VSEPR Theory and Molecular Geometry

Module 2: Impact of Molecular Structure on Compound Properties

• This module examines how molecular structure influences the properties of compounds.

2.1 Lewis Structure and Molecular Geometry

• Lewis structures (also called Lewis diagrams or electron-dot diagrams) are simplified visual representations of valence electrons in atoms or molecules.
• They depict valence electrons as dots or crosses, and the maximum number of electrons is 8, following the octet rule.
• Understanding Lewis structures is helpful in predicting the geometry and properties of substances.

Lewis Structure of Neutral Atoms

• Valence electrons are shown as dots or crosses around the elemental symbol.
• The maximum is 2 electrons per side forming a total of 8 electrons, following the octet rule.
• Examples:
  • Sodium (Na) with one valence electron: Na.
  • Chlorine (Cl) with 7 valence electrons: :Cl:
  • Neon (Ne) with 8 valence electrons: :Ne:

Learning Check

• Students are asked to draw Lewis structures for potassium (K), carbon (C), and oxygen (O), and identify their group and valence electron numbers.
  • Potassium (K): Group 1, 1 valence electron
  • Carbon (C): Group 14, 4 valence electrons
  • Oxygen (O): Group 16, 6 valence electrons

Lewis Structure for Ionic Compounds

• A simple ion is an atom that gains or loses electrons to acquire a positive or negative charge.
• Ionic bonding occurs when valence electrons transfer from one atom to another.
• The atom that loses electrons (cation) gains a positive charge and its Lewis structure has zero dots.
• Example: Sodium (Na) losing one electron to form Na⁺.
• Magnesium (Mg) losing two electrons to form Mg²⁺.

Examples

• Magnesium (Mg) has two valence electrons and loses them to become Mg²⁺.
• Nitrogen (N) has five valence electrons and gains three to become N³⁻.

Lewis Structure for Ionic Compounds

• Cations (positive ions) are on the left, and anions (negative) are on the right.
• The charges and subscripts indicate the proportion in the compound.
• Examples include NaCl, MgCl₂, and K₂O; their Lewis structures are; Na⁺ [:Cl⁻], Mg²⁺ [::] [:Cl⁻]₂, 2K⁺ [:O²⁻].

Lattice Structure of Ionic Compounds

• Lewis structures represent electron transfer, but do not show the actual arrangement of ions in real-life ionic compounds.
• Ions arrange in a lattice structure due to attractive forces between opposite charges.
• The structure is three-dimensional and involves a repeating pattern of positive and negative ions.
• The ratio of ions is constant, but the lattice geometry varies based on the size of the ions.
• The formula of the ionic compound represents the ratio of ions present in the solid crystal lattice.
• Sodium chloride (NaCl) exemplifies a 1:1 ratio of Na⁺ and Cl⁻ ions.

Lewis Structure of Covalent Molecules

• Covalent bonds form when non-metal atoms share electrons to achieve a stable octet.
• Single, double, or triple bonds form based on the number of shared electrons between atoms involved.
• The Lewis structure shows all valence electrons in a molecule, including shared electrons (bonding pairs) and unshared electrons (lone pairs).
• Non-bonding electrons influence molecular shape and properties.

Calculating Lewis Structures for Covalent Molecules

• Determine the total number of valence electrons.
• Draw the skeletal structure, placing atoms in their typical positions within molecules.
• Use lines for bonds, crosses or dots for unbonded electrons, to ensure each atom has an octet.
• Add any unbonded electron pairs to complete the octets of atoms involved.
• Form double or triple bonds to fill octets if needed.
• Example: Calculating the Lewis structure for CCl₄.

Molecular Geometry of Covalent Compounds

• Molecular geometry is the 3D arrangement of atoms in a molecule.
• This arrangement affects properties such as polarity, reactivity, and phase.
• The position of atoms stems from interactions (attraction and repulsion) among atoms and electrons.

Valence Shell Electron Pairs Repulsion Theory(VSEPR)

• VSEPR is a theory for predicting molecular geometry.
• It's based on Valence electrons repelling each other, causing atoms in a molecule to arrange themselves to minimize these repulsions.
• Molecules will attempt to maximize the distances between electrons in the structure.

Types of Molecular Geometry

• Describes different molecular shapes depending on the number of atoms.
• The shapes of molecules influence many properties including polarity and the strength of intermolecular forces.
• Examples are linear, trigonal planar, trigonal pyramidal, bent, and tetrahedral.

2.2 Polarity of Molecules

• Polarity occurs when electrons aren't evenly shared in a chemical bond.
• Electronegativity differences between atoms in a molecule cause electrons to spend more time around one atom than another.
• This makes one end of the bond slightly negative and the other slightly positive (dipole).
• A molecule with polar bonds may be nonpolar if the bonds are symmetrically arranged, canceling out the dipole moments.
• Some examples of polar molecules are water (H₂O), ammonia (NH₃), and hydrogen bromide (HBr).
• Nonpolar examples are oxygen (O₂), carbon dioxide (CO₂), and methane (CH₄).

Polarity of a Bond

• Types of bonds based on electronegativity difference:
  • Ionic: Complete electron transfer; large electronegativity difference.
  • Polar covalent: Unequal electron sharing; moderate electronegativity difference.
  • Nonpolar covalent: Equal electron sharing; small electronegativity difference.

Polarity of a Molecule

• Polar molecules result when the molecule's polar bonds are not symmetrically arranged, leading to a net dipole moment.
• Nonpolar molecules occur when polar bonds are symmetrically arranged, thus canceling the net dipole moment.

Examples of Polar and Nonpolar Molecules

• Examples include water, carbon dioxide, methane, etc.

2.3 Intermolecular Forces of Attraction (IMFA)

• Intermolecular forces are attractive forces between molecules, not within them.
• They are weaker than chemical bonds.
• Intermolecular forces determine bulk properties of substances, like boiling and melting points.

Types of Intermolecular Forces

• Dispersion forces (London forces): Weakest forces that occur in all molecules, including nonpolar.
• Dipole-dipole forces: Moderate forces that occur between polar molecules.
• Hydrogen bonds: Strongest forces, occur between molecules with H bonded to N, O, or F.

Effect of Intermolecular Forces on Boiling and Freezing Points

• Stronger intermolecular forces lead to higher boiling and freezing points.
• This is due to the greater energy required to overcome these forces to change the state of matter.

Summary of Intermolecular Forces

• Intermolecular forces are categorized by strength (London, dipole-dipole, hydrogen bonds).
• They are crucial for determining substances' physical properties like boiling and freezing points and other physical properties of matter.

Solubility

• Solubility is a physical property describing a substance's ability to dissolve in another substance (solvent).
• Water is a universal solvent because it can dissolve many polar substances due to its polarity and ability to form hydrogen bonds.

Summary

• The module's key takeaways center on the role of molecular structure in determining the properties of chemical compounds. These properties are dependent on the interactions within and between molecules. Electron configuration and bond type are important factors.