Structure of Atoms

Bohr theory

Presented by Neil Bohr.

1. Electrons of an atom can have only very distinct energy values. Therefore, the electrons are restricted to specific “energy levels” or “stationary states”.
2. Electrons in these energy levels rotate about the nucleus in fixed orbits, without radiating or absorbing any energy.
3. When an electron receives energy, it moves from a lower to a higher energy level.
4. Emission of radiant energy (electromagnetic radiation) is due to the movement of the electrons from the higher level to the former lower level of energy.

Bohr Theory can’t be used to describe the energy characteristics of atoms containing many electrons. Consequently, in the Modern Atomic Theory, four quantum numbers were introduced to describe the energy levels.

Quantum numbers

Principal quantum number

The number given in Bohr’s original stationary states corresponds to the principal quantum number ($n$). Relates to the distance between the nucleus and the principal energy levels. Maximum is $7$.

Secondary quantum number

Represents the various secondary sub-levels within the main energy level. Relates to the shape of the electron cloud. Denoted by $l$. Values goes from $0$ to $n-1$.

Magnetic quantum number

Represents the direction of maximum extension in space of the electron cloud in the sub shells p, d, and f. which has the dumb - bell shape. Denoted by $m$. Values goes from $-l$ to $l$ (in total $2l + 1$ different values).

Spin quantum number

Represents the electron spin. $+\frac{1}{2}$ for clockwise and $-\frac{1}{2}$ for counterclockwise spin.

Principles

Pauli’s exclusion principle

No two electrons in a single atom can have all four-quantum numbers the same.

Aufbau Principle

Sub-levels with the lowest energy are filled before those with higher energy.

Hund’s Rule

The sublevels p, d, and f are occupied by single electrons before any pairing of electrons with opposed spins take place.

Atomic bonding

Primary (strong) bonds

Form when valence electrons are present (outer shell not completely filled).

Ionic

Forms between highly electropositive elements (metals) and highly electronegative elements (non-metals). Non-directional.

Covalent

Forms within electronegative elements that are close to each other in the periodic table. Directional.

Metallic

Forms in metals. Arrangement of positive ion cores in a sea of electrons. Non-directional.

Secondary (weak) bonds

Form when there are no valence electrons (outer shell filled completely).

Van Der Waal’s

Forms between inert gases and between symmetric molecules. Forms because of the dipolarity caused inside molecules. Non-directional.

Hydrogen

Forms between polar covalently bonded molecules containing $\ce{H}$. Polarity is due to the asymmetry of the molecule. Directional. $$

Properties of bonds

Property	Order
Hardness, Melting point, Crystalline percentage	Ceramics > Metals > Polymers
Directionality	Covalent > Ionic > Metallic
Brittleness	Ionic > Covalent > Metallic