Transcript ppt

Chapter 2:
Atomic Structure and Interactive Bonding
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Why Study Atomic Structure and Interactive
Bonding?
An important reason to have an understanding of
inter-atomic bonding in solids is that, in some
instances, the type of bond allows us to explain a
material’s properties.
For example, consider carbon, which may exist as
both graphite and diamond. Whereas graphite is
relatively soft and has a greasy feel to it,
diamond is the hardest known material. This
dramatic disparity in properties is directly
attributable to a type of interatomic bonding
found in graphite that does not exist in diamond.
2.4 Periodic Table
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All the elements have been classified
according to electron configuration.
Elements are situated with increasing
atomic number.
Seven horizontal rows are called
periods.
All elements that are arrayed in a
given column or group have similar
valence electron structures, as well
as chemical and physical properties.
2.4 Periodic Table (Contd.)
2.4 Periodic Table (Contd.)
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Electropositive elements  capable of giving up
their few valence electrons to become positively
charged ions.
Electronegative elements  readily accept
electrons to form negatively charged ions, or
sometimes they share electrons with other
atoms.
Figure 2.7 displays electronegativity values
assigned to various elements.
As a general rule, electronegativity increases in
moving from left to right and from bottom to top.
2.4 Periodic Table (Contd.)
Ch. 2: ATOMIC BONDING IN SOLIDS
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While there are only 118 or so elements listed on the
periodic table, there are obviously more substances in
nature than 118 pure elements. This is because atoms can
react with one another to form new substances called
compounds. When two or more atoms chemically bond
together, the resulting compound is unique both chemically
and physically from its parent atoms.
Let's look at an example.
The element sodium is a silver-colored metal that reacts so
violently with water that flames are produced when sodium
gets wet. The element chlorine is a greenish-colored gas
that is so poisonous that it was used as a weapon in World
War I. When chemically bonded together, these two
dangerous substances form the compound sodium chloride,
a compound so safe that we eat it every day - common
table salt!
NaCl
+
Bonding Forces and Energies
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Consider two isolated atoms:
• When the atoms are at large inter-atomic
separation distance, the atoms do not exert any force on each
other.
• When the distance is decreased, an attractive force FA starts to
act pulling atoms closer.
• FA increases as the atoms gets closer.
• But as the atoms get closer a repulsive force FR begin to act.
• The net force FN between the two atoms is given by:
FN = FA + FR
• At some inter-atomic distance ro, FR exactly equals FA and FN
becomes Zero
FN = 0 = FA + FR
• ro is called the equilibrium inter-atomic separation distance at
which atoms enter into bonding
ro ≈ 0.3 nm
Force vs. Separation Distance
Energy vs. Separation Distance
2.5 Bonding Forces and Energies (Contd.)
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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
E =  F dr
For atomic systems,
r
E N   FN dr
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r
r
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  FA dr   FR dr
 EA  ER
EN, EA, and ER : the net, attractive, and repulsive
energies for two isolated and adjacent atoms.
2.5 Bonding Forces and Energies (Contd.)
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Figure 2.8b plots potential energies.
The net curve, which is sum of attractive and
repulsive energies, has a potential energy well
around its minimum.
E0  Bonding energy at minimum point, energy
required to separate these two atoms to an
infinite separation.
A similar yet complex condition exists for solid
materials because force and energy interactions
among many atoms must be considered.
Nevertheless, a bonding energy, analogous to E0
above, may be associated with each atom.
2.5 Bonding Forces and Energies (Contd.)
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A number of material properties depend on E0, the curve
shape, and bonding type. For example,
• Materials having large bonding energies typically also
have high melting temperature.
• At room temperature, solids formed for large bonding
energies, whereas for small energies the gaseous state is
favored, liquids prevail when energies are of
intermediate magnitude.
• The mechanical stiffness (modulus of elasticity) is
dependent on the shape of the force-versus-interatomic
separation curve.
• The coefficient of thermal expansion (how much expands
or contracts per degree change in temperature) is related
to the shape of its E0-versus-r0 curve.
Types of Atomic & Molecular Bonds
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Primary Atomic Bonds
• Ionic Bonds
• Covalent Bonds
• Metallic Bonds
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Secondary Atomic & Molecular Bonds
• Permanent Dipole Bonds
• Fluctuating Dipole Bonds
Ionic Bonding
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Large inter-atomic forces are created
by the “coulombic” effect produced
by positively and negatively charged
ions.
Ionic bonds are “non-directional”.
The “cation” has a + charge & the
“anion” has the - charge.
The cation is much smaller than the
anion.
IONIC BONDING IN NaCl
NaCl
IONIC BONDING
The attractive energy EA is a function of inter-atomic
separation distance.
EA = -A/r
Where
A = (z1e)(z2e)/40
z1 and z2 are the valance of the two ions.
e is the electron charge (1.6 x 10-19 C)
o is the permittivity of vacuum (8.85 x 10-12 F/m)
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Repulsive energy for two isolated ions: ER = B / (rn)
A, B, and n are constants and depends on ionic
system, n is approximately 8.
Energy and Force relationship
Energy is related to force as:
E = ∫F dr
and thus
F = dE/dr
Useful Equations
EA = -A/r
FA = dE/dr = -A [d/dr(1/r)]
FA = A/r2
Since A = (z1e)(z2e)/4o
or
FA = (z1e)(z2e)/4or2
r=
√ (z1e)(z2e)/4oFA
Covalent Bonding
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Large inter-atomic forces are created
by the sharing of electrons to form
directional bonds.
The atoms have small differences in
electronegativity & close to each
other in the periodic table.
The atoms share their outer s and p
electrons so that each atom attains
the noble-gas electron configuration.
H2 Molecule
Primary Interatomic Bonding (Contd.)
Examples of covalent bonding:
Nonmetallic elemental molecules (H2, Cl2, F2, etc.)
Dissimilar atoms ( CH4, H2O, HNO3, and HF )
Elemental solids ( Diamond (Carbon), silicon,
germanium )
Number of covalent bond = { 8 – (No. of valence
electrons) }
Chlorine :8-7=1 can bond to only one other atom.
Carbon: 8-4=4 can bond to four other atoms.
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Covalent bonds may be very strong (diamond) or
very weak (bismuth). See Table 2.3.
Primary Interatomic Bonding (Contd.)
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It is possible to have interatomic bonds that are partially
ionic and partially covalent, and, in fact, very few
compounds exhibit pure ionic or covalent bonding.
For a compound, the degree of either bond type depends
on the relative positions of the constituent atoms in the
periodic table (Figure 2.6) or the difference in their
electronegativities (Figure 2.7).
The percent ionic character of a bond between elements
A and B (A being the most electronegative) may be
approximated by the expression:
% ionic character = {1 – exp[-(0.25)(XA – XB)2]} x 100
Where XA and XB are the electronegativities for the respective
elements.
Metallic Bonding
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Large inter-atomic forces are created
by the sharing of electrons in a
delocalized manner to form strong
non-directional bonding.
Cloud of electrons surrounding the positively charged
cores
Secondary Atomic & Molecular Bonds [Van
der Waals Bonds]
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Physical Bonds (no electron involvement).
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Weak electrostatic bonds.
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Bonding occurs between atomic or molecular
electric dipoles.
Bonding Energies are low. The bonds are
relatively weak.
Secondary Atomic & Molecular Bonds
[Van der Waals Bonds]
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Fluctuating Induced Dipole Bonds
• Weak electric dipole bonding can take place among
atoms due to an instantaneous asymmetrical
distribution of electron densities around their nuclei.
• This type of bonding is termed fluctuation since the
electron density is continuously changing.
• Example: Bonding between Argon atoms
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Polar Molecule-Induced Dipole Bonds
• Weak intermolecular bonds are formed between atoms
or molecules which possess permanent dipoles.
• A dipole exists in a molecule if there is asymmetry in its
electron density distribution.
Fluctuating Dipole Bonds
Bonding Between Argon electric dipoles
Argon atom
+
Argon Electric Diploe
≈
Van der waals bond
Polar Molecule-Induced Dipole Bonds
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Permanent dipole exist in some molecules.
Such molecules are called Polar Molecules.
Polar Molecules can induce dipoles in
adjacent non-polar molecules and bonding
can take place.
Example: HCl molecule
Cl
H
+
-
Permanent Dipole Bonds
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Van der waal bonds also occur between
permanent polar molecules.
The bonding energies are higher than the
fluctuating induced or polar molecule
induced bonds.
The strongest Secondary Bonding is
Hydrogen Bond.
Examples of Hydrogen Bonding:
HF, H2O, NH3