures.
NOTE

ELECTROPOSITIVE ELEMENTS. Elements that are capable of
giving up their few valence electrons to become positively
charged ions.
ELECTRONEGATIVE ELEMENTS. Elements situated on the right
side of the table they readily accept electrons to form
negatively charged ions, or sometimes they share electrons with
other atoms.
NOTE

Mechanically, the metallic elements exhibit varying degrees of
ductility—the ability to be plastically deformed without
fracturing (e.g., the ability to be rolled into thin sheets).
Most of the nonmetals are either gases or liquids, or in the solid
state are brittle in nature. Furthermore, for the Group IVA
elements [C (diamond), Si, Ge, Sn, and Pb], electrical conductivity
increases as we move down this column.
The Group VB metals (V, Nb, and Ta) have very high melting
temperatures, which increase in going down this column.
NOTE

BONDING ENERGY. This represents the energy
required to separate two atoms to an infinite
separation
NOTE

At small separation distances, each atom exerts forces on the
others. These forces are of two types, attractive (FA) and
repulsive (FR), and the magnitude of each depends on the
separation or interatomic distance (r).
The origin of an attractive force FA depends on the particular
type of bonding that exists between the two atoms.
Repulsive forces arise from interactions between the negatively
charged electron clouds for the two atoms and are important
only at small values of r.
NOTE

The net force FN between the two atoms is just the sum of both
attractive and repulsive components.
NOTE
When FA and FR are equal in magnitude but opposite in sign,
there is no net force and a state of equilibrium exists.

Sometimes it is more convenient to work with the potential
energies between two atoms instead of forces. Mathematically,
energy (E) and force (F) are related as
NOTE
May also be expressed as:

NOTE
May also be expressed as:
For Atomic Systems

Materials having large bonding energies typically also have high
melting temperatures.
At room temperature, solid substances are formed for large
bonding energies.
For small energies, the gaseous state is favored.
Liquids prevail when the energies are of intermediate
magnitude.
The mechanical stiffness (or modulus of elasticity) of a material
is dependent on the shape of its force–versus–interatomic
separation curve.
NOTE

Three different types of primary or chemical bond are found in
solids—ionic, covalent, and metallic.
For each type, the bonding necessarily involves the valence
electrons;
The nature of the bond depends on the electron structures of
the constituent atoms.
In general, each of these three types of bonding arises from
the tendency of the atoms to assume stable electron
structures, like those of the inert gases, by completely filling
the outermost electron shell.
NOTE

IONIC BONDING. It is always found in compounds composed of both
metallic and nonmetallic elements, elements situated at the horizontal
extremities of the periodic table.
Atoms of a metallic element easily give up their valence electrons to the
nonmetallic atoms.
In the process, all the atoms acquire stable or inert gas configurations
(i.e., completely filled orbital shells) and, in addition, an electrical charge
—that is, they become ions.
The attractive bonding forces are coulombic—that is, positive and
negative ions, by virtue of their net electrical charge, attract one
another.
NOTE

For two isolated ions, the attractive energy EA is a function of
the interatomic distance according to
NOTE

Here 𝜀0 is the permittivity of a vacuum (8.85 × 10−12 F/m), |Z1|
and |Z2| are absolute values of the valences for the two ion
types, and e is the electronic charge (1.602 × 10−19 C).
NOTE

Ionic bonding is termed nondirectional—that is, the
magnitude of the bond is equal in all directions around an
ion.
Ionic bonding is termed nondirectional—that is, the
magnitude of the bond is equal in all directions around an
ion.
Bonding energies, which generally range between 600 and
1500 kJ/mol, are relatively large, as reflected in high
melting temperatures.
NOTE

COVALENT BONDING. It is found in materials whose atoms have small
differences in electronegativity—that is, that lie near one another in the
periodic table.
For these materials, stable electron configurations are assumed by the
sharing of electrons between adjacent atoms
Two covalently bonded atoms will each contribute at least one electron
to the bond, and the shared electrons may be considered to belong to
both atoms
Covalent bond is directional—that is, it is between specific atoms and
may exist only in the direction between one atom and another that
participates in the electron sharing.
NOTE

Covalent bonds may be very strong, as in diamond, which is very
hard and has a very high melting temperature, 3550°C (6400°F),
or they may be very weak, as with bismuth, which melts at about
270°C (518°F).
Most covalently bonded materials are electrical insulators, or, in
some cases, semiconductors.
Mechanical behaviors of these materials vary widely.
It is difficult to predict the mechanical properties of covalently
bonded materials on the basis of their bonding characteristics.
NOTE

BOND HYBRIDIZATION IN CARBON
Often associated with the covalent bonding of carbon (as well
other nonmetallic substances) is the phenomenon of
hybridization—the mixing (or combining) of two or more atomic
orbitals with the result that more orbital overlap during bonding
results.
Bonding hybrid orbitals are directional in nature—that is, each
extends to and overlaps the orbital of an adjacent bonding atom.
Furthermore, for carbon, each of its four sp3 hybrid orbitals is
directed symmetrically from a carbon atom to the vertex of a
tetrahedron. The angle between each set of adjacent bonds is
109.5°.

For diamond, its carbon atoms are bonded to one another with
sp3 covalent hybrids—each atom is bonded to four other
carbon atoms.
Diamond’s carbon–carbon bonds are extremely strong, which
accounts for its high melting temperature and ultrahigh
hardness (it is the hardest of all materials).
Many polymeric materials are composed of long chains of
carbon atoms that are also bonded together using sp3
tetrahedral bonds; these chains form a zigzag structure
because of this 109.5° interbonding angle.

Other types of hybrid bonds are possible for carbon, as
well as other substances. One of these is sp2 , in which an s
orbital and two p orbitals are hybridized.
To achieve this configuration, one 2s orbital mixes with two
of the three 2p orbitals—the third p orbital remains
unhybridized.

METALLIC BONDING. the final primary bonding type, is found in metals
and their alloys.
NOTE

With this model, these valence electrons are not bound to any
particular atom in the solid and are more or less free to drift
throughout the entire metal.
With this model, these valence electrons are not bound to any
particular atom in the solid and are more or less free to drift
throughout the entire metal.
The remaining nonvalence electrons and atomic nuclei form what
are called ion cores, which possess a net positive charge equal in
magnitude to the total valence electron charge per atom.
NOTE

The free electrons shield the positively charged ion cores from
the mutually repulsive electrostatic forces that they would
otherwise exert upon one another;
Metallic bond is nondirectional in character.
In addition, these free electrons act as a “glue” to hold the ion
cores together.
Metallic bonding is found in the periodic table for Group IA and
IIA elements and, in fact, for all elemental metals.
NOTE

Metals are good conductors of both electricity and heat as a
consequence of their free electrons.
At room temperature, most metals and their alloys fail in a
ductile manner—that is, fracture occurs after the materials
have experienced significant degrees of permanent
deformation. This is implicitly related to the characteristics of
the metallic bond
NOTE

SECONDARY
BONDING OR VAN
DER WAALS
BONDING

These are weak in comparison to the primary or chemical bonds; bonding
energies range between about 4 and 30 kJ/mol.
Secondary bonding exists between virtually all atoms or molecules, but
its presence may be obscured if any of the three primary bonding types
is present.
Secondary bonding forces arise from atomic or molecular dipoles.
SECONDARY BONDS OR VAN DER WAALS BONDING
NOTE

The bonding results from the coulombic attraction between the positive
end of one dipole and the negative region of an adjacent one.
NOTE

FLUCTUATING INDUCED DIPOLE
BONDS
A dipole may be created or induced in an atom or molecule that is
normally electrically symmetric—that is, the overall spatial
distribution of the electrons is symmetric with respect to the
positively charged nucleus.
All atoms experience constant vibrational motion that can cause
instantaneous and short-lived distortions of this electrical symmetry
for some of the atoms or molecules and the creation of small electric
dipoles.

BONDS
One of these dipoles can in turn produce a displacement of the
electron distribution of an adjacent molecule or atom, which induces
the second one also to become a dipole that is then weakly attracted
or bonded to the first; this is one type of van der Waals bonding.

BONDS
These attractive forces, which are temporary and fluctuate
with time, may exist between large numbers of atoms or
molecules.
The liquefaction and, in some cases, the solidification of the
inert gases and other electrically neutral and symmetric
molecules such as H2 and Cl2 are realized because of this
type of bonding.

POLAR MOLECULE–INDUCED DIPOLE
BONDS
Permanent dipole moments exist in some molecules by virtue
of an asymmetrical arrangement of positively and negatively
charged regions; such molecules are termed polar molecules.
Polar molecules can also induce dipoles in adjacent nonpolar
molecules, and a bond forms as a result of attractive forces
between the two molecules.

POLAR MOLECULE–INDUCED DIPOLE
BONDS
The magnitude of this bond is greater than for fluctuating
induced dipoles.

PERMANENT DIPOLE BONDS
Permanent dipole describes the partial charge separation that
can occur within a molecule along the bond that forms
between two different atoms.
Dipoles generally occur between two nonmetals that share
electrons as part of their bond.
For example, a molecule of hydrogen chloride, HCl has a large
permanent dipole. This is because there is a difference in
electronegativity between H and Cl: Chlorine is more
electronegative than hydrogen and pulls the electrons in the
covalent bond towards itself.

PERMANENT DIPOLE BONDS
The strongest secondary bonding type, the hydrogen bond, is
a special case of polar molecule bonding.

Govern physical phenomena
solubility, surface tension, vapor pressure, volatility, viscosity, etc.
Applications in MSE:
adhesives
surfactants
emulsifiers
desiccants
NOTE

Sometimes it is illustrative to represent the four bonding types—ionic,
covalent, metallic, and van der Waals—on what is called a bonding
tetrahedron—a three-dimensional tetrahedron with one of these
“extreme” types located at each vertex.
NOTE

For many real materials, the atomic bonds are mixtures of two or more of
these extremes.
Three mixed bond types—covalent–ionic, covalent–metallic, and
metallic–ionic—are also included on edges of this tetrahedron.
For mixed covalent–ionic bonds, there is some ionic character to most
covalent bonds and some covalent character to ionic ones. As such,
there is a continuum between these two extreme bond types.
This type of bond is represented between the ionic and covalent bonding
vertices.
The degree of either bond type depends on the relative positions of the
constituent atoms in the periodic table or the difference in their
electronegativities
NOTE

The wider the separation (both horizontally—relative to Group IVA—and
vertically) from the lower left to the upper right corner (i.e., the greater
the difference in electronegativity), the more ionic is the bond.
Conversely, the closer the atoms are together (i.e., the smaller the
difference in electronegativity), the greater is the degree of covalency.
Percent ionic character (%IC) of a bond between elements A and B (A
being the most electronegative) may be approximated by the expression:
where XA and XB are the electronegativities for the respective elements.
NOTE

Another type of mixed bond is found for some elements in Groups IIIA,
IVA, and VA of the periodic table (viz., B, Si, Ge, As, Sb, Te, Po, and At).
Interatomic bonds for these elements are mixtures of metallic and
covalent.
These materials are called the metalloids or semi-metals, and their
properties are intermediate between the metals and nonmetals.
In addition, for Group IV elements, there is a gradual transition from
covalent to metallic bonding as one moves vertically down this column
—for example, bonding in carbon (diamond) is purely covalent,
whereas for tin and lead, bonding is predominantly metallic.
NOTE

Mixed metallic–ionic bonds are observed for compounds
composed of two metals when there is a significant difference
between their electronegativities.
The larger this electronegativity difference, the greater the
degree of ionicity.
NOTE

Many common molecules are composed of groups of atoms
bound together by strong covalent bonds, including
elemental diatomic molecules (F2, O2, H2, etc.), as well as a
host of compounds (H2O, CO2, HNO3, C6H6, CH4, etc.).
In the condensed liquid and solid states, bonds between
molecules are weak secondary ones.
Molecular materials have relatively low melting and boiling
temperatures.
NOTE

Most materials that have small molecules composed of a few
atoms are gases at ordinary, or ambient, temperatures and
pressures.
However, many modern polymers, being molecular materials
composed of extremely large molecules, exist as solids; some
of their properties are strongly dependent on the presence of
van der Waals and hydrogen secondary bonds.
NOTE

BONDING TYPE-
MATERIAL
CLASSIFICATION
CORRELATIONS

Some correlations have been drawn between bonding type
and material classification.
Ionic bonding is associated with ceramics, covalent bonding is
associated with polymers, metallic bonding is associated with
metals, and van der Waals bonding is associated with
molecular solids.
NOTE

Material Science and Engr. GROUP2 (2).pdf

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