Chapter 2 Atomic Structure and Interatomic Bonding

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Chapter 2Atomic Structure andInteratomic
Bonding
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Why study this?

• An important reason is that, in some instances,
  the type of
• bond allows us to explain a materials properties.

relatively soft
This dramatic difference is directly
attributable to the type of interatomic bonding.
graphite
Carbon
hardest known material
diamond
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2.1 Introduction

• Some of the important properties of solid
• materials depend on geometrical atomic
• arrangements, and also the interactions that
• exist among constituent atoms or molecules.
• Fundamental concepts
• Electrons in atoms
• Periodic table
• Atomic Bonding in Solids

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2.2 Fundamental concepts

• Atomic number of an element equal to the number
• of protons (??) in the nucleus.
• Atomic mass may be expressed as the sum of the
• masses of protons and neutrons (??) within the
  nucleus.
• Atoms of the same element can have different
  numbers
• of neutrons, which are called isotopes (???).
• The atomic weight of an element corresponds to
• the weighted average of the atomic masses of the
• atoms naturally occurring isotopes.

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2.2 Fundamental concepts

• The atomic mass units (amu) may be used for
• computations of atomic weight. 1 amu is defined
  as
• 1/12 of the atomic mass of 12C atom.
• 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
  6.02?1023
• atoms or molecules.
• 1 amu/atom1 g/mol

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2.3 Electrons in atoms
During late 19th century it was realized that
many phenomena involving electrons in solids
could not be explained in terms of classical
mechanics. Quantum mechanics (????) is
established to describe the phenomena in atomic
and subatomic systems.

• Bohr atomic model

Electrons are assumed to revolve around the
atomic nucleus in discrete orbitals, and the
position of any particular electron is more or
less well defined in terms of its orbital.
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Quantized energy Energies of electrons are
not continuous, and electrons have only specific
values of energy. An electron may change
energy, but in doing so it must make a quantum
jump either to an allowed higher energy (with
absorption of energy) or to a lower energy (with
emission of energy).
The allowed electron energy state for the Bohr
hydrogen atom
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Concept Check

• Why are the atomic weights of the elements
  generally
• not integers? Cite two reasons.

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Concept Check

• Why are the atomic weights of the elements
  generally
• not integers? Cite two reasons.

1. Atomic mass is not integer! Atomic mass
are most often expressed by amu. Atomic
mass(MprotonMneutron)/amu
amu1/12MC12(MprotonMneutron)/2 2. Atomic
weight weighted average of atomic masses of
natural isotopes
Isotope Atomic mass Abundance Abundance
Isotope Atomic mass Standard Range
28Si 27.976 926 532 46(194) 92.2297(7) 92.2192.25
29Si 28.976 494 700(22) 4.6832(5) 4.694.67
30Si 29.973 770 171(32) 3.0872(5) 3.103.08
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Wave-mechanical model
Bohr model was eventually found to have some
significant limitations because of its inability
to explain several phenomena involving electrons.

• The electron is considered to exhibit both
  wavelike
• and particle-like characteristics.
• An electron is no longer treated as a particle
  moving
• in a discrete orbit but rather, position is
  considered to
• be the probability of an electrons being at
  various
• locations around the nucleus.
• In other words, position is described by a
  probability
• distribution or electron cloud.

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Wave-mechanical
Bohr
Comparison of the (a) Bohr and (b)
wave-mechanical atom models for the H atom in
terms of electron distribution.
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Quantum numbers

• Using wave mechanics, every electron in an
• atom is characterized by four parameters called
• quantum numbers.
• Principal quantum number, n
• Second quantum number, l
• Third quantum number, ml
• Fourth quantum number, ms

the spin orientations
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Quantum numbers

• Principle quantum number (????)
• Specify the shells
• K, L, M, N, O (n1, 2, 3, 4, 5)
• Related to the distance of an electron from the
• nucleus, or its position

• Second quantum number (????)
• Specify the subshells
• s, p, d, f (l0, 1, 2, 3n-1)
• Related to the shape of the electron subshell

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Quantum numbers

• Third quantum number (????)
• determine the number of energy states for each
• subshell
• ml 0, 1, l

• Fourth quantum number (?????)
• related to spin moment of an electron
• ms 1/2, -1/2
• one for each of the spin orientations

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Quantum numbers
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Quantum numbers

• The smaller the n, the
• lower the energy level.
• 1slt2slt3s
• Within each shell, the
• energy of a subshell level
• increases with the value
• of the l quantum number.
• 4slt4plt4dlt4f
• There may be overlap
• in energy state of one
• shell with states in an
• adjacent shell.
• 3dgt4s

the relative energies of the electrons for the
various shells and subshells
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Electron configurations

• Pauli exclusion principle (???????)
• Each electron state can hold no more than two
  
• electrons, which must have opposite spins.

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Electron configurations

• Not all possible states in an atom are
  filled with
• electrons. For most atoms, the electrons prefer
  to fill
• up the lowest possible energy states in the
  electron
• shells and subshells.
• When all the electrons occupy the lowest possible
  
• energies, an atom is said to be in its ground
  state.
• Electron configuration (structure of an atom)
• represents the manner in which the energy states
• are occupied in an atom.

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Electron configurations
Electron configurations for Na is 1s22s22p63s1
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Electron configurations for some elements
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Electron configurations

• Valence electrons (???) refer to the electrons
• occupy the outermost filled shell.
• They participate in the bonding between atoms.
• Many of the physical and chemical properties of
• solids are based on these valence electrons.

Na atom
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Electron configurations

• Stable electron configurations
• are the state in which valence
• electron shell are completely
• filled.

Ar1s22s22p63s23p6
Some atoms that have unfilled valence shells
can become stable electron configurations by
gaining or losing electrons to form charged
ions, or by sharing electrons with other atoms.
Basis for chemical reaction, also for atomic
bonding in solids
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Concept Check
Give electron configurations for the Fe3 (26)
and S2- (16) ions.
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2.3 The Periodic Table
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The Periodic Table
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The Periodic Table

• Electropositive element metallic in nature and
  give
• up electrons in chemical reaction to produce
• positive ions.
• Electronegative element nonmetallic in nature
  and
• accept electrons in chemical reaction to form
• negative ions.
• Electronegativity describes the tendency of an
• atom to gain an electron.

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Electronegativity
Electronegativity increases in moving from
left to right and from bottom to top.
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2.4 Atomic bonding in solids

• Bonding Forces and Energies
• consider the interaction between two isolated
  atoms
• as they are brought into close proximity from an
• infinite separation.

FNFAFR
FN net force (??) FA attractive force (??) FR
repulsive force (??)
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Equilibrium state
FAFR0 equilibrium spacing r0
Bonding energy (E0) represents the
energy required to separate two equilibrium
atoms to an infinite separation.
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Atomic bonding in solids

• Primary bonds
• Ionic bonds
• Covalent bonds
• Metallic bonds
• Secondary bonds
• van der Waals
• Hydrogen bonding

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Ionic Bonding
An ionic bond is created between two unlike
atoms with different electronegativities.
When sodium donates its valence electron to
chlorine, each becomes an ion attraction
occurs, and the ionic bond is formed.
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Ionic Bonding

• Features of ionic bonding
• Nondirectional ionic bonding
• the magnitude of the bond is
• equal in all directions around an ion.
• Bonding energy range 600-1500 kJ/mol
• Most ceramic materials

high melting temperatures hard and brittle
electrically and thermally insulative
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Covalent Bonding
Carbon atom has four valence electrons and H
atom has only one valence electron. H 1s1
1s2 C 2s1p3 2s22p6
Covalent bonding in a CH4 molecule
Covalent bonding requires that electrons be
shared between adjacent atoms in such a way that
each atom has its outer sp orbital filled.
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Covalent Bonding

• Directional
• It exists only in the direction between one
  atom and
• another that participates in the electron
  sharing.
• Number of possible covalent bonds is determined
  by
• the number of valence electrons. (8-N)
• Bonding energy
• Strong as diamond (very hard, Tm 3550C)
• Weak as bismuth (Tm 270C)
• Nonmetallic elemental molecules (H2, O2, N2)
• molecules containing dissimilar atoms (CH4,
  H2O)
• elemental solids (diamond, Si)

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Covalent Bonding

• Mixed between ionic bond and covalent bond
• It is possible to have interatomic bonds that
  are partially
• ionic and partially covalent.
• For a compound, the degree of either bond
  type depends
• on the difference in their electronegativities.
  The greater
• the difference in electronegativity, the more
  ionic the
• bonds. And vise versa.
• The percentage ionic character of a bond
  between
• elements A and B may be approximated by

ionic character1-exp-(0.25)(XA-XB)2?100
XA and XB electronegativity for respective
elements
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Metallic Bonding
The metallic bond forms when atoms give up
their valence electrons, which then form an
electron sea. Valence electrons are not
bound to any particular atom in the solid and
are free to drift throughout the entire metal.

The positively charged atom cores are bonded by
mutual attraction to the negatively charged
electrons.
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Metallic Bonding

• Features of ionic bonding
• Nondirectional
• Bonding energy
• may be weak as Hg (Tm -39 ?C)
• may be strong as W (Tm 3410 ?C)
• Materials
• metals and alloys
• Good conductor of electricity and heat
• as a result of free electrons

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Atomic bonding in solids

• Primary bonds
• Ionic bonds
• Covalent bonds
• Metallic bonds
• Secondary bonds
• van der Waals
• Hydrogen bonding

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Secondary bonding
Secondary bonds are relatively weak in
comparison to the primary ones, and bonding
energies are typically on the order of only 10
kJ/mol.
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Secondary bonding
Secondary bonding arises from the coulombic
attraction of the electric dipoles contained in
atoms or molecules.
An Electric dipole exists when there is some
separation in positive and negative portions of
an atom or molecule .
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van der Waals bonding
Fluctuating dipole bonds Permanent dipole bonds
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Fluctuating Dipole Bonds
electrically symmetric
induced atomic dipole
Constant vibrational motion (existed in all
atoms) can cause temporary and short-lived
distortions of the electrical symmetry to induce
a dipole. The liquefaction and the
solidification of the inert gases and other
electrically neutral and symmetric molecules such
as H2 and Cl2 are realized by this type of
bonding.
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Permanent Dipole Bonds
Permanent dipole moments exist in polar
molecules, which have an asymmetrical
arrangement of positively and negatively charged
regions.
HCl molecule
A permanent dipole moment arises from net
positive and negative charges that are
respectively associated with the hydrogen and
chlorine ends of the HCl molecule.
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Hydrogen bonds

• Hydrogen bond, is a special case of polar
  molecule
• bonding. It occurs between molecules in which
• hydrogen is covalently bonded to fluorine (as in
  HF),
• oxygen (as in H2O), and nitrogen (as in NH3).

Schematic representation of hydrogen bonding in
hydrogen fluoride (HF)
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Summary

• A survey of the fundamentals of atomic structure
• Bohr and wave-mechanical model

Quantum number Electron configurations Periodic
table
Primary bonds (ionic, covalent and metallic)
Secondary bonds van der Waals hydrogen bonds