INTRAMOLECULAR VS INTERMOLECULAR FORCES
So far in Chemistry 11, we have encountered covalent and ionic bonds. They are
both the results of strong intramolecular (intra means within) forces that hold two
atoms tightly together. There are also other weaker forms of bonding interactions
that are intermolecular (inter means between). We will discuss intramolecular
forces in more detail as well as examples of intermolecular forces and how their
bond strengths compare.
Intramolecular Forces
Ionic Bonding
Ionic bonds are formed by the attraction of cations and anions. The cation
donates its electrons to the anion, resulting in full octets for both ions.
The strength of ionic bonds can be estimated by measuring the melting
temperatures of ionic compounds: the higher the melting temperature,
the stronger the ionic bond. Crystals of ions form a specific arrangement
called an ionic lattice such as the structure NaCl.
Covalent Bonding
The bonds of covalent compounds are extremely strong but they do not form
crystal lattices like ionic compounds. So although it would be extremely difficult to
break a covalent bond, it is not particularly difficult to break up a chunk of a
covalent compound (brittle solids remember?). Double and triple bonds are
incrementally stronger than single covalent bonds.
Polarity in Covalent Compounds
Electrons in covalent compounds are not necessarily shared equally. We talked
about electronegativity or electron affinity as the attraction an atom has for
electrons so atoms with higher electronegativity tends to have more electrons
around them. An unequal sharing of electrons between two atoms creates a dipole
moment which is represented in a Lewis structure by drawing an arrow across a
bond pointed towards the atom with the higher electronegativity.
A covalent compound in which its atoms share electrons unevenly is called a polar
covalent compound whereas a compound with equal sharing is nonpolar.
Some molecules share electrons evenly such as the homonuclear diatomic
molecules H2 and N2. Their electron densities will look like the diagram below on
the left.
To determine the polarity of a polyatomic compound, we must find all the dipole
moments and add them together using vector addition. If all the dipole moments
are equal in magnitude and have opposite directions (they cancel out), then we
say that there is no net dipole moment and the compound is nonpolar, otherwise
it is polar.
CO2 is nonpolar while H2O is polar. This is because the dipoles of CO2 are equal and
in opposite directions so they cancel and result in no net dipole. The dipoles of H2O
are not in opposite directions so they don’t cancel completely and result in a net
dipole. The shapes of molecules have to do with the nonbonding electrons atoms
have and are beyond the scope of this course.
Practice: draw arrows representing the dipole moments of the following molecules.
State the polarity of the molecule.
a) O2
b) CH4
c) CH2FCl
d NH3
Intermolecular Forces
Dipole-Dipole Interactions
Dipoles moments in polar covalent compounds cause partial charges on atoms. For
example, in the molecule HF, F has a higher electronegativity which attracts more
electrons to it so it gains a partial negative charge while H is left with a partial
positive charge. The partially positive H atoms will find another partially negative F
atom to form a dipole-dipole interaction due to the electrostatic attraction of
opposite charges.
Hydrogen Bonds
Hydrogen bonds are a special form of dipole-dipole interaction in water because of
the attraction between the partial positive H atoms and partial negative O atoms
are particularly strong. Recall that in the Lewis structure of H2O, O has 4
nonbonding electrons which contribute to a larger partial negative charge than
that of the H atom in HCl or HF.
London Force
Another type of intermolecular force is the London force. Remember that
electrons move constantly so when two atoms
are next to each other their electron densities will repel each other, causing
separation of charge within each atom called temporary dipoles.
Relative Strengths of Intramolecular and Intermolecular Interactions
Intramolecular forces result in the strongest interactions (covalent and ionic bonds)
between two atoms. The strongest intermolecular force is the hydrogen bond
because in a way it is between a covalent bond and a dipole-dipole interaction.
Following the hydrogen bond are dipole-dipole interactions and the weakest of all
interactions are the London forces.
Unit 3 Study Guide with Hebden Reference Pages
History of the Atom 139-144
- Democritus, Aristotle, Dalton, Thomson Rutherford, Bohr, Schrodinger,
Heisenberg
- Plum-pudding model, nuclear model, planetary model, electron cloud model
Atomic Structure 144-151 #13-25
- Proton, electron, neutron, nucleus, atom, molecule, ion, cation, anion,
charge, isotope
- Bohr diagrams of atoms and ions
- Calculating atomic mass using percent abundance and mass of isotopes
- Calculating percent abundance given atomic mass and mass of isotopes
Electron Configuration 151-158 #26-29, p192 #115, 118, 119, 124, 126
-
Shell, subshell, orbital, principle quantum number, isoelectronic, energy level
Electron configurations of atoms, anions and cations
Core notation
The sequence of electron addition: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p
The sequence of electron removal in the same period: p then s then d
Calculating valence electrons from electron configuration
Periodic Table 158-164, 189-191 #87-105, 107, 111, 123 b, c, f, 125
- Mendeleev
- Groups (families), periods, nuclear notation, atomic mass, atomic number,
mass number
- Homonuclear diatomic molecules
- Key groups on the periodic table and their properties: alkali metals, alkaline
earth metals, halogens, noble gases
- General properties of metals, nonmetals and semiconductors (metalloids)
Lewis Structures 183-188 #86
- Lewis structures of atoms, ions and covalent compounds
- Exceptions to the octet rule
- Resonance structures, formal charges
- Polarity, dipole moments, net dipole moments
Bonding, Atomic Radius, Ionization Energy and Electronegativity 165-178 #40, 41,
45, 47-51, 53, 55, 58-61, 74-77
-
Intramolecular forces: ionic bond, covalent bond
Intermolecular forces: dipole-dipole, hydrogen bond, London force
Description & relative strengths
Key trends: electronegativity, ionization energy, atomic radius and valence
electrons
o How do they affect reactivity?