1. CHAPTER-6 6.1 Introduction (electromagnetic waves)
6.2 Bohr model and quantum mechanical model).
6.3 Quantum numbers, energy levels and orbitals.
6.4 Electronic configuration in atoms and ions.
6.5 Orbital diagrams and orbital shapes of atoms.

2. Electric Fields Gravity Fields Magnetic Fields

3. 6.1. Electromagnetic Waves
The electric and magnetic fields vibrate at right angles to the direction the wave travels so it is a transverse wave.

4. are distances – S.I. unit is the meter

5. Electromagnetic Radiation
Wavelength () : distance between consecutive crests.
Frequency () : number of waves that pass a given point in one second (SI Unit is s-1 or Hz).

6. The Electromagnetic Spectrum
Light travels in space as a wave.
In vacuum, speed of light is constant and given the symbol “c”, c = 3.00 x 108 m/s.
Light waves have amplitude, frequency, and wavelength.

7. The Electromagnetic Spectrum
The electromagnetic spectrum consists of all the different wavelengths of electromagnetic radiation, The only region in the entire electromagnetic spectrum that our eyes are sensitive to is the visible region.
Gamma rays have the shortest wavelengths, < 0.01 nanometers (about the size of an atomic nucleus). This is the highest frequency and most energetic region of the electromagnetic spectrum. Gamma rays can result from nuclear reactions.

8. The Electromagnetic Spectrum
X-rays range in wavelength from 0.01 – 10 nm (about the size of an atom). They are generated, when matter is irradiated by a beam of high-energy charged particles such as electrons.
Ultraviolet radiation has wavelengths of 10 – 310 nm (about the size of a virus). Stars produce a lot of ultraviolet light.
Visible light covers the range of wavelengths from 400 – 700 nm (from the size of a molecule to a protozoan). Our sun emits the most of its radiation in the visible range, which our eyes perceive as the colors of the rainbow. Our eyes are sensitive only to this small portion of the electromagnetic spectrum.

9. The Electromagnetic Spectrum
Infrared wavelengths span from 710 nm – 1 millimeter (from the width of a pinpoint to the size of small plant seeds). At a temperature of 37 oC, our bodies give off infrared wavelengths with a peak intensity near 900 nm.
Radio waves are longer than 1 mm waves, have the lowest energy. Radio stations use radio wavelengths of electromagnetic radiation to send signals that our radios then translate into sound. Radio stations transmit electromagnetic radiation, not sound. Our radios receive the electromagnetic radiation, decode the pattern and translate the pattern into sound.

10. Sample Calculation – wavelength/frequency conversion
Calculate the frequency of visible light having a wavelength of 485 nm?
Remember to use S.I. units in your calculations!
 = c / 
= (3.00 x 108 m/s) ÷ (485 x 10-9 m)
= x 1014 s-1
What colour of visible light is this?

11. History of Optics & Light Studies
Ibn Alhazen is considered the “Father of Optics.” He wrote the “Book of Optics”, which correctly explained and proved the modern theory of vision. His experiments included ones on lenses, mirrors, refraction, reflection, and the dispersion of light into its constituent colors. He studied the electromagnetic aspects of light, and argued that rays of light are streams of energy particles traveling in straight lines.
Ibn Alhazen
(965 – 1039)
“Father of Optics”

12. 6.2. Bohr model and quantum mechanical model
In 1913, Bohr developed a quantum model for the hydrogen atom.
Proposed a model that the electron
in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.
Niels Henrik David Bohr
Oct. 7, 1885 – Nov. 18, 1962
Danish Physicist
In atomic physics, the Bohr model depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus — similar in structure to the solar system, but with electrostatic forces providing attraction, rather than gravity. This was an improvement on the earlier cubic model (1902), the plum-pudding model (1904), the Saturnian model (1904), and the Rutherford model (1911). Since the Bohr model is a quantum-physics based modification of the Rutherford model, many sources combine the two, referring to the Rutherford-Bohr model.
The Nobel Prize in Physics 1922 for the investigation of the structure of atoms and of the radiation emiting from them.
Solar System Model has electrons moving around the nucleus.

13. How are the electrons arranged in a Bohr Model of an atom?
Electrons orbit the nucleus.
Electrons are arranged in specific pathways called energy levels.
There’s a fixed number of energy levels.
Each energy level is capable of holding certain number of electrons.
An electron cannot be found between energy levels.

14. Bohr’s Quantum Model for Hydrogen
The electron in hydrogen occupies discrete energy levels.
The atom does not radiate energy when the electron is in an energy level.
When an electron falls to a lower energy level, a quantum of radiation is released with energy equal to the difference between energy levels.
An electron can jump to higher energy levels if the atom absorbs a quantum of radiation with sufficient energy.

15. Bohr’s Explanation of Line Spectra
Neils Bohr developed a mathematical model that could explain the observation of atomic line spectra.
He proposed that electrons orbit the nucleus in certain “allowed” orbits - or energy levels. That is, the electron’s energy is QUANTIZED (not continuous).
Working on a model of single-electron atoms, Bohr derived an equation to calculate the energy of an electron in the nth orbit of such an atom:
where Z = atomic number (nuclear charge) and n = electron energy level
- Ve sign means energy of the electron bound to nucleus is lower than it would be if electron were at infinite distance.

16. The Hydrogen Line Spectrum

17. What’s Wrong with Bohr’s Model ?
Although it works great for single-electron atoms, Bohr’s model fails for atoms with 2 or more electrons!
It was a huge leap forward, but was fundamentally flawed.
Ultimately, the failure of Bohr’s model lay in the fact that he treated the electron as a charged particle orbiting the nucleus like a planet around the sun.
Electrons are more complicated …

18. 6. 3. QUANTUM NUMBERS
Total four quantum numbers were developed to better understand the movement and pathways of electrons in its orbital within an atom. Each quantum numbers indicates the trait of electron in atom.
1- The principal quantum number (n): It has integral values: 1, 2, 3…. The principal quantum number is related to the size and energy of the orbital.
An increase in n means higher energy, because the electron is less tightly bound to the nucleus, and the energy is less negative. Electrons with same value of n are said to be in the same “electron shell”

19. 2- The angular momentum quantum number (l)
2- The angular momentum quantum number (l). It has integral values from 0 to n ─1 for each value of n. This quantum number is related to the shape of atomic orbitals.

20. They range from 0 to “n–1”, Example: If n = 3, l can be 0, 1 or 2.
The value of l for a particular orbital is commonly assigned a letter: 0 is called s; 1 is called p; 2 is called d; 3 is called f.
3- The magnetic quantum number (m): It has integral values between +l and─l, including zero. The value of m is related to the orientation of the orbital in space relative to the other orbitals in the atom.

21. 4- Electron Spin Quantum Number (ms)
A fourth quantum number, ms describes electron spin (either +½ or –½)
Each electron in atom has a unique set of these four quantum numbers.
Electrons in orbitals with same n and l values are said to be in the same subshell.
Electrons with all three numbers the same, n, l , and ml, are in the same orbital.
A spinning negative charge creates a magnetic field. The direction of spin d determines the direction of the field.

22. Pauli Exclusion Principle
Electrons have negative charge and repel each other. How are the electrons in an atom distributed?
Wolfgang Pauli proposed that no two electrons in a given atom can be described by the same four quantum numbers!
The first three quantum numbers determine an orbital – therefore spins must be opposite!
Practical result is that each orbital can hold a maximum of two electrons, with opposite spins.

23. Aufbau Principle
In the ground state, the electrons occupy the lowest available energy levels. An atom is in an excited state if one or more electrons are in higher energy orbitals.
In atoms with more than one electron the lower energy orbitals get filled by electrons first!
This is the Aufbau principle, which is named after the German word which means “to build up”.
When one describes the locations of the electrons in an atom, start with the lowest energy electron and work up to the highest energy electron.

24. Hund’s Rule
If the degenerate orbitals are available, then electrons would like to be unpaired (separate) as long as possible to minimize electron-electron repulsions within the orbitals.
A set of orbitals is said to be “degenerate” if the orbitals possess the same energies. For example, all three “2p” orbitals on energy level 2 are degenerate. All five “3d” orbitals on energy level 3 are also degenerate.
When each degenerate orbital has one electron, electrons will then pair, spins opposed, until that sub-shell is filled.

25. Summary of Distribution Rules
Electrons distribute to lower energy levels until they are filled, before occupying higher levels. (Aufbau Principle)
Electrons will spread out as much as possible within a sub-shell. (Hund’s Rule)
Electrons will pair up, two to an orbital, spins opposed, until that sub-shell is filled. (Pauli Exclusion)
Every electron will have unique set of 4 quantum numbers.

26. Particle Theory of Light
In 1900, Max Planck turned the world of physics on its head by presenting the particle theory of light.
Planck proposed that light is composed of particles (quanta) each carrying a fixed amount of energy.
The amount of energy per quantum is directly proportional to the frequency of the light.
This hypothesis was later extended by Albert Einstein. Einstein presented light as small discrete particles of energy called photons.
Max Karl Ernst Ludwig Planck
(April 23, 1858 – October 4, 1947)
German Physicist
The Nobel Prize in Physics 1918 for the discovery of energy quanta.

27. Wave-Particle Dual Nature of Electrons
Einstein’s famous equation, E = mc2, suggested that energy and mass are related – matter can be converted directly into energy.
Louis de Broglie (1924) made the leap that if light can behave as wave/particles, then so can matter!
Wavelength = h/p p=momentum
An atomic model must make use of the wave-nature of electrons to be complete! This is what Bohr had missed .

29. 6.4. Electron Configurations
Electrons orbitals are defined by their quantum numbers, n, l and ml .
Each electron in an atom has a unique set of 4 quantum numbers.
No two electrons can have the same “address”, i.e., the same 4 quantum #’s
Rules define how multiple electrons will be distributed among the possible energy levels

30. Examples of Electron Configurations
Helium has 2 electrons in the 1s orbital, He: 1s2
Carbon has 6 electrons, C: 1s2 2s2 2p2
Calcium has 20 electrons, Ca: 1s2 2s2 2p6 3s2 3p6 4s2
Calcium cation, Ca2+: 1s2 2s2 2p6 3s2 3p6
Noble Gas “shorthand” Notation.
Ca: 1s2 2s2 2p6 3s2 3p6 4s2
Ca: [Ar] 4s2
Ar: 1s2 2s2 2p6 3s2 3p6

31. Orbital Diagrams

32. 6.5. s, p and d Orbital shapes

33. 6.6 Electron arrangements in mono-atomic ions.
Certain monatomic ions form when atoms gain or lose electrons to achieve a stable electron configuration for their highest energy electrons. 

34. 6.7.Periodicity in Periodic Table
“Periodicity” refers to similarities in behavior and reactivity due to similar outer shell electron configurations.
All the Alkali Metals have one unpaired valence electron; all the Noble Gases have completely filled sub-shells.
We will examine periodic trends in atomic radius, ionization energy, electronegativity, and electron affinity.

35. Atomic Size Atomic radii increase within a group (column)
as the principal
quantum number of
Outermost shell
increases.
Atomic size decreases
across a row (period)
from Left to Right,
because the effective
nuclear charge

36. IONIZATION ENERGY remove an electron from the ground State of atom
IONIZATION ENERGY: Minimum energy required to
remove an electron from the ground State of atom
(molecule) in the gas phase.
M (g) + h  M+ + e
Sign of the ionization energy is always positive, for
example, energy is required for ionization to occur.

37. Electron Affinity
1) Electron affinity is the energy change which occurs when an electron is accepted by an atom in the gaseous state.
A(g) + e A-(g)

38. Electronegativity: The ability of an atom in a bond to pull on the electron. (Linus Pauling)
When the atoms are the same they pull on the electrons equally. Example, H-H.
When the atoms are different, the atoms pull on the electrons unevenly. Example, HCl.

39. Electronegativities of Some Elements
Element Pauling scale F
Cl 3.0 O N S C H
Na 0.9
Cs 0.7
Most electronegative element is F which is (EN 4.0).
Least electronegative stable element is Cs (EN 0.7).
Left to right Electronegativity increases
Up to down Electronegativity decreases

40. Negative Ions Positive ions Positive ions are always smaller that
the neutral atom.
Loss of outer shell
electrons.
Negative Ions
Negative ions are
Always larger than
The neutral atom.
Gaining electrons.

41. Up to down Electronegativity decreases
Left to right Electronegativity increases