HomeClass Notes1,200,000CA670,000

CH110 Lecture Notes - Lecture 5: Pauli Exclusion Principle, Aufbau Principle, Electron Configuration

11 views2 pages

turquoisewasp140

18 Nov 2020

School

Department

Course

Professor

Daniel Santorelli

For unlimited access to Class Notes, a Class+ subscription is required.

Document Summary

Uses [ ] symbol to represent a noble gas configuration. The [ ] symbol represents the core electron configuration only. The valance electron configuration is written as usual (valence electrons = electrons in the highest n shell) The periodic table is arranged in order of increasing electrons. The orbitals of the atoms are filled with these electrons. Each orbital has a maximum capacity of 2 electrons (pauli exclusion principle) The atoms in the periodic table are in their ground states (most stable) Conditions for electron arrangement in atomic ground states: each electron in the atom must occupy the most stable available orbital, no two electrons can have identical quantum numbers (pauli exclusion principle, orbital capacities: Each row is labelled with the highest principle quantum number. Each column (group) is labelled with its group number. The actinides and lanthanides (grey) do not have column numbers. Each element shows its least stable occupied orbitals.

Related Documents

CH110 Lecture Notes - Lecture 4: Pauli Exclusion Principle, Periodic Trends, Atomic Number

turquoisewasp140

CH110 Lecture Notes - Lecture 2: Principal Quantum Number

turquoisewasp140

CH110 Lecture 8: Ch110 lecture 8

turquoisewasp140

Related Questions

I need the answers with explanation

Slater’s rules are a way to estimate the effective nuclear charge experienced by an electron. In this approach, the “shielding constant,” S, is calculated. The effective nuclear charge is then the difference between S and the atomic number, Z.

The shielding constant, S, is calculated using the following rules:

(1) The electrons of an atom are grouped as follows: (1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d), and so on.

(2) Electrons in higher groups (to the right) do not shield those in the lower groups.

(3) For ns and np valence electrons

a) Electrons in the same ns, np group contribute 0.35 (for 1s 0.30 works better).

b) Electrons in n − 1 groups contribute 0.85.

c) Electrons in n − 2 groups (and lower) contribute 1.00.

(4) For nd and nf electrons, electrons in the same nd or nf group contribute 0.35, and those in groups to the left contribute 1.00.

As an example, let us calculate Z* for the outermost electron of oxygen:

Here is the calculation for a d electron in Ni:

and for an s electron Ni:

(Here 3s, 3p, and 3d electrons are in the (n − 1) groups.)

a) Calculate Z* for F and Ne. Relate the Z* values for O, F, and Ne to their relative atomic radii and ionization energies.

b) Calculate Z* for one of the 3d electrons of Mn, and compare this with Z* for one of the 4s electrons of the element. Do the Z* values give us some insight into the ionization of Mn to give the cation?