The classical picture of the atom
J.J. Thomson’s Cathode Tube
The Atom : J. J. Thomson (1856-1940)
The Atom based on Thomson’s experiment
Mass of electron
Rutherford Experiment
Rutherford Experiment
The Nucleus
Modern View
3.2. Electromagnetic Radiation and Quantization
Spectrum
Electromagnetic radiation
Electromagnetic Radiation
ELECTROMAGNETIC RADIATION
Electromagnetic Radiation - Characteristics
Radio in the 909kHz. What wavelength does it correspond to?
Nature of Matter
Photoelectric effect
Dual Nature of Light
De Broglie 1924
Diffraction
How to test the wave properties of an electron?
How to test the wave properties of an electron?
Diffraction
Conclusion
Atomic Spectrum of Hydrogen
Table 3.4. The atomic spectrum of hydrogen
Atomic Spectrum of Hydrogen
3.3.2: The Bohr Model
The Bohr Model
The Bohr Model
Wave Function and Atomic Orbitals
De Broglie
2.2 SChRONDINGER EQUATION
Quantum Mechanical Description of the Atom
The Schrödinger equation
Schrodinger Wave Equation
Cartesian and Spherical Coordinate
The wavefunction
Homework-2
Wave Equation for the Hydrogen Atom
Quantum numbers :
Radial and Angular Wave Function for 1s derived from Schrodinger Equation
s orbitals
Physical Meaning of Orbitals
p orbitals
d orbitals
f orbitals
Schrödinger Equation
Heisenberg uncertainty principle
The Hydrogen Atom : summary
Polyelectronic Model
Self-Consistent Field Method
https://www.youtube.com/watch?v=A6DiVspoZ1E
Many Electron Atoms
Electron Spin and Pauli Principle
History of the Periodic Table
The Aufbau Principle
Valence electrons
Rules
Rules
Rules
Rules
Hund’s Rule
Penetration Effect
Penetration Effect
Slater’s Rules The rules were devised semi-empirically by John C. Slater and published in 1930
Slater’s Rules The rules were devised semi-empirically by John C. Slater and published in 1930
Slater’s Rules for determining S for a specific electron
Slater’s Rules for determining S for a specific electron
Slater’s Rules for determining S for a specific electron
Solution
Solution
Comparison of The effective nuclear charge
Periodic Properties of Atoms : Ionization Energy
Trend in Atomic Properties : Ionization Energy
Trend in Atomic Properties : Ionization Energy
Trend in Atomic Properties : Ionization Energy
Trend in Atomic Properties : Atomic Radius
END
Alkali Metals – 1A
Trend in Atomic Properties : Ionization Energy
16.33M

Atomic structure and properties. (Chapter 3)

1.

General Chemistry I
Atomic Structure and Properties
Dr. Ould Ely
School of Science and Technology
1

2.

Chapter 3
Picture of the Atom
Electromagnetic radiation and Atomic Spectra
The Nature of Electron and Atomic Orbitals
Many-electron atoms
Atomic properties and Periodicity
Nuclear chemistry
2

3.

Part I
3.1.1 Atomic concept,
3.1.2 Subatomic particles,
3.1.3 Atomic structure: first ideas
3

4. The classical picture of the atom

Dalton Atomic Theory
1. Elements are made of tiny particles called atoms
2. The atoms of a given elements are identical
3. Chemical compounds are formed when atoms combine
with one another. A given compound has the same relative
numbers and types of atoms
4. Chemical reaction involve reorganization of the atoms.
The atom themselves are not changed.
4

5. J.J. Thomson’s Cathode Tube

• Charge-to-mass ratio
5

6. The Atom : J. J. Thomson (1856-1940)

e/m = -1.76 x 108 C/g
Experiment date
1898-1903
6

7. The Atom based on Thomson’s experiment

• A ray of particles is produced
between two metallic electrodes.
• These particles are negatively charged
• Since electrons could be produced
from electrodes made of various
types of metals, all atoms must
contain electrons
• e/m = -1.76 x 108 C/g
• Atoms = neutral! Positive charges are
located somewhere.
7

8. Mass of electron

Mass of a single electron
e= -1.6x10-19 C
m = 9.11 x 10-31 kg (Millikan)
8
http://www.youtube.com/watch?v=XMfYHag7Liw

9. Rutherford Experiment

Ernest Rutherford – 1911
• With Thomson Model :
a particles should travel
through the atom
without deflection.
http://sun.menloschool.org/~dspence/chemistry/atomic/ruth_expt.html
9

10. Rutherford Experiment

10

11. The Nucleus

Ernest Rutherford – 1911
Conclusion : Dense positive center with electrons far from the
nucleus
Its great density is
dramatically
demonstrated by the
fact that a piece of
nuclear material about
the size of a pea would
have a mass of 250
million tons
11

12. Modern View

A
Z
X
12

13. 3.2. Electromagnetic Radiation and Quantization

• 3.2.1: Electromagnetic Radiation
• 3.2.2: Quantization
• 3.2.3: The Atomic Spectrum of Hydrogen
13

14. Spectrum

14

15. Electromagnetic radiation

MRI
X-ray
Light
Microwave
Travel like a wave
Travel with the speed of light

16. Electromagnetic Radiation

Electromagnetic Radiation = a way for energy to travel.
2 oscillating fields (H and E)
16

17. ELECTROMAGNETIC RADIATION

17

18. Electromagnetic Radiation - Characteristics

l = wavelength = distance
between two peaks or two
troughs in a wave. (m)
= frequency = number of
waves / s at a specific
point of space. (s-1 or Hz)
l 1/
l c
Because speed = c
= 3x108 m/s
The radiation with the shortest
wavelength has the highest frequency
18

19. Radio in the 909kHz. What wavelength does it correspond to?

l = c/ = 330 m
C = 2.998 108 ms-1
= 909. 103 s-1
19

20. Nature of Matter

At the end of the 19th century :
Matter ≠ Energy
Matter = particles and Energy = electromagnetic radiations
Max Planck and the black body radiation :
Classic : matter can absorb or emit any
quantity of energy no maximum
infinite intensity at very low wavelength.
Quantum : Energy could only be gained
or emitted in whole number multiples of
h . h = Plank’s constant = 6.626x10-34Js
DE = nhn
20

21. Photoelectric effect

Albert Einstein Theory :
Energy itself is quantified and radiation could be seen as a stream of particles (photons)!
Ephoton = hn =
hc
l
Photoelectric effect
When UV radiation hits a metal surface, electrons are ejected –
photoelectric effect. (in 1905 explained by Albert Einstein using
a quantum approach)
h = + EKE
- work function – minimum energy required to remove the
electron
EKE – kinetic energy of the ejected electron
21

22.

When copper is bombarded with high-energy electrons, X
rays are emitted. Calculate the energy (in joules) associated
with the photons if the wavelength of the X rays is 0.154 nm.
E=hx
E=hxc/l
E = 6.63 x 10-34 (J•s) x 3.00 x 10 8 (m/s) / 0.154 x 10-9 (m)
E = 1.29 x 10 -15 J
22

23. Dual Nature of Light

Energy – Mass relationship :
A particle but also a wave :
E mc 2
E
m 2
c
Summary :
- Energy is quantized
- Only discrete units of energy (quanta) could be transferred
- Dual nature of light
23

24. De Broglie 1924

l Proportional to h/m
l = h/m
H :Planck Constant
M : masse
: velocity
24

25. Diffraction

What is the wavelength for an electron?
Me = 9.11x10-31 kg
Ve = 1.0x107 m/s
le
cc
cc
1 J = 1 kg.m2/s2
6.626x10-34Js
2
kgm
6.626x10-34
-11
s
le =
=
7.3x10
m
-31
7
( 9.11x10 kg) (1.0x10 m/ s)
The electron has a WL similar to the spacing of atoms in a crystal.
Confirmed for Ni crystal.
Diffraction : result of light scattered from a regular array of points or
lines.
25

26. How to test the wave properties of an electron?

26

27. How to test the wave properties of an electron?

27

28. Diffraction

When X-rays are scattered by ordered atoms Diffraction pattern.
28

29. Conclusion

All matter exhibits both particulate and wave
properties.
Large particles : mainly particle
Small particles : mainly wave
Intermediate particles (electron) : both
29

30. Atomic Spectrum of Hydrogen

When a high energy
discharge is passed
through H2 H-H breaks
excited H atoms.
Release of energy
Emission spectrum.
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