1. The Development of the Periodic Table and Electron Configuration

2. The Development of the Periodic Table
Who is given the most credit for the development of the periodic table?
Dmitri Mendeleev – his table, an expression of periodic law, had predictive power. However, it did not explain these predictions. (Difference between a law (summarizes behavior) and a theory (explains behavior)).
The modern periodic table, which arranges the elements according to their atomic number (number or protons), is vitally important to understand the properties of macroscopic substances.

3. The Periodic Table (Continued)
For the AP exam, you will need to be able to make predictions about elements (and compounds) based on the elements position on the table.
Periodic properties: a property that is predictable based on the elements position on the table.
Can you name some of the periodic properties?
Atomic radius, ionization energy, electronegativity, electron affinity, metallic characteristics.

4. Electron Configuration
Quantum-mechanical theory:
Describes the behavior of electrons in atoms.
Helps us understand and describe chemical behavior
Electrons configuration – provide a method of describing the electron distribution in the atom.
Ground state electron configuration – the electron arrangement of an atom or ion in the lowest energy level (most stable configuration)

5. Single electron vs. Multi-electron atoms
Unlike the shell model, the quantum mechanical model takes sub-level splitting (for multi-electron atoms) into consideration. This means that sub-levels within a principle energy levels have different energies.
In order to understand why sublevels split, we must understand three key concepts associated with the energy of an electron in the vicinity of a nucleus:
1.) Coulombs law: describes the interactions between charged particles.
2.) Shielding : describes how one electron can shield another electron from the
full charge of the nucleus.
3.) Penetration: describes how one atomic orbital can overlap spatially with
another, thus penetrating into a region that is close to the nucleus (and
therefore less shielded from nuclear charge).

6. Coulombs Law
Coulombs law: the potential energy (E) of two charged particles, depends on their charges (q1 and q2) and their separation (r)
q1q2
E = r
For like charges, the potential energy (E) is positive and decreases as particles get farther apart (as r increases)
For opposite charges, the PE is negative and becomes more negative as the particles get closer together (as r decreases)
The magnitude of the interaction between charged particles increases as the charges of the particles increases. Consequently, an electron with a charge of 1- is more strongly attracted to a nucleus with a 2+ charge than it would be to a nucleus with a charge of 1+

7. Let’s Try a Practice Problem
According to Coulomb’s law, what happens to the potential energy of two oppositely charges particles as they get closer together?
(a) Their potential energy decreases.
(b) Their potential energy increases.
(c) Their potential energy does not change.
(a) Their potential energy decreases, because as the distance between the two oppositely charged particles become closer, their potential energy becomes more negative. This is a more stable interaction.

8. Shielding
For multielectron atoms, they experience both the positive charge from the nucleus (attractive force) and the negative charge from other electrons (repulsive force).
Effective nuclear charge (Zeff): the attractive charge of the nucleus + the repulsive charge from electron shielding (or screening). Zeff = nuclear charge + core electron charge.
The Zeff of lithium: (3+) + (2-) = 1+
The 3+ is the nuclear charge, and the 2- is the overall charge of the two electrons that occupy the 1s sublevel that are shielding lithium’s valence electron from the full nuclear charge.

9. Penetration
The outer electrons, can penetrate into a region occupied by the inner electrons, and when it does it experience a greater nuclear charge and therefore (according to Coulomb’s law) a lower energy.
Penetration of electrons, explains why sub-levels that may be thought to be higher in energy are actually lower in energy. For example, 4s is lower in energy that 3d.

10. Aufbau Principle
Electrons will enter the lowest energy levels first. (F.Y.I. aufbau is a German word that means “build up”)
Copy down the diagram that I put on the board. This diagram will help you write electron configurations in sub-level notation.
s sub-levels = 1 orbital
p sub-levels = 3 orbitals
d sub-levels = 5 orbitals
f sub-levels = 7 orbitals
Each orbital can only hold a total of two electrons, and when they do, the electrons must have opposite spins: Pauli’s exclusion principle.

11. Writing Electron Configurations
The electron configuration (in sub-level notation) for sodium can be written in one of two ways:
Na 1s22s22p63s1 or
Na [Ne] 3s1

12. Let’s Try a Practice Problem
Write the electron configurations for each element. Cl Si Sr O
(a) Cl 1s22s22p63s23p5 or [Ne] 3s23p5
(b) Si 1s22s22p63s23p2 or [Ne] 3s23p2
(c) Sr 1s22s22p63s23p64s23d104p65s2 or [Kr] 5s2
(d) O 1s22s22p or [He] 2s22p4

13. pgs #’s 42, 46, 56, 58, & 60 Read pgs