Chemical Bonding and Intermolecular Forces Notes

Ionic bonding occurs when there's a significant difference in electronegativity between two atoms, typically greater than 1.8.
It usually happens between reactive metals and reactive non-metals.
During ionic bond formation:
- Metals lose electrons and become positively charged ions (cations).
- Non-metals gain electrons and become negatively charged ions (anions).
Oppositely charged ions attract each other through strong electrostatic forces, forming the ionic bond.
The number of electrons lost or gained during ionic bonding is the same as the element's valency.
- For Example:
  - In sodium chloride (NaCl) formation, sodium (Na) transfers one electron to chlorine (Cl), creating $Na^+$ and $Cl^-$ .
  - The electronegativity difference between Na (0.9) and Cl (3.0) is 2.1.
  - $Na^+$ and $Cl^-$ achieve the electronic configurations of noble gases neon (Ne) and argon (Ar) respectively but carry ionic charges.
  - Magnesium oxide (MgO) forms because the electronegativity difference between Mg (1.2) and O (3.5) is 2.3, which greater than 1.8.
  - Mg loses two electrons, and oxygen gains two electrons.

Lowest energy main shell and subshell that electrons occupy.
Directions of electron spins in the same orbital of an atom.
Number of orbitals in the s, p, and d subshells, and their shapes.
Electronic configuration, valence electrons, and Lewis symbol for elements: $^{13}Al$ , $^{15}P$ , $^{16}S$ , $^{19}K$ , $^{35}Br$ , and $^{36}Kr$ .
Definition of electronegativity.
Factors determining the type of chemical bonding.
Classification of chemical bonding types and the strongest force among them.
Intermolecular forces.

Also known as electrovalent bonding.
Results from electron transfer from an atom with low ionization energy to one with high electron affinity.
Involves electrostatic attraction between cations and anions.

Values for various elements are provided such as Hydrogen (2.1), Lithium (1.0), Beryllium (1.5), etc.
Groups are classified (1A, 2A, 3A, etc.) with corresponding elements and electronegativity values.

Most bonds are not purely ionic or covalent but possess both characters.
London dispersion forces are the weakest intermolecular forces found in all substances.
Electronic structures and Lewis structures of elements like Carbon, Sodium, and Chlorine.

Electron affinity (EA) is the energy released when an electron is added to a gaseous atom.
Ionization energy (IE) is the energy required to remove an electron from a gaseous atom.
Elements with high EA readily gain electrons to form anions.
Elements with low IE readily lose electrons to form cations.

Weak forces between molecules.
- Hydrogen bonding.
- Van der Waals forces:
  - Ion-dipole interaction.
  - Dipole-dipole interaction.
  - London dispersion forces.
Impact of chemical bonding and intermolecular forces on structures and physical properties.
Ionic bonds are generally stronger than covalent bonds:
- Ionic\ bond > Covalent\ bond
Hydrogen bonds are the strongest intermolecular force, followed by van der Waals forces:
- Hydrogen\ bond > van\ der\ Waals\ forces\ (Ion-dipole > dipole-dipole > London\ dispersion)

Depends on elements involved (metals or non-metals).
- Ionic bonding (Metal + Non-metal).
- Covalent bonding (Non-metal + Non-metal).
- Metallic bonding (In metals).
Forces between molecules are weak compared to forces within molecules.

Higher first ionization energy generally corresponds to greater electronegativity.
Greater electronegativity means stronger attraction to electrons.

Elements of the second period may have exceptions when there are an odd number of valence electrons, too few, or too many valence electrons.
The octet rule does not apply to d-block elements.

Every orbital in a subshell is singly occupied before any one orbital is doubly occupied with opposite spins.
All electrons in singly occupied orbitals have the same spin.

Electrons are filled in the lower energy atomic orbitals before filling higher energy ones.
Pauli's exclusion principle states that no more than two electrons can occupy the same orbital, and two electrons in the same orbital must have opposite spins.

Aufbau principle: Electrons first fill lower energy atomic orbitals.
Order of filling orbitals: 1s 2s 2p 3s 3p 4s 3d 4p,…
Electron spin: Each electron spins on an axis with two possible directions (↑ or ↓).
Electrons in the same orbital must have opposite spins.
Pauli's exclusion principle: No more than two electrons per orbital, with opposite spins.

Aufbau principle, Pauli's exclusion principle, and Hund's rule govern electron filling in atomic orbitals.