Atomic Structure and Bonding - PowerPoint PPT Presentation

1 / 33
About This Presentation
Title:

Atomic Structure and Bonding

Description:

... is because of the type of interatomic bonding in. graphite and diamond. ... There are interatomic bonds that are partially ionic and partially covalent. ... – PowerPoint PPT presentation

Number of Views:1627
Avg rating:3.0/5.0
Slides: 34
Provided by: sib3
Category:

less

Transcript and Presenter's Notes

Title: Atomic Structure and Bonding


1
Atomic Structure and Bonding
  • Chapter 2
  • Callister, 2000.

2
Outline
  • Atomic Structure
  • Electron Configurations and Periodic Table
  • Atomic Bonding in Solid Materials
  • -----Primary interatomic bonding
  • -----Secondary bonding

3
What 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)

6
Atomic 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
7
The 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.

12
(No Transcript)
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.

14
The whole picture
15
Electron 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.

17
Periodic 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.

18
(No Transcript)
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.

20
(No Transcript)
21
Atomic 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.
22
Net force is zero (equilibrium state) r0equilibri
um spacing (0.3 nm)
Bonding energy (minimum net energy)
23
When 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
25
Primary 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.

29
A 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.
30
Metallic 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.
31
Secondary 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..
Write a Comment
User Comments (0)
About PowerShow.com