VSEPR Theory, Polarity, and Molecular Structure

• Lewis Structures depict the arrangement of atoms and bonds in a molecule.
  • Count the total valence electrons of all atoms in the molecule.
  • Place the least electronegative atom in the center (typically the least abundant).
  • Arrange remaining atoms symmetrically around the central atom.
  • Follow the octet rule (8 valence electrons) and duet rule (2 valence electrons for hydrogen).
    • Exceptions: Beryllium (Be) prefers 4 valence electrons; Boron (B) prefers 6 valence electrons; elements in rows 3-7 can exceed the octet rule.
  • Check formal charge to optimize the Lewis structure.
    • Formal Charge = (# of valence electrons for an atom) - (# of electrons assigned to the atom in Lewis Structure).
    • Structures with lower formal charges are more stable, and exceeding the octet rule can sometimes improve formal charge.

Molecular Shapes

• Molecular shape is determined by the arrangement of atoms around the central atom.
• Bond angles are the angles between two bonds.
• Bonding domains are areas where bonds exist (single, double, or triple bonds count as one domain each).
• Unshared pairs (lone pairs) are electrons not involved in bonding.
• Valence Shell Electron Pair Repulsion (VSEPR) Theory: Valence electrons around the central atom arrange themselves to minimize repulsion, influencing molecular shape.
  • Molecules with resonance have the same shape for all resonance structures.
  • For molecules with multiple central atoms, determine the shape around each central atom.

Specific Molecular Structures

• Linear: atoms arranged in a line, 180° bond angle. Always the shape of molecules made of only 2 atoms or with 2 bonding domains around the central atom.
• Trigonal Planar: 2D triangle shape, 120° bond angle, 3 bonding domains around the central atom.
• Tetrahedral: 3D shape with 4 bonding domains around the central atom, 109.5° bond angle, 4 bonds and no lone pairs.
• Trigonal Pyramidal: 3D shape with triangular sides, 107° bond angle, 1 lone pair and 3 bonding domains.
• Bent: bonding domains are bent away from lone pairs, approximately 104.5° bond angle, 1 or 2 lone pairs and 2 bonding domains. This shape is always polar.

Polarity

• Bond Polarity: Determined by the difference in electronegativity between two atoms.
  • ≤ 0.4: Nonpolar covalent bond

  • 0.4 and ≤ 1.7: Polar covalent bond

  • 1.7: Ionic bond

• Molecular Polarity: Depends on bond polarity and molecular symmetry.
  • Bonds must be distributed in a non-symmetrical arrangement around the central atom.
  • If more than one central atom, assess the symmetry of the overall molecule.
  • Tips: Bent and trigonal pyramidal shapes are always polar; linear, tetrahedral, and trigonal planar shapes with the same atom around the central atom are nonpolar; different atoms around the central atom result in a polar molecule. Trigonal pyramidal is polar because the lone pair pushes down.

Hybrid Orbitals

• Hybrid Orbitals: Atomic orbitals formed by combining two or more different orbitals, preparing for covalent bonds.
• Electron Domains: The sum of bonding domains and lone pairs around the central atoms.
  • sp: one s orbital and one p orbital = 2 sp hybrid orbitals (linear arrangement). Examples: BeCl2, C2H2.
  • sp2: one s orbital and two p orbitals = 3 sp2 hybrid orbitals (trigonal planar arrangement). Examples: BF3, C2H4.
  • sp3: one s orbital and three p orbitals = 4 sp3 hybrid orbitals (tetrahedral geometry). Examples: CH4, NH3, H2O.
• Sigma (σ) Bond: A shared pair of electrons between 2 atoms; present in single bonds. Form from hybrid orbitals.
• Pi (π) Bond: A bond above or below a sigma bond; present in double bonds (one sigma, one pi).

Intermolecular Forces

• Intermolecular Forces: Attractive forces between molecules, affecting whether a substance is a solid, liquid, or gas.
  • Affects physical change.
• Dipole-Dipole Forces: Attractive forces between polar molecules.
  • Partial positive end of one molecule attracts the partial negative end of another.
• Hydrogen Bonding: Occurs between a hydrogen atom and a highly electronegative atom (F, O, or N).
  • Stronger than dipole-dipole forces because F, O, and N are highly electronegative, and hydrogen is weakly electronegative, creating strong partial charges.
• London Dispersion Forces (LDF): Occur between all molecules, including noble gases and nonpolar molecules.
  • Result from temporary, uneven distribution of electrons, creating temporary dipoles.
  • Strength increases with the number of electrons and molecular size.

Boiling Point

• Boiling Point: The temperature at which a substance transitions from a liquid to a gas.
• Higher boiling point indicates stronger intermolecular forces.
• To determine the substance with the highest boiling point:
  • The stronger the intermolecular force, the higher the boiling point.
  • If the same type of force, the substance with more electrons has a higher boiling point.

Intramolecular Forces

• Intramolecular Forces: Forces within a molecule (e.g., covalent bonds).
  • Stronger than intermolecular forces.
  • Relate to chemical changes.

Types of Solids

• Characterized by physical properties.
  • Liquid: Boiling point, volatility, vapor pressure.
  • Solids: Electrical conductivity, hardness, melting point.
• Crystalline Solid: Particles arranged in a repeating pattern.

Crystalline Solid Types

• Ionic Solids: Ions held together by ionic bonds, arranged in a crystal lattice.
  • Properties: Hard, high melting point, conducts electricity when dissolved in water.
• Molecular Solids: Molecules held together by intermolecular forces.
  • Properties: Not as hard as ionic solids, low melting point, generally do not conduct electricity, can dissolve. Nonpolar molecular solids can sublime.
• Covalent Solids/Atomic Solids: Atoms covalently bonded together. Properties vary depending upon the atom.
• Allotropes: Different forms of an element due to different arrangements of atoms (e.g., diamond and graphite). Diamond has only covalent bonds between atoms, graphite has London dispersion forces between sheets of carbon.
• Metallic Crystals: Metal atoms held together by metallic bonds. Metals have access to d orbitals and are more likely to give away electrons.
  • Electron Sea Model: Metal cations are stationary in a "sea" of mobile valence electrons.
    • Properties: Shiny, malleable, ductile, good conductors of electricity and heat, high melting points.

Mass of Atoms

• Atomic Mass Unit (amu): Unit for atomic mass. 1 amu = 1.66 x 10^-24 grams = 1/12 mass of Carbon-12.
• The number below the element symbol on the periodic table is the average mass of one atom in amu. This number in grams is the mass of 6.022 x 10^23 atoms.
• Use factor label method for conversions between amu and number of atoms.

Significant Figures

• Captive zeroes (between two non-zeroes) are significant. Leading zeros are never significant. Trailing zeroes are significant only with a decimal point.
• For numbers in scientific notation, the digits in N are significant.
• Exact numbers have infinite significant figures.
• Adding/Subtracting: Answer should have as many decimal places as the measurement with the least amount of decimal places.
• Multiplying/Dividing: Answer should have as many significant figures as the least precise measurement.

The Mole

• Mole: 6.022 x 10^23 of anything (atoms, compounds, molecules, ions)
• Avogadro’s number: 6.022 x 10^23
• Molar Mass: Mass of 1 mole of a substance in grams.
• Molar mass of ions: The same as the mass of a normal atom (mass of electron is virtually insignificant).
• Conversion: Atoms <-> Moles <-> Grams

Molecular Mass

• Molecular Mass: Sum of atomic masses in one molecule.
  • Multiply the atomic mass of each atom by the number of those atoms in the molecule, then add all values together.

Percent Composition

• Percent Composition: Percentage by mass of each element in a compound.
  • Formula: (mass of element in 1 mole of compound / mass of 1 mole of compound) x 100

Chemical Formula

• Chemical Formula: Subscripts indicate the number of each type of atom in one molecule or the number of moles of each type of atom in one mole of molecules.

Empirical Formula

• Empirical Formula: Simplest whole number ratio of atoms in a molecule (cannot have a fraction of an atom).
  • Steps: Convert grams (or percent composition) to moles for each element -- Divide each by the smallest number of moles -- Round to whole numbers -- If necessary, multiply by a common factor to get whole number ratios.

Molecular Formula

• Molecular Formula: Actual number of atoms in the molecule.

Hydrates

• Hydrates: Ionic compounds attached to water molecules (e.g., MgSO4·7H2O - magnesium sulfate heptahydrate).
• Anhydrous Salt: Ionic compound without water in its structure.
• Determining the formula of a hydrate: Weigh hydrate sample -- Heat to boil off water -- Weigh anhydrous compound -- Calculate mass of water lost -- Convert masses of anhydrous salt and water to moles -- Calculate mole ratio.