Chapter 9: Covalent Bonding and Molecules

Covalent Molecules

• Covalent bonds

  • Sharing of electrons, two atoms share electrons

  • Occurs between nonmetal and metalloids

• Octet rule - Atoms are stabilized by having 8 electrons in the valence shell

• Lewis structure - Shows the arrangement of covalently bonded atoms.

  • Dashes are used to represent shared electrons.

  • Lone pairs are two dots beside each other.

• In this image:

  • Nonbonding pairs = 6 pairs, 12 total electrons

    • Aka lone pairs

  • Bonding pairs = 1 pair, 2 total electrons

Covalent Double and Triple Bonds

• Double bonds - two dashes represent 4 shared electrons.

• Triple bonds - three dashes represent 6 shared electrons.

Octet Rule Exceptions

• Incomplete octets have incomplete valence levels.

• Expanded octets often contain 5 or 6 valence.

  • Over 8 electrons.

• Complete octets have 8 electrons.

Drawing Lewis Structures

• Step 1: Find the total amount of electrons.

• Step 2: Draw the base.

• Step 3: Fill the octets of the outer atoms.

• Step 4: Fill the octet of the central atom.

Molecules and Charge

• Polyatomic ions - groups of atoms with an overall charge.

• Formal charges - the atom with a specific charge.

  • Formal charge is an estimate of the total electron charge around bonded atoms.

  • Formal charge assumes that two atoms share electrons.

• Calculating formal charges:

  

Drawing Lewis Structures for Polyatomic Ions

• For polyatomic lewis structures, consider the charges when finding the number of valence electrons.

• If an ion has a negative charge it means that there are extra electrons present.

Resonance Structures

• A set of structures that show how electrons are distributed around a molecule or ion.

• Used when a single Lewis structure is insufficient.

Electronic and Molecular Geometry

• Electronic geometry - arrangement of electrons around the central atom.

  • Look at the groups/bonds.

  • Bonds = dashes. Double and triple bonds count as one set.

• Molecular geometry - shape caused by the arrangement of atoms.

  • Considers nonbonding and bonding pairs.

• Three electronic geometry shapes:

  • Linear

  • Trigonal Planar

  • Tetrahedral

• Five molecular geometry shapes:

  • Linear

  • Trigonal Planar

  • Tetrahedral

  • Bent

  • Trigonal Pyramidal

Lineae

Two Electron “Groups”: Linear

• Electronic Geometry = Linear

  • Angle: 180 degrees

• Molecular Geometry = Linear

• 2 electron groups, 2 bonding groups, 0 lone pairs

Trigonal Planar

Three Electron “Groups”: Trigonal Planar

• Electronic geometry = Trigonal Planar

  • Angle: 120 degrees

• Molecular geometry = Trigonal Planar

• 3 electron groups, 3 bonding groups, 0 lone pairs

Three Electron “Groups” and One Lone Pair: Trigonal Planar

• Electronic geometry = Trigonal Planar

  • Angle: 120 degrees

• Molecular geometry = Ben

  • If there are lone pairs, the MG will be bent.

• 3 electron groups, 2 bonding groups, 1 lone pair

Tetrahedral

Four Electron “Groups”: Tetrahedral

• Electronic geometry = Tetrahedral

  • Angle: 109.5 degrees

• Molecular geometry = Tetrahedral

• 4 electron groups, 4 bonding groups, 0 lone pairs

Four Electron “Groups”: Tetrahedral

• Electronic geometry = Tetrahedral

  • Angle: 109.5 degrees

• Molecular geometry = Trigonal Pyramidal

• 4 electron groups, 3 bonding groups, 1 lone pairs

Four Electron “Groups”: Tetrahedral

• Electronic geometry = Tetrahedral

  • Angle: 109.5 degrees

• Molecular geometry = Bent

• 4 electron groups, 2 bonding groups, 2 lone pairs

Polar Bonds and Molecules

• Polar covalent bond - atoms do not share electrons evenly.

Electronegativity

• Electronegativity (EN) - How strong electrons in bonded systems are “pulled”.

  • Electronegativity increases from left to right.

  • Electronegativity decreased from top to bottom.

  • Fluorine has the strongest electronegativity = 4.0

  • Does not include noble gases

Comparing Covalent, Polar Covalent, and Ionic Bonds

• Calculating difference in atom electronegativity:

  • higher value - smaller value

• Spectrum of polarity

  • non-polar =======> polar

• Non-polar Covalent < 0.5

• Polar Covalent 0.5 - 2.0

• Ionic (not polar) > 2.0

Molecules with Dipoles

• Molecular dipole - an overall polarity in a molecule.

  • AKA net dipole

  • The molecule as a whole is polar => it has a net dipole.

  • Must consider the shape.

• A molecule with a dipole has one side with a slight positive charge, while the other side has a slight negative charge.

• Polar bonds

  • If net dipole is present, the molecule is polar.

  • If there is no net dipole, the molecule is non-polar.

  • The dipole in opposite directions will have an equal force.

  • The dipoles cancel out and the molecule is non-polar.