Molecular Structures, Shapes, and Polarity

Objective

  • Predict Lewis structures of molecules and ions and draw their three-dimensional structures on paper.
  • Build molecular model sets based on Lewis structures.
  • Predict bond angles, electron-pair geometry, molecular geometry, and molecular polarity given the formula of a compound or polyatomic ion.

Introduction

  • Valence electrons are the electrons in the outermost energy level (valence shell) of an atom.
  • Valence electrons participate in chemical bonding, unlike core electrons (inner electrons) which are unavailable for bonding.
  • Valence electrons can be transferred between atoms (typically from a metal to a non-metal), leading to the formation of ions (cations and anions).
  • The attraction between oppositely charged ions forms the basis for ionic bonding.
  • Covalent bonding involves the sharing of valence electrons between two or more atoms (typically non-metals), forming molecules and polyatomic ions.
  • Lewis structures are 2-D representations of bonded atoms.
  • Molecules exist in 3-D space, and representing their 3-D shape on paper is important for visualization.

Drawing Lewis Structures

  • Lewis structures are drawn for covalently bonded atoms (molecules) or covalently bonded atoms with a net charge (polyatomic ions).
  • Lewis structures contain two types of electrons:
    • Lone pairs (LP): valence electrons not involved in bonding, represented as dots.
    • Bonding electrons: shared electrons, represented as lines connecting atoms.
  • Up to 3 bonds (single, double, or triple) can be formed between two atoms.

Diatomic Lewis Structure Example

  • Example of a simple "XY" molecule: X - Y:
    • Shared electrons (bonds) count toward completing the valence of both bonded atoms.

Ammonia (NH3) Lewis Structure

  • Example:
    H - N - H \atop | \quad \,\, \, \, H
  • This is a valid 2-D Lewis structure but not the correct 3-D structure.
    • Nitrogen is the central atom (bonded to at least 2 other atoms).
    • Simple molecules have one central atom; complex molecules can have more.
    • Expanded formulas can show connectivity in complex molecules.
    • Atoms bonded to only one other atom are terminal atoms.
    • Hydrogen is always a terminal atom.
    • Nitrogen has a lone pair (LP).

Drawing Lewis Structures: Step-by-Step

  • Lewis structures are only drawn for covalently bonded species (molecules or polyatomic ions).
  • Consider nonmetals for covalent bonding; metals that tend to form covalent bonds are omitted.

Step 1: Count Valence Electrons

  • Count the total number of valence electrons in the molecular formula.
  • For polyatomic ions, add or subtract electrons based on the ion's charge.
  • The total electrons in the final structure should match the total valence electrons.

Step 2: Determine the Skeletal Structure

  • Choose the least electronegative (EN) atom as the central atom.
  • The central atom is often the atom that shares electrons more and has a lower EN value.
  • Electronegativity increases across a period and decreases down a group.
  • Do not connect atoms in a circular way unless necessary.

Step 3: Connect Atoms with Single Bonds

  • Connect all atoms to the central atom using single bonds.

Step 4: Complete Valence Shells

  • Add lone pairs (LP) to each atom to complete their valence shells based on the following completion criteria:
    • H: 2 electrons
    • Be: 4 electrons
    • B: 6 electrons
    • C, N, O, F: 8 electrons (octet)

Step 5: Verify Electron Count

  • Add up all electrons (LP and bonds) in the structure and check against the count from Step 1.
  • Ensure the total electron count equals the total valence electrons.
  • If the total electron count is higher than the count from Step 1, proceed to Step 6.

Step 6: Adjust Structure by Forming Multiple Bonds

  • Pick two neighboring atoms with at least one set of LP.
  • Add an extra bond between the two atoms and remove a set of LP from each bonded atom.
  • This reduces the total valence count by 2 electrons.
  • Repeat this step if necessary to match the total valence count.
  • If multiple structures are possible, choose the one with a more symmetrical bonding configuration.

Carbon Dioxide (CO₂) Example

  • Applying steps 1-6 to CO_2 yields valid Lewis structures.
  • The symmetrical structure is preferred:
    :O = C = O:

Resonance Structures

  • When more than one valid Lewis structure exists, they are called resonance structures.
  • Resonance structures can be equal or unequal.
  • For CO_2, the symmetrical structure is better than the asymmetrical one (:O rianle C≡O:).
  • The more balanced structure is reported as the best Lewis structure.

Carbonate Ion (CO₃²⁻) Example

  • Applying steps 1-6 to the polyatomic ion carbonate (CO_3^{2-}) requires including 2 extra electrons in the total valence count.
  • Equal resonance structures are obtained, indicated by a double-arrow symbol:
    [ :O - C(=O) - O: ]^{2-}
  • The charge of the polyatomic ion is placed outside the bracket.
  • The angular relationship between carbon and oxygen atoms is determined by the electron geometry of the central carbon atom.

Classify Electron Geometry (EG) of a Molecule

  • Covalently bonded species (molecules and polyatomic ions) exist in 3-D space.
  • Shapes of simple molecules (with one central atom) are determined by maximizing the space between electron groups (LPs or bonded groups).
  • Electron groups shape into a state of maximum separation to minimize electron-electron repulsions (VSEPR principle).
  • VSEPR theory can be demonstrated using balloons tied together to represent electron groups.

Electron Geometry Classifications

  • Electron group shapes resemble balloon shapes and are classified based on the number of groups around the central atom.
  • This is called Electron Geometry (EG) classification.
  • Molecules with more than one central atom can have different EG classifications for each central atom.
  • Double and triple bonds are treated as one electron group.

Common Electron Geometries

  • Linear: 2 electron groups, 180° angle.

  • Trigonal Planar: 3 electron groups, 120° angle.

  • Tetrahedral: 4 electron groups, 109.5° angle.

  • Tetrahedral shapes are represented on paper using triangles and dashes:

    • Solid lines: bonds in the plane of the paper.
    • Solid triangle (wedge): bond sticking out of the plane.
    • Hatched triangle (dashes): bond going into the plane.

Classify Molecular Shape (MS) of a Molecule

  • Molecular Shape (MS) or molecular geometry is another classification for simple molecules with one central atom.
  • Molecules with more than one central atom may have different MS for each central atom.
  • To determine MS, first find the correct electron geometry (EG) and then consider the geometric relationship between the bonded atoms.
  • Electron geometry includes all lone pairs, while molecular shape considers the geometrical relationship between bonded atoms.

Molecular Shape Classification

  • Molecular shape is classified according to each electron geometry type.

Examples

  • Linear: CO_2
  • Trigonal Planar: NO_3^−
  • Bent (120°): SO_2
  • Tetrahedral: CCl_4
  • Bent (109.5°): H_2O
  • Trigonal Pyramidal: NH_3

Drawing 3-D Structures

  • 3-D structures are drawn on paper using lines, triangles, and dashes to represent bonds.

  • Lone pairs are shown in bond positions to better illustrate angular relationships.

    • Methane (CH_4):
      • EG: tetrahedral
      • MS: tetrahedral
    • Ammonia (NH_3):
      • EG: tetrahedral
      • MS: trigonal pyramid
    • Water (H_2O):
      • EG: tetrahedral
      • MS: bent (109.5°)

  • When building molecules with model sets, position the molecule to place the maximum number of bonds (or atoms) on a plane, then draw the "in" and "out" electrons (bonds or LP) using wedges and dashes.

Determine Bond and Molecular Polarity of a Molecule

  • Covalent bonds can be polar or nonpolar based on how strongly atoms attract shared electrons.
  • Unequal sharing of electrons results in a polar bond with positive and negative ends or "poles".
  • Polarity is based on Electronegativity (EN) differences of bonded atoms.

Electronegativity

  • If there is a chemical bond between identical atoms (e.g., N_2), the bond is non-polar.
  • If the EN difference is ≤ 0.4, the bond is considered non-polar (e.g., C-H bond).
  • If the EN difference is > 0.4, the bond is polar.
  • Use partial charge symbols δ+ and δ- to denote charges (lower EN gets δ+, higher EN gets δ−).

Molecular Polarity

  • Polarity refers to having 2 opposite magnetic charge poles.
  • A polar molecule contains a net magnetic pole and has special physical properties.
  • Liquid crystal displays use polar molecules that behave differently toward light when exposed to a changing external magnetic field.
  • To find the net dipole for a molecule, consider all polar bonds and determine if bond polarities cancel each other out.
  • Locate the center of negative and positive poles.
  • If the two pole centers are separated, the molecule is polar.
  • If the two centers coincide, the molecule is non-polar, even if its bonds are polar.

Net Polarity Examples

  • To determine polarity, find the one-point pole center given the partial charges.
  • Use a net dipole symbol +--> to show the location of resulting poles (arrow end: negative pole, cross end: positive pole).
  • Electrons are pulled from the positive pole toward the negative pole.
  • If the poles coincide, there is no net dipole symbol (non-polar).

Examples of Net Polarity

  • Structure A: non-polar
  • Structure B: polar
  • Structure C: non-polar
  • Structure D: polar
  • Structure E: non-polar
  • Structure F: polar