Foundational topics: bond polarity, acid-base behaviour of molecules, and hydrogen bonding.
2.1 Polar Covalent Bonds: Electronegativity
Most bonds are neither fully ionic nor fully covalent, but are somewhere between the two extremes. Such bonds
are called polar covalent bonds the bonding electrons are attracted more strongly by one atom than the other
so that the electron distribution between atoms is not symmetrical.
Bond polarity is due to differences in electronegativity (EN), the intrinsic ability of an atom to attract the shared
electrons in a covalent bond. >> Electronegativity generally increases from left to right across the periodic table,
and decreases from top to bottom. >> Carbon has an electronegativity value of 2.5.
Non-Polar Covalent Bonds: ∆EN < 0.5
Polar Covalent Bonds: ∆EN = 0.5 – 2.0
Ionic Bonds: ∆EN > 2.0
A crossed arrow is used to indicate the direction of
bond polarity. By convention, electrons are displaced in the
direction of the arrow. The tail of the arrow is electron-poor (
) and the head of the arrow is electron-rich ( ).
Electrostatic Potential Maps: use colour to indicate
electron-rich (red) and electron-poor (blue) regions.
Inductive effect: an atom’s ability to polarize a bond (it’s simply the shifting of electrons in a σ bond in
response to the electronegativity of nearby atoms).
>> E.g. metals such as lithium and magnesium, inductively donate electrons, whereas reactive non-
metals, such as oxygen and nitrogen, inductively withdraw electrons.
2.2 Polar Covalent Bonds: Dipole Moments
Molar polarity results from the vector summation of all individual bond polarities and lone-pair
contributions in the molecule. Strong polar substances are often soluble in polar solvents like water,
whereas less polar substances are insoluble in water. *LIKE DISSOLVES LIKE*
Net molecular polarity is measured by a quantity called the dipole moment.
Dipole moment, μ: the magnitude of the charge Q at either end of the molecular dipole times the
distance r between the charges, μ = Q X r. Dipole moments are expressed in debyes (D), where
1 D = 3.336 X 10 coulomb meter (C ∙ m) in SI units.
For example, the unit charge on an electron is 1.60 X 10 C. Thus, if one positive charge and one
negative charge are separated by 100 pm (a bit less than the length of a typical covalent bond), the
dipole moment is 1.60 X 10 C · m, or 4.80 D. Molecules have large dipole moments if they are:
Ionic compounds (largest dipole moments)
Contain largely electronegative atoms (e.g. oxygen and nitrogen) have large dipole moments.
Have lone-pair electrons
Symmetrical structures have zero dipole moments because the individual bond polarities and lone-pair contributions exactly
2.3 Formal Charges
Formal charges are a formalism and don’t imply the presence of actual ionic charges in a molecule. Instead, they’re a device
for electron “bookkeeping” and can be thought of in the following way: a typical covalent bond is formed when each atom
donates one electron. Although the bonding electrons are shared by both atoms, each atom can still be considered to “own”
one electron for bookkeeping purposes.
To express the calculations in a general way, the formal charge on an atom is equal to the number of valence electrons in a
neutral, isolated atom minus the number of electrons owned by that bonded atom in a molecule. The number of electrons in
the bonded atom, in turn, is equal to half the number of bonding electrons plus the nonbonding, lone-pair electrons.
Most substances can be represented unambiguously by Kekule line-bond structures. The true structure of a substance is the
intermediate between the two (a resonance hybrid, which has characteristics of both). Neither structure is correct by itself.
The only difference between resonance forms is the placement of the π and nonbonding valence electrons. The atoms
themselves occupy exactly the same place in both resonance forms, the connections between atoms are the same, and the
three-dimensional shapes of the resonance forms are the same.
2.5 Rules for Resonance Forms Rule 1: Individual resonance forms are imaginary, not real. The real structure is a composite, or resonance hybrid of the
different forms >> They have single, unchanging structures, and they do not switch back and forth between resonance forms.
The only difference between these and other substances is in the way they must be represented in drawings on paper.
Rule 2: Resonance forms differ only in the placement of their π or nonbonding electrons. Neither the position nor the
hybridization of any atom changes from one resonance form to another.
>> In the acetate ion, for instance, the carbon atom is sp2-hybridized and the oxygen atoms remain in exactly the same place
in both resonance forms. Only the positions of the π electrons in the C = O bond and the lone-pair electrons on oxygen differ
from one form to another. This movement of electrons from one resonance structure to another can be indicated by using
curved arrows. A curved arrow always indicates the movement of electrons, not the movement of atoms. An arrow shows that
a pair of electrons moves from the atom or bond at the tail of the arrow to the atom or bond at the head of the arrow.
Rule 3: Different resonance forms of a substance don’t have to be equivalent.
When two resonance forms are non-equivalent, the actual structure of the resonance hybrid resembles the more stable form
more than it resembles the less stable form.
Rule 4: Resonance forms obey normal rules of valency. A resonance form is like any other structure: the octet rule still applies
to second-row, main-group atoms. For example, one of the following structures for the acetate ion is not a valid resonance
form because the carbon atom has five bonds and ten valence electrons:
Rule 5: The resonance hybrid is more stable than any individual resonance form. In other words, resonance leads to stability.
Generally speaking, the larger the number of resonance forms, the more stable a substance is because its electrons are spread
out over a larger part of the molecule and closer to more nuclei.
2.6 Drawing Resonance Forms
2.7 Acids and Bases: The Bronsted-Lowry Definition
>The most important of all concepts related to electronegativity and polarity is that of acidity and basicity.
A Brønsted–Lowry acid is a substance that donates a hydrogen ion, H
A Brønsted–Lowry base is a substance that accepts a hydrogen ion, H +
(The name proton is often used as a synonym for H because loss of the valence electron from a neutral hydrogen atom leaves
only the hydrogen nucleus—a proton.)
When gaseous HCl dissolves in water, a polar HCl molecule acts as an acid and donates a proton, while a water molecule acts
as a base and accepts the proton, yielding chloride ion (Cl ) and hydronium ion (H O ). This reaction is reversible. The chloride
3 ion (the product when the acid HCl loses a proton) is called the conjugate base of the acid, and the hydronium ion (the
product when the base H20 gains a proton), is called the conjugate acid of te base.
Recall: water can act either as an acid or as a base, depending on the circumstances.
2.8 Acid and Base Strength
Acids differ in their ability to donate H . Stronger acids (such as HCl) react almost completely with water, whereas weaker
acids (such as acetic acid, CH3CO2H) react only slightly. The exact strength of a given acid HA in water solution is described
using the acidity constant (Ka) for the acid-dissociation equilibrium. *The concentration of solvent is ignored in the equilibrium
Stronger acids have their equilibria toward the right and thus have larger acidity constants, whereas weaker acids have their
equilibria toward the left and have smaller acidity constants.
Acid strengths are normally expressed using pK vaaues rather than K valuas, where the pK is theanegative common logarithm
of the Ka:
>A stronger acid (larger Ka) has a smaller pKa, and a weaker acid (smaller Ka) has a larger pKa.
2.9 Predicting Acid-Base Reactions from pK Values a
Two ways of remembering:
1) An acid will donate a proton to the conjugate base of a weaker acid, and the conjugate base of a weaker acid will
remove the proton from a stronger acid.
2) The product conjugate acid in an acid–base reaction must be weaker and less reactive than the starting acid and the
product conjugate base must be weaker and less reactive than the starting base.
2.10 Organic Acids and Organic Bases
Organic acids are characterized by the presence of a positively polarized hydrogen atom (blue in electrostatic potential
maps), and are of two main kinds: those acids that contain a hydrogen atom bonded to an electronegative oxygen atom (O—
H), and those that contain a hydrogen atom bonded to a carbon atom next to a C = O bond (O=C—C—H).
Organic bases are characterized by the presence of an atom (reddish in electrostatic potential maps) with a lone pair of
electrons that can bond to H . Nitrogen-containing compounds are the most common organic bases and are involved in
almost all metabolic pathways, but oxygen-containing compounds can also act as bases when reacting with a sufficiently
strong acid. Note: some oxygen-containing compounds can act both as acids and as bases, depending on the circumstances.
2.11 Acids and Bases: The Lewis Definition A Lewis acid is a substance that accepts an electron pair
A Lewis base is a substance that donates an electron pair
>The donated electron pair is shared between the acid and the base in a covalent bond.
The fact that a Lewis acid is able to accept an elec+ron pair means that it must have either a vacant, low-energy orbital or a
polar bond to hydrogen so that it can donate H (which has an empty 1s orbital). Thus, the Lewis definition of acidity includes
many species in addition to H . +
Other examples of Lewis acids:
metal cations (e.g. Mg ) because they accept a pair of electrons when they form a bond to a base.
compounds of group 3A elements (e.g. BF and AlC3 ) because3they have unfilled valence orbitals and can accept electron
pairs from Lewis bases
Many transition-metal compounds (e.g. TiCl , FeCl , ZnCl , and SnCl )
4 3 2 4
H O2(with its two pairs of nonbonding electrons on oxygen) acts as a Lewis base by donating en electron pair to an H in
forming the hydronium ion, H O . 3 +
most oxygen- and nitrogen-containing organic compounds can act as Lewis bases because they have lone pairs of electrons.
a divalent oxygen compound has two lone pairs of electrons
a trivalent nitrogen compound has one lone pair
Some compounds can act as both acids and bases, just as water can. E.g., alcohols and carboxylic acids act as acids when they
donate an H but as bases when their oxygen atom accepts an H . +
2.12 Non-covalent Interactions Between Molecules
There are several different types of noncovalent interactions (called intermolecular forces, or van der Waals forces) such as:
dipole-dipole forces, dispersion forces, and hydrogen bonds.
Dipole-dipole forces: occur between polar molecules as a result of electrostatic interactions among dipoles. The forces can be
either attractive or repulsive, depending on the orientation of the molecules –attractive when unlike charges are together and
repulsive when like charges are together. The attractive geometry is lower in energy, and therefore predominates.
Dispersion forces: occur between all neighbouring molecules and arise because the electron distribution within molecules is
constantly changing. Although uniform on a time-averaged basis, the electron distribution even in nonpolar molecules is likely
to be nonuniform at any given instant.
on one side of a molecule may have a slight excess of electons relative to the opposite side, giving it a temporary dipole
causes a nearby molecule to adopt a temporarily opposite dipole, with the result that a tiny attraction is induced between the
temporary molecular dipoles have only a fleeting existence and are constantly changing, but their cumulative effect is often
strong enough to hold molecules close together so that a substance is a liquid or solid rather than a gas. Hydrogen bond: an attractive interaction between a hydrogen bonded to an electronegative O or N atom and an unshared
electron pair on another O or N atom.
-very strong dipole-dipole interaction involving polarized O-H or N-H bonds.
Hydrogen bonding has enormous consequences for living organisms. E.g.:
-causes water to be a liquid rather than a gas at ordinary temperatures
-hold enzymes in the shapes necessary for catalyzing biological reactions
-cause strands of DNA to pair up and coil into the double helix that stores genetic information.
Hydrophilic (“water-loving”): a substance that is not strongly attracted to water. E.g. table sugar have a number of ionic
charges or polar –OH groups in their structure so they can form hydrogen bonds
Hydrophobic substances (e.g. vegetable oil) do not have groups that form hydrogen bonds, so their attraction to water is
limited to weak dispersion forces.
17.2 Properties of Alcohols and Phenols
Alcohols and phenols can be thought of as organic derivatives of water in which one of the water hydrogens is replaced by an
organic group: H—O—H versus R—O—H and Ar—O—H.
20.2 Structure and Properties of Carboxylic Acids
25.3 D, L Sugars
25.5 Cyclic Structures of Monosaccharides: Anomers
Chapter 3: Organic Compounds: Alkanes and Their Stereochemistry
There are more than 50 million known organic compounds. Each of these compounds has its own physical properties, such as
melting point and boiling point, and each has its own chemical reactivity.
Organic compounds can be classified into dozens of families according to their structural features and the members of a given
family often have similar chemical behaviour.
Alkanes are relatively unreactive and not often involved in chemical reactions. Functional groups: a group of atoms within a molecule that has a characteristic chemical behaviour. A given functional group
behaves in nearly the same way in every molecule that it’s part of.
The group typically reacts independent of the rest of the molecule
Functional Groups with Carbon-Carbon Multiple Bonds (Hydrocarbon Functional Groups)
Alkenes, alkynes, and arenes (aromatic compounds) all contain carbon-carbon multiple bonds.
Alkenes have a double bond
Alkynes have a triple bond
Arenes have alternating double and single bonds in a six-membered ring of carbon atoms
Because of their structural similarities, these compounds also have chemical similarities Functional Groups with Carbon Singly Bonded to an Electronegative Atom
Alkyl halides (haloalkanes), alcohols, ethers, alkyl phosphates, amines, thiols, sulfides, and disulfides all have a carbon atom
singly bonded to an electronegative atom – Halogen, oxygen, nitrogen, or sulfur. Alkyl halides have a carbon atom bonded to halogen (—X)
Alcohols have a carbon atom bonded to the oxygen of a hydroxyl group (—OH)
Ethers have two carbon atoms bonded to the same oxygen
Organophosphates have carbon atom bonded to the oxygen of a phosphate group (—OPO ) 32-
Amines have a carbon atom bonded to a nitrogen
Thiols have a carbon atom bonded to the sulfur of an –SH group
Sulfides have two carbon atoms bonded to the same sulfur
Disulfides have carbon atoms bonded to two sulfurs that are joined together
The bonds are polar, with the carbon atom bearing a partial positive charge, and the electronegative atom bearing a partial
Functional Groups with a Carbon-Oxygen Double Bond (Carbonyl Groups)
Carbonyl groups (C=O) are present in a large majority of organic compounds and in practically all biological molecules. These
compounds behave similarly in many aspects, but different repending on the identity of the atoms bonded to the carbonyl-
Aldehydes have at least one hydrogen bonded to the C=O
Ketones have two carbons bonded to the C=O
Carboxylic acids have an –OH group bonded to the C=O
Esters have an ether-like oxygen bonded to the C=O
Thioesters have a sulfide-like sulfur bonded to the C=O
Amides have an amine-like nitrogen bonded to the C=O
Acid chlorides have a chlorine bonded to the C=O
The carbonyl carbon atom bears a partial positive charge, and the oxygen bears a partial negative charge Alkanes