Topic 2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

 

2.2 Atomic orbitals: s, p and d

 

2.3 Electronic configuration

 

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

 

2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

Candidates should be able to:

(a) explain the formation of the emission line spectrum of atomic hydrogen in the Lyman and Balmer series using Bohr’s Atomic Model.

2.2 Atomic orbitals: s, p and d

Candidates should be able to:

(a) deduce the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3, including the 4s orbitals;
(b) describe the shape of the s and p orbitals.

2.3 Electronic configuration

Candidates should be able to:

(a) predict the electronic configuration of atoms and ions given the proton number (and charge);
(b) define and apply Aufbau principle, Hund’s rule and Pauli Exclusion Principle.

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

Candidates should be able to:

(a) identify the position of the elements in the Periodic Table as
(i) block s, with valence shell configurations s1 and s2,
(ii) block p, with valence shell configurations from s2p1 to s2p6,
(iii) block d, with valence shell configurations from d1s2 to d10s2;
(b) identify the position of elements in block f of the Periodic Table.

 

2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

Candidates should be able to:

(a) explain the formation of the emission line spectrum of atomic hydrogen in the Lyman and Balmer series using Bohr’s Atomic Model.

2.2 Atomic orbitals: s, p and d

Candidates should be able to:

(a) deduce the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3, including the 4s orbitals;

(b) describe the shape of the s and p orbitals.

2.3 Electronic configuration

Candidates should be able to:

(a) predict the electronic configuration of atoms and ions given the proton number (and charge);

(b) define and apply Aufbau principle, Hund’s rule and Pauli Exclusion Principle.

Aufbau principle

The aufbau principle (also called the building-upprinciple or the aufbau rule) states that in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels. For example, the 1s shell is filled before the 2s subshell is occupied.

Hund’s rule

The German Physicist, Fredrich Hund, formulated it around 1927, which allows to easily draw electron energy diagrams for a ground state multi- electron atom. Hund’s rule states: – Every orbital in a sublevel is occupied before they start doubling up -All electron in a single occupied orbital have the same spin (in order to maximize the total spin).
Hund´s rule is sometimes used to describe people getting seated on a bus. Everyone will first get their own seat, sitting alone. Not until there are no other seats people will start sitting next to a stranger.

Pauli Exclusion Principle

in an atom or molecule, no two electrons can have the same four electronic quantum numbers. As an orbital can contain a maximum of only two electrons, the two electrons must have opposing spins. This means if one is assigned an up-spin ( +1/2), the other must be down-spin (-1/2).

Electron filling in atomic orbitals following Aufbau principle [or building-up principle], Pauli exclusion principle [or Antisymmetry principle] and Hund’s rule of maximum multiplicity.
For copper

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

Candidates should be able to:

(a) identify the position of the elements in the Periodic Table as
(i) block s, with valence shell configurations s1 and s2,
(ii) block p, with valence shell configurations from s2p1 to s2p6,
(iii) block d, with valence shell configurations from d1s2 to d10s2;

(b) identify the position of elements in block f of the Periodic Table.







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