Chapter 2: Polar Covalent Bonds, Acids & Bases

polar covalent bonds

bonding electrons attracted more strongly by one atom than by the other (ionic character); electron distribution between atoms is not symmetrical

electronegativity (EN)

intrinsic ability of an atom to attract the shared electrons in a covalent bond; based on an arbitrary scale

bond polarity

produced from differences in electronegativity

F

the most electronegative atom (4.0)

Cs

the least electronegative atom (0.7)

nonpolar covalent bonds

atoms with similar EN <0.5

ionic bonds

difference in EN >2

inductive effect

shifting of electrons in a bond in response to EN of nearby atoms

polar bond examples

C bonded to N, O, Si, F, Cl, Br, most halogens & metals

nonpolar bond examples

C bonded to H, B, S, P

electrostatic potential maps

show calculated charge distributions; colors indicate electron-rich (red) & electron poor (blue) regions; arrows indicate direction of bond polarity

dipole moment

net molecular polarity, due to difference in summed charges

strongly polar substances

are soluble in polar solvents like water

nonpolar substances

are insoluble in water

molecule is nonpolar

if bond dipoles point in opposite directions (cancels out)

formal charge

the charge on an atom in a molecule or a polyatomic ion

formal charge equation

FC = number of VE - (½ all shared e^- + all unshared e^-)

heteroatoms

atoms other than C

resonance forms

represent the electrons in a molecular structure that are delocalized; multiple structures of molecules that are connected by a double-headed arrow; different forms of a substance do not have to be equivalent

resonance forms differ

only in the placement of their π bonds or nonbonding electrons

resonance forms must

be valid Lewis structures; the octet rule generally applies

when resonance forms are not equivalent

a structure may be more/less “important” (contribute more/less to the hybrid)

curved arrows

symbol used to show the redistribution of valence electrons (VE); indicate the breaking & forming of bonds

more important contributors

are more stable and contribute to the hybrid

resonance contributors

multiple valid Lewis structures drawn for a single molecule or ion that, when combined, describe the molecule's true electronic structure

resonance hybrid

the actual, more stable structure of the molecule that is a weighted average of contributing structures, with individual resonance contributors existing only as theoretical concepts

Brønsted-Lowry theory

defines acids and bases by their role in reactions that transfer protons (H⁺) between donors and acceptors

Brønsted-Lowry acid

proton donor

Brønsted-Lowry base

proton acceptor

amphiprotic

both a Brønsted-Lowry acid and base

stronger acids

have larger K_a values (& smaller pK_a)

weaker acids

have smaller K_a values (& larger pK_a)

equilibrium favors

the side with the weaker acid

pK_a values

are useful for predicting whether a given acid-base reaction will take place

organic acids

characterized by the presence of positively polarized hydrogen atom

the more acidic the compound

the more stable the resultant anion (conjugate base)

organic bases

MUST HAVE atom with lone pair of electrons that can bond to H⁺

Lewis acids

electron pair acceptors; includes metal cations (electron “sinks”)

Lewis bases

electron pair donors; can accept protons as well as Lewis acids

Brønsted-Lowry acids are not Lewis acids

because they cannot accept an electron pair directly (only a proton would be a Lewis acid)

noncovalent interactions

dispersion/london/Van der Waals forces, dipole-dipole forces, & hydrogen bonds

dipole-dipole forces

occur between polar molecules as a result of electrostatic interactions among dipoles; forces can be attractive or repulsive depending on orientation of the molecules

dispersion forces

occur between all neighboring molecules & arise because the electron distribution within molecules that is constantly changing

hydrogen bonding forces

most important noncovalent interaction in biological molecules; forces are result of attractive interaction between a hydrogen bonded to an electronegative atom & an unshared electron pair on another O or N atom