1. Atomic Structure and Periodicity
CHAPTER-7
Atomic Structure and Periodicity

2. EXAMPLE: Very less wavelength (negligible)
Wavelength in electromagnetic
radiation range
Very less wavelength (negligible)

3. 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 .

4. 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”

5. 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.

6. 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 + and─, including zero. The value of m is related to the orientation of the orbital in space relative to the other orbitals in the atom.

8. s, p and d Orbital shapes

9. 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 subs hell.
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.

10. The first four levels of orbitals in the hydrogen atom are listed with their quantum numbers in Table 7.2. Note that each set of orbitals with a given value of (sometimes called a sub shell) is designated by giving the value of n and the letter for .
Thus an orbital where n 2 and 1 is symbolized as 2p. There are three 2p orbitals, which have different orientations in space. We will describe these orbitals in the next section.
For principal quantum level n 5, determine the number of allowed subs hells (different values of ), and give the designation of each.

11. 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.

12. 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

13. 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.

14. 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.

15. 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.

16. 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

17. Orbital Diagrams

18. 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.

19. 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

20. 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, it requires energy for ionization to occur.

21. 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)

22. 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.

23. 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

24. 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.

25. THE END