Chapt 2: Polar Covalent Bonds - Acids and Bases

Polar covalent bonds

bonding electrons are attracted more strongly by one atom than by the other

• electron distribution between atoms is not symmetrical

Electronegativity

The ability of an atom to attract the shared electrons in a covalent bond

Bond Polarity

differences in electronegativity

• More than 2 = ionic bonds

• less than 2 = covalent bonds

Electronegativity trends

increases across a period and up a group

• Fluorine is the most EN

Inductive effect

shifting of electrons in a sigma (σ) bond in response to EN of nearby atoms

“Like dissolves like“

Polar substances are soluble in polar solvents, the same goes for nonpolar substances and solvents

Dipole moment (μ)

net molecular polarity, due to difference in summed charges

μ

magnitude of charge Q at end of molecular dipole times distance r between charges, μ=Q * r in debyes (D)

Dipole moment of an average covalent bond

1 D = 3.336 × 10^-30 C*m, or 4.8D

In symmetrical molecules, the dipole moments

cancel each out in the opposite direction

Formal charge

½ (# of valence electrons - # of bonding electrons - # of nonbonding electrons)

Resonance

describes electron delocalization in a molecule, where

Rules for Resonance Forms

1. Individual resonance forms are imaginary, not real

2. Resonance forms differ only in the placement of their pi or nonbonding

3. Different resonance forms of a substance don’t have to be equivalent

4. Resonance forms obey normal rules of valency

5. The resonance hybrid is more stable than any individual resonance

Any three atoms grouping with a p orbital on each atom

has two resonance forms

Bronsted-Lowry theory defines acids and bases by their role

in reactions that transfer protons (H+) between donors and acceptors

Bronsted-Lowry acid

substance that donates a hydrogen ion, H+

Bronsted-Lowry base

substance that accepts a hydrogen ion, H+

Conjugate base (A-)

product that results from deprotonation of a Bronsted-Lowry acid. Also a measure related to the strength of the acid

Conjugate acid

product that results from protonation of a Bronsted-Lowry base

Acidity Constant (K_a)

Measure of acid strength (for the reaction of an acid (HA) with water to form hydronium ion)

K_a=

[H₃O⁺][A^-] / [HA]

pK_a

negative logarithm of the K_a

Stronger acids have a ___ pK_a

smaller

Weaker acids have a ___ pK_a

larger

Ion product of water

K_w= [H₃O⁺][OH^-]

Organic acids

characterized by the presence of positively polarized hydrogen atom

Two main kinds of organic acids

1. Hydrogen atom bonded to a electronegative oxygen atom (O-H)

2. A hydrogen atom bonded to a carbon atom next to a C-O (double) bond

Organic bases

have an atom with a lone pair of electrons that can bond to H⁺

Oxygen-containing compounds can react as

bases with a strong acid or as acids with strong bases

Lewis acid

electron pair acceptors

Lewis bases

electron pair donors

Bronsted acids are not Lewis acids because

they cannot accept an electron pair directly

Examples of Lewis acids

group 3A elements (BF₃) and transition-metal compounds (TiCl₄)

Curved arrow means

that a pair of electrons move from the atom at the tail of the arrow to the atom at the head of the arrow

Oxygen- and nitrogen- containing organic compounds are

Lewis bases; they have lone pairs of electrons

Non-covalent interactions

One of a variety of nonbonding interactions between molecules

Dipole-Dipole forces

Occurs between polar molecules as a result of electrostatic interactions among dipoles

• depending on orientation of the molecules, the forces can either be attractive or repulsive

Dispersion forces

Occur between all neighboring molecules

• arise due to constant change distribution within molecules

Hydrogen Bond forces

the result of attractive interaction between a hydrogen bonded to an electronegative O or N atom and an unshared electron pair on another O or N atom