This document provides an overview of atomic structure and theory. It begins by discussing how an understanding of atomic structure relates to materials properties. It then reviews the atomic models of Democritus, Dalton, Thomson, Rutherford, Bohr, and others. Key concepts covered include the nucleus, protons, neutrons, electrons, atomic number, mass number, and energy levels. The document also discusses atomic bonding mechanisms like ionic bonding, covalent bonding, and metallic bonding. Finally, it introduces topics like energy bands and the quantization of energy and light.

1. Electrical Materials and Technology
PCE3302
Chapter One
Review of atomic theory of matters
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2. Introduction
 Selection of appropriate material for the intended purpose depends on the
knowledge of properties of materials. i.e.,
 Physical properties
 Electrical
 Mechanical
 Thermal and
 Chemical
 These properties of materials depend on the atomic structure of that material.
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3. Atomic Structure
All matter is composed of atoms.
 Understanding the structure of atoms is critical to understanding the
properties of matter
 Atom – smallest particle of an element that can exist alone
Two regions of an atom
Nucleus
 Center of atom
 Protons and neutrons
Electron “cloud”
 Area surrounding nucleus containing electrons
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4. cont’d…………
 Proton – Positive charge (+), 1 atomic mass unit (amu); found in the nucleus
 amu -Approximate mass of a proton or a neutron
 Neutron – Neutral charge (0), 1 amu; found in the nucleus
 Electron – Negative charge (-), mass is VERY small
 In a neutral atom, the number of protons equals the number of electrons, so the
overall charge is zero (0)
 Example/ Helium, with an atomic number of 2, has 2 protons and 2 electrons
when stable
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5. Counting Atoms
Atomic Number
 Number of protons in nucleus
 The number of protons determines identity of the element!!
 Mass Number (Atomic Mass)
 Number of protons + neutrons
Units are g/mol
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6. Atomic model
Democritus 460 BC
 Greek Philosopher
 Suggested world was made of two things – empty space and “atomos”
 Atomos – Greek word for uncuttable
2 Main ideas
 Atoms are the smallest possible particle of matter
 There are different types of atoms for each material
 His ideas did agree with later scientific theory, but did
not explain chemical behavior, and was not based
on the scientific method – but just philosophy
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7. Cont’d……
John Dalton’s Atomic Theory (1766 - 1844)
 Dalton was an English school teacher who performed
many experiments on atoms.
 All matter is made of atoms.
 Atoms of one element are all the same.
 Atoms cannot be broken down into smaller parts
 Compounds form by combining atoms
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8. Dalton’s Early Atomic Mode
“Billiard Ball” model
He envisioned atoms as solid, hard spheres, like billiard(pool) balls
So he used wooden balls to model them
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9. J.J. Thomson (1856 - 1940)
 Discovered the electron
 He was the first scientist to show the atom was made of even smaller things
 Used the Cathode ray tube to discover electrons
1897
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10. Thomson’s Experiment
Cathode ray tube
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Voltage source
Vacuum tube
Metal Disks

11. cont’d………
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Voltage source

12. cont’d………
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Voltage source

13. cont’d………
Passing an electric current makes a beam appear to move from the
negative to the positive end
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Voltage source

14. cont’d………
Passing an electric current makes a beam appear to move from the
negative to the positive end
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Voltage source

15. cont’d………
By adding an electric field
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+
-
Voltage source

16. cont’d………
 Adding an electric field cause the beam to move toward the positive plate.
 Thomson concluded the beam was made of negative moving pieces.
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+
-
Voltage source

17. Thomson’s “Plum Pudding” Atom Model
Plum and pudding atom modeling
..DownloadsCathode Rays Lead to Thomson's Model of the Atom.mp4
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18. Ernest Rutherford (1871 - 1937)
Discovered the nucleus of a gold atom
with his “gold foil” experiment
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19. cont’d………
Using J.J Thomson’s Plum Pudding atomic model, Rutherford
predicted the alpha particles would pass straight though the gold foil.
That’s not what happened.
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20. cont’d………
Gold Foil Experiment Results
Most alpha particles go straight through the gold foil A few alpha
particles are sharply deflected
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21. Rutherford’s Conclusion
The atom is mostly empty space.
There is a small, dense center with a positive charge.
Rutherford discovered the nucleus in atoms
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22. cont’d………
Rutherford’s Atomic Model
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23. Rutherford’s Contribution to the Atomic Theory
The atom is mostly empty space.
The nucleus is a small, dense core with a positive charge.
..DownloadsRuther's Alpha Scattering Experiment2.mp4
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24. Bohr Model (1871 - 1937)
In 1913, the Danish scientist Niels Bohr proposed an improvement.
In his model, he placed each electron in a specific energy level.
According to Bohr’s atomic model, electrons move in definite orbits
around the nucleus, much like planets circle the sun.
These orbits, or energy levels, are located at certain distances from the
nucleus.
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25. Cont’d…….
Additionally, the electrons can jump from a path in one level to a path in
another level (depending on their energy).
He proposed the following
 Protons and neutrons are in the nucleus
 Electrons can only be certain distances from the nucleus
 The electrons orbit the nucleus at fixed energy levels
 The electrons must absorb or emit a fixed amount of energy to travel
between these energy levels
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26. Atomic bonding
Chemical bonds :-forces that hold together atoms when they combine
 produce by interaction of electrons in outer levels of two or more atoms
 can form by sharing or transferring electrons
 All atoms will try to achieve a stable octet
Two Major Types of Bonding
Ionic Bonding
 Form when electrons are transferred from one atom to another
 Forms ionic compounds ,
 Transfer of e- 11/13/2022 26

27. Ionic Bonding
 Two neutral atoms close to each can undergo an ionization process in
order to obtain a full valence shell
 Due to ionization, electrons are transferred from one atom to the
other
 The atom giving up the electron becomes positively charged (cation)
 The atom taking up the electron becomes negatively charged (anion)
 The ions are bonded due to coulombic forces of attraction
 Generally, metallic atoms donate electrons to non-metallic atoms
 Examples: NaCl, KCl, MgBr2 etc.
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28. Ionic Bonding
Ionic bond between sodium and chloride forming NaCl
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29. Covalent Bonding
 The outer electron levels of atoms, which are close to each other, can
interact
 The interaction leads to a sharing of electrons between the atoms
 One pair of electrons shared => single covalent bond
 Two pairs of electrons shared => double covalent bond and so on
 The shared electrons are said to be delocalized i.e. they do not belong
to any particular atom
 Generally, between non-metallic atoms
Examples: H2, CO2, C6H12O6 and other molecules
polar molecule :- have covalent bonds between but atoms do not share
electrons equally
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30. Cont’d….
Single hydrogen atom (left) and two hydrogen atoms forming a covalent
bond with a shared electron pair (right)
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31. Metallic bonds
 Atoms come together, electrons from outer shell of atoms share space
with neighboring atoms.
 The electrons can move freely within the atom orbitals.
 Sharing of ‘free’ electrons among a lattice of positively charged ions
(An array of positive ions in a sea of electrons)
 Electrostatic attractive forces between delocalized electrons and
positively charged metal ions.
 Between metallic atoms
 Examples: Ni, Fe and other metals
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32. Cont’d….
 The outer electrons are so weakly bound to metal atoms that they are free
to roam across the entire metal.
 This results in a lattice of positively ions in a sea of communal electrons
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33. Energy Bands
Electrons try to occupy the
lowest energy band possible
Not every energy level is a
legal state for an electron to
occupy
These legal states tend to
arrange themselves in bands
Allowed
Energy
States
Disallowed
Energy
States
Increasing
Electron
Energy
Energy Bands
}
}

34. Energy Bands
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EC
EV
Valence Band
Conduction Band
Last filled energy
band at T=0K
First unfilled energy
band at T=0K

35. Band Diagrams
Band Diagram Representation
Energy plotted as a function of position
EC  Conduction band
 Lowest energy state for a free electron
EV  Valence band
 Highest energy state for filled outer shells
EG  Band gap
 Difference in energy levels between EC and EV
 No electrons (e-) in the bandgap (only above EC or below EV)
 EG = 1.12eV in Silicon
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EG
EC
EV
Increasing electron energy
Increasing voltage

36. Energy Bands In Solids
Explains and distinguish between conductors, insulators and semi
conductors
 Conduction Band – small energy can remove an electron from an atom
 Valence Band – Below the conduction band
 Filled Band- Near the nucleus
 Insulators- Conduction band is empty but valence is almost
completely filled+ wide energy gap between the two bands (Poor
conductivity)
 Conductors- Overlapping of both conduction and valence bands
 Conduction es move almost twice as fast as holes
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37. Example 1
Aluminum (Al) (no periodic table)
Protons = 13
Electrons =
Neutrons = 14
Atomic Number =
Atomic Mass =
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38. Quantization of Energy
 Energy is quantized in some systems, meaning that the system can
have only certain energies and not a continuum of energies, unlike the
classical case
The quantum view of the atom
Observation: An atom will only absorb or release light at discrete
frequencies
Explanation:
 Absorption or emission of light is caused by electron energy
transitions within the atom
 The energy carried by light is connected with its frequency.
 Electrons are only allowed to move between discrete energy levels in
the atom
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39. Quantization
Classical Physics:
Quantities are continuous.
Quantum Physics:
Some quantities are limited
to a discrete set of values.
Example: charge, Q = N.e

40. Light is also quantized
 Light seems to carry energy in discrete packets
 We call these packets photons
 In this sense, light behaves like a particle
 But light also has many wave-like properties
 Particle-wave duality: it is both at the same time!
The photoelectric effect
 Metal plate illuminated by a light beam
 If UV light is used, electrons are emitted
 If visible light is used, no electrons are emitted no matter how bright
the light
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41. Cont’d…..
Explanation:
 A minimum amount of energy is needed to remove an electron from
the metal
 The energy of UV light must be greater than the energy of visible light
 Energy and frequency are related by
E = hn
where h is called Plank’s constant
h = 6.625 x 10–34 J s
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42. Cont’d……
Reading Assignment
Schrödinger Equation
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