Polar Covalent Bonds

Resonance Structure

• Electrons will flow from a region of high negative to charge to a region of low negative charge

  • negative to positive

• Start your curved arrows where there is higher electron density

• Uncharged structures contribute more to the real strutcure

• Negative charges are preferred on electron negative atoms

• The most stable resonance contributor maximizes the number of atoms with octets and places any negative charge on a more electronegative atom

• If you start neutral end resonance structure has to have a + and a - charge

  • if you start with - charge, end with a - charge

  • if you start with a + charge, end with a + charge

Rules:

1) lone pair → pi bond

2) pi bond → lone pair

3) pi bond → pi bond

When comparing resonance structures, evaluate them in this order:

1) Octet rule satisfied?

• Structures where every atom has a complete octet are almost always more stable.

• Any structure with an atom lacking an octet (especially C, N, O, F) is a minor contributor.

2) Number of formal charges minimized?

• Fewer charges → greater stability.

• If two structures both satisfy octets, the one with the least separation of charge is favored.

3) Charges placed on appropriate atoms?

• Negative charge prefers O > N > C.

• Positive charge prefers C > N > O.

• A resonance form that violates this is less stable.

4) Charge delocalization present?

• Structures that allow charge delocalization (spread-out charge) are more stable.

• A localized charge is less favorable.

5) No unreasonable charge separation?

• A positive charge on oxygen is usually unstable.

• A negative charge on carbon is less stable than on oxygen or nitrogen.

The structure that best satisfies these criteria is the most stable contributor.

Rules for determining which resonance structure is more important:

1) Minimize charge

2) Full octet preferred

3A) if a negative charge is present, place on atom which stablizes it best (least basic)

3B) if there has to be an atom with a less than full octet (carboncation), place on atom best able to stabilize it

• Avoid unnecessary + or – charges if a structure can be drawn without them.

• If charges must exist, place them where they make the most chemical sense:

  • Negative charges on more electronegative atoms (like O, N).

  • Positive charges on less electronegative atoms (like C).

Polar vs Non-polar bonds

A molecule is nonpolar if its electron distribution is equal, resulting in no significant electrical charge difference across the molecule.

• It is symmetric

• All outer atoms are identical

• No lone pairs on the central atom (in most cases)

• If dipoles cancel

A molecule is polar if the distribution of electron density is uneven across its bonds, resulting in a dipole moment that creates partial positive and negative charges.

• It has lone pairs on the central atom

• Outer atoms are not identical

• The shape is not symmetric

• If dipoles add up to a net dipole

• A polar bond has a dipole arrow pointing toward the more electronegative atom.

• If the molecule is perfectly symmetric, dipoles cancel.

• Symmetry = nonpolar

• Asymmetry = polar

Formal Charge

• Adding a bond → + charge

• Removing a bond → - charge

Intermolecular Forces

Dipole-Dipole → forces occur between polar molecules as a result of electrostatic interactions among dipoles.

• to attract one another when they orient with unlike charges together

• to repel one another when they orient with like charges together.

• occcur between polar molecules.

Dispersion-Forces → occur between all neighboring molecules and arise because the electron distribution within molecules is constantly changing.

• attractive dispersion forces in nonpolar molecules are caused by temporary dipoles

  • present in all molecules, even nonpolar ones.

• temporary molecular dipoles have only a fleeting existence and are constantly changing

  • caused by temporary fluctuations in electron distribution.

• often strong enough to hold molecules close together so that a substance is a liquid or solid rather than a gas.

  • Strength increases with:

    • Larger atoms

    • More electrons

    • Longer, more spread‑out molecules

Hydrogen Bond → an attractive interaction between a hydrogen atom bonded to an electronegative O or N atom and an unshared electron pair on another O or N atom.

• a very strong dipole–dipole interaction involving polarized O–H or N–H bonds.

• Occurs when H is bonded to:

  • N

  • O

  • F

  • and interacts with a lone pair on another N, O, or F.

Ion-Dipole → Occur between an ion and a polar molecule.

• Very strong.