Unit 5: Chemical Bonding and Shapes of Molecules

Valence Shell: The outermost shell of an atom that contains the valence electrons, which are involved in chemical bonding.
Valence Electrons: Electrons in the outermost shell that determine an element's bonding behavior.
Octet Theory: Atoms tend to gain, lose, or share electrons to achieve a stable configuration of 8 electrons in their valence shell (similar to noble gases).

Ionic Bond: A type of chemical bond formed when one atom donates an electron to another atom, resulting in the formation of positively charged cations and negatively charged anions.
Properties:
- High melting and boiling points.
- Good electrical conductors in molten or aqueous states.
- Soluble in polar solvents like water.
- Hard and brittle solids.

Covalent Bond: A bond formed when atoms share pairs of electrons.
Coordinate Covalent Bond (Dative Bond): A type of covalent bond where both shared electrons come from the same atom.
- Example: $\text{NH}_4^+$ (Ammonium ion), where nitrogen donates a lone pair to bond with a hydrogen ion.

Low melting and boiling points.
Poor electrical conductivity (except in molten state or when dissolved in water).
Insoluble in polar solvents, soluble in non-polar solvents.
Usually softer and more flexible compared to ionic compounds.

Lewis Structures represent the arrangement of valence electrons around atoms in a molecule.
Examples:
- $\text{H}_2\text{O}$ (Water): Oxygen shares two pairs of electrons with hydrogen atoms.
- $\text{CO}_2$ (Carbon Dioxide): Carbon forms two double bonds with oxygen atoms.

Resonance: Occurs when a molecule can be represented by two or more valid Lewis structures.
- Example: Ozone (O₃) has two resonance structures where the double bond is distributed between the oxygen atoms.

Valence Shell Electron Pair Repulsion (VSEPR) Theory: Electron pairs around a central atom repel each other, determining the molecule’s shape.
Shapes of some simple molecules:
- BeF₂ (Beryllium Fluoride): Linear (180° bond angle).
- BF₃ (Boron Trifluoride): Trigonal planar (120° bond angle).
- CH₄ (Methane): Tetrahedral (109.5° bond angle).
- CH₃Cl (Methyl Chloride): Tetrahedral.
- PCl₅ (Phosphorus Pentachloride): Trigonal bipyramidal (120° and 90° bond angles).
- SF₆ (Sulfur Hexafluoride): Octahedral (90° bond angle).
- H₂O (Water): Bent (104.5° bond angle).
- NH₃ (Ammonia): Trigonal pyramidal (107° bond angle).
- CO₂ (Carbon Dioxide): Linear (180° bond angle).
- H₂S (Hydrogen Sulfide): Bent.
- PH₃ (Phosphine): Trigonal pyramidal.

Valence Bond Theory: Describes covalent bond formation as the overlap of atomic orbitals. Electrons in overlapping orbitals form a bond by sharing electron pairs.

Hybridization: The concept of mixing atomic orbitals to form new hybrid orbitals.
- sp Hybridization: Linear arrangement (e.g., BeCl₂).
- sp² Hybridization: Trigonal planar arrangement (e.g., BF₃).
- sp³ Hybridization: Tetrahedral arrangement (e.g., CH₄).

Bond Length: The distance between the nuclei of two bonded atoms.
Ionic Character: The degree to which a bond is ionic (greater difference in electronegativity leads to higher ionic character).
Dipole Moment: A measure of the separation of positive and negative charges in a molecule; determines the molecule's polarity.

Van der Waals Forces: Weak intermolecular forces that include:
- Dispersion Forces: Caused by temporary dipoles.
- Dipole-Dipole Interactions: Between polar molecules.
Molecular Solids: Held together by these weak forces, resulting in lower melting and boiling points.

Hydrogen Bonding: A strong type of dipole-dipole interaction involving hydrogen and a highly electronegative atom (N, O, or F).
- Applications: Explains the high boiling point of water and the structure of DNA.

Metallic Bonding: The bonding in metals where valence electrons are delocalized and shared among a lattice of metal atoms.
Properties of Metallic Solids:
- Good electrical and thermal conductors.
- Malleable and ductile due to the mobility of electrons.
- High melting and boiling points.

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