Title: Atomic Structure and Bonding
1Atomic Structure and Bonding
- Chapter 2
- Callister, 2000.
2Outline
- Atomic Structure
- Electron Configurations and Periodic Table
- Atomic Bonding in Solid Materials
- -----Primary interatomic bonding
- -----Secondary bonding
-
3What is the importance of atomic structure?
- Some properties of solid materials depend on
geometrical atomic - arrangements and interactions among the atoms,
which eventually - are controlled by the subatomic structure of the
materials. Therefore, - we will learn about the subatomic structure,
elctronic configurations, - and major bondings holding the atoms together.
- For example Carbon (pure) can exist as graphite
and diamond. - Graphite is soft and greasy feel to it,
- Diamond is the hardest known material.
- This difference is because of the type of
interatomic bonding in - graphite and diamond.
-
4- Atom nucleus (protonsneutron) and electrons
- Electrons are negatively charged, protons are
positively charged - particles.
- Charge 1.60x10-19 C
- For an electrically neutral atom the number of
protons equal to that of - electrons.
- Neutrons are electrically neutral.
- Masses of these particles are very small and
- Mass of proton mass of neutrons1.67x10-27 kg
- Mass of electron 9.11x10-31 kg
- Each element is characterized by the number of
protons in the nucleus, - Atomic Number (Z). Z1 (hydrogen) -94 (plutonium)
(naturally existing - elements)
- Atomic mass (A) for an atom is the sum of the
masses of protons and nucleus (electrons are not
considered, because ...?)
5-
AZN - For all atoms of an element the number of protons
are the same, but the number of neutrons may
vary, which may vary the atomic mass. The atoms
with two or more atomic masses are called
isotopes. Then they average the different masses
of the isotopes to find out the atomic weight of
the element. - The atomic mass unit (amu) is used for the
computation of atomic weight. - Scale 1 amu1/12 of the atomic mass of the
Carbon (C) - (A12.00000 for carbon 12 isotope)
- 1 amu/atom 1 g/mol ( 1 mol of a
substance6.023x1023 atoms) - For example Fe
- A55.85 amu/atom or 55.85 g/mol (this is most
commonly used form)
6Atomic Models
- Bohr Atomic Model is the simplest model of the
atom in which electrons are assumed to be
positioned around the nucleus in discrete
orbitals and position of the electron is more or
less well defined in its orbital.
The energies of the electrons are quantized
electrons are permitted to have specific levels
of energy (energy levels or states). An electron
can change its energy level, for ex., To a higher
level by absorbing energy To a lower level by
emitting energy
7The zero reference is the unbound or free
electron. Single electron in H can exist only in
one of the states acc. to Bohr atom model.
8- Bohr atomic model describe the electrons in terms
of their positions (orbitals) and energy
(quantized energy levels by Rydberg equation).
Bohrs model was not able to explain
quantitatively the spectra of the atoms more
complex than hydrogen and the model could not
have been modified. - Limitations of Bohr model was resolved by
wave-mechanical model of the atom. In this model,
electrons exhibit wavelike and particle like
characteristics and position is considered as the
probability of an electrons being at various
locations or electron cloud. -
NOTE Atomic spectra When an electric discharge
(spark) passes through a gas (H2), it excites or
energizes the atoms of the gas. More specifically
it excites the electrons of the atoms. The atoms
then emit the absorbed energy in the form of
light as the electrons return to a lower energy
state. When a narrow beam of light is passed
through a prism, a spectrum of colors as
individual lines can be seen. This is called
atomic spectrum or emission spectrum.
9- Electron Waves in Atoms
- Schrodinger (1187-1961), Austrian physicist
- He found that electron waves are standing waves,
similar to guitar string, they can have many
different wave forms and patterns. Each of these
waveforms are called orbitals. Most of the waves
have different levels of energy. Energy change
within an atom is simply the result of an
electron changing from a wave pattern with one
energy to another. - Ground state The most stable state of an atom
and atoms electrons have waveforms with the
lowest possible energies. - The shapes of the wave patterns are important
since the amplitude of the wave is related to the
likelihood of finding the electron there. This is
important in chemical bonding...... - The characteristics of a wave can be described by
using Quantum Numbers. -
10- Quantum Numbers
- Every electron in an atom is characterized by
four numbers. - Principal quantum number (n) n1 to infinity and
it is shell number. - n1 (first shell) n second shell, so forth... or
n1 shell designation K, - n2 shell designation L, so forth....
- The principal quantum number is related to the
size of the electron wave (how far the wave can
extend from the nucleus) and the energy of the
orbital, i.e., as n increases the energy level of
the orbital also increases. - 2) The Secondary Quantum Number (l) signifies the
subshells. In other words, l divides the shells
into subshells. - For a given n, l can range from l 0 to l n-1.
- When n1, then l 0 when n1, there is a single
subshell - Subshells are defined using a letter code
- Value of l 0 1 2 3 4 5....
- Letter s p d f g h
11- 2p means subshell with n2 and l 1.
- This number determines the shape of the orbital,
and also affects the energy. In atoms with two or
more electrons, the subshells within a given
shell differ slightly with the energy of the
subshell increasing with increasing l. Therefore,
s subshell has the lowest energy, p is the next
lowest followed by d, then f so on.. - 3) The magnetic quantum number (ml) divides the
subshells into individual orbitals or energy
states. - For a given l, ml can range from l to l
- When l0 (s subshell) then ml is zero.
- When l1 (p subshell), then ml 1, 0, -1.
- Therefore s subshell has 1, p has 3, d has 5, f
has 7 orbitals...or energy states. - In the absence of magnetic field, the states in
each subshell have similar enrgies but in the
presence of magnetic field, each state has
slightly different levels of energy.
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13- Paulis Exclusion Principle
- This principle states that no two electrons can
have the identical values for all four of their
quantum numbers. - Suppose the 1s orbital of an atom has two
electrons. n1, l0 and ml0 for the two
electrons. - Therefore, there is need for the fourth quantum
number. - 4) Spin quantum number (ms)
- One of the electron has ms of 1/2 and the other
of -1/2. - The maximum number of electrons in an orbital is
two and when two - Electrons are in the same orbital they must have
different spin numbers.
14The whole picture
15Electron configurations
- The electron configuration of an atom represents
the manner in which the states are occupied. - Conventionally, the number of electrons in each
subshell is indicated by a superscript after - the shell-subshell designation.
- For example H 1s1
- He 1s2
- Na 1s22s22p63s1
16- The electrons occupying the outermost filled
shell are called valence electrons. The electrons
participate in bonding and therefore they are
very important. Moreover many of the physical
and chemical properties of solids are based on
these valence electrons. - Some atoms have stable electron configurations
that is the states within the outermost shell are
completely filled. For example, Ne, Ar, Kr and
He. These are inert, or noble, gases, which are
unreactive chemically. - Some atoms of the elements having unfilled
valence shells assume stable electron
configurations by gaining or losing electrons to
form charged ions or by sharing electrons with
other atoms. This is in fact the basis of the
chemical reactions and atomic bonding. - Some of the elements have s and p orbitals
forming a hybrid spn orbital, where n is the
number of p orbitals involved (1,2 or 3). The
reason for such a hybrid orbitalis the lower
energy state for the valence electrons. For
example sp3 hybrid orbital control Carbon
chemistry. 3A, 4A, and 5A group elements of the
periodic table form hybrid orbitals.
17Periodic Table
- Elements are classified according to electron
configuration in this table. - The elements are placed with increasing atomic
number in seven - horizontal rows, called periods.
- All elements in the same column, or group, have
similar valence - electron configurations, as a result similar
properties. These properties - change gradually and systematically across each
period moving - horizontally.
-
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19- Group 0 inert gases (filled electron shells)
- Group VIIA (halogens) and VIA elements one and
two electrons deficient respectively from having
stable configurations. - Group IA and IIA are alkali (except H) and
alkaline earth metals, with one and two excessive
electrons respectively from stable
configurations. - Elements in three long periods, Groups from IIIB
to IIB are transition metals, with partially
filled d electron states and in some cases one or
two electrons in the next higher shell. - Elements in Group IIIA, IVA and VA have
characteristics between metal and nonmetals
because of their valence electron configurations. - Electropositive elements are the elements capable
of giving up their electrons to become positively
charged ions (located on the left of the table.) - Electronegative elements are the ones ready to
accept electrons to form negatively charged ions
or to share their electrons. Atoms are more
electronegative if their outer shells are almost
full. -
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21Atomic Bonding in Solids
- Consider the interaction of two isolated atoms as
they are getting closer from an infinite
separation - Interactions grow up as they approach.
- Interaction forces. Attraction (FA) and repulsion
(FR) - Magnitude of the interaction is a fcn. of
interatomic distance.
Where FN is the net force between the two atoms.
22Net force is zero (equilibrium state) r0equilibri
um spacing (0.3 nm)
Bonding energy (minimum net energy)
23When there are more than two atoms, force and
energy interactions among many atoms have to be
considered and there will be bonding energies
among each atom analogous to E0. There are
number of material properties depending the
magnitude of bonding energy and shape of the
curve. For example melting point is higher for
the materials having larger bonding energies.
24- Mechanical stiffness of a material is dependent
on the shape of the - force versus interatomic separation curve. Slope
of the curve at - rr0 position is steep for very stiff materials.
Slope is shallower for - flexible materials.
- Linear coefficient of thermal expansion of the
material is also related - with the shape of the curve. Deep and narrow
curves correlates with - low coefficient of expansion.
-
Bonds (in general)
Primary Bonds
Secondary Bonds
metallic
ionic
covalent
25Primary Interatomic Bonds
- Ionic bonding found in compounds formed by
metallic and nonmetallic - elements. Atoms of a metallic element easily give
up their valence - electrons to the nonmetallic atoms in the search
of a stable or inert gas - configurations and an electrical charge as they
become ions. - For example,
Sodium can assume the electron structure of neon
by a transfer of its one valence 3s electron to
a chlorine atom, which has an electronic configur
ation of argon. In the structure of NaCl, Na and
Cl exist as ions.
26- There is an attractive force between the positive
and negative ions - (Coulombic forces).
- For two isolated ions, the attractive and
repulsive energies are - calculated as
-
A,B, and n are constants. The n is approximately
8. The ionic bonding is nondirectional, that is
the magnitude of the bond is equal in all
directions. The predominant bonding in ceramics
is ionic. Bonding energies 600-1500 kJ/mol
(high)
27- Covalent bonding stable electron
- configurations are assumed by
- sharing of electrons between adjacent
- atoms. Shared electron is considered
- to belong to both atoms.
- For example CH4 (methane)
- H feels like helium electron
- configuration, while C feels like neon
- electron configuration.
- Covalent bonding is directional, and it
- forms between two specific atoms
- and may exist only in the direction
- between one atom and another.
Many nonmetallic elemental molecules as well as
molecules containing dissimilar atoms are
covalently bonded. H2, Cl2, F2, H2O, HNO3, HF,
GaAs, etc.....
28- The number of covalent bonds possible for a
particular atom is determined by the number of
valence electrons. For 7 valence electrons, an
atom can have maximum 1 more bond (completing the
valence orbital electron number to eight), like
Cl atom. - N number of valence electrons
- 8-Nnumber of covalent bonds.
- Covalent bonds may be extremely strong (like in
diamonds) or may be weak (like in Bismuth). - Polymeric materials are covalently bonded
materials. - There are interatomic bonds that are partially
ionic and partially covalent. The degree of
either bond is controlled by the
electronegativities of the composing atoms. As
the electronegativity difference gets higher, the
bonding becomes more ionic. As the difference in
electronegativity becomes smaller, the bonding is
more covalent.
29A is the most electronegative element in the
equation.
Metallic Bonding is found in metals and theri
alloys. Metallic materials have one, two or at
most three valence electrons and these valence
electrons are not bound to any particular atom
in the solid and they are more or less free to
move throughout the entire metal. The
nonvalence electrons and atomic nuclei form ion
cores, which has a net positive charge equal in
magnitude to the total valence electron charge
per atom. Metallic bond is nondirectional.
Bonding may be weak or strong ranging from 68
kJ/mol for Hg to 850 kJ/mol for Tungsten.
30Metallic bonding can be seen between all
elemental metals. Metalling bonding explains
the heat and electric conductivity of the
metallic Materials as well as their ductility.
31Secondary Bonding or van der Waals Bonding
- Secondary bonds are weaker than primary bonds
(bonding energies in the order of 10 kJ/mol). - These bonds exist due to polarity of the atoms or
molecules. If there is a separation of positive
and negative sides of the atoms or molecules,
then there is an electric dipole. The bonding
results from the coulombic attraction of the
positive and negative ends.
Hydrogen bonding is a special type of secondary
bonding that form between some molecules, which
have hydrogen in their composition. For example,
water. Dipole interactions may occur between
induced dipoles, induced dipoles and polar,
polar molecules.
32- Fluctuating induced dipole bonds When the
overall electron distribution is symmetric, a
dipole may be created in an atom. This type of
bonds are weakest and they are also known as van
der Waals bonding. - Polar molecule-Induced Dipole Bonds Polar
molecules have asymmetrical arrangement of
positively and negatively charged regions. The
atoms may induce polarity to other atoms and
creating this type of interaction. HCl as an
example.
33- Permanent Dipole Bonds stronger than the
previous types of bonds. Hydrogen bonding is a
good example for this type. HF, H2O, NH3, etc..