Intermolecular Forces

Intermolecular Forces

The solid and liquid phases of matter for a given compound or atom are the direction consequence of attractive forces between the molecules or atoms. If no attractive forces existed, then a collection of molecules or atoms would remain in the gas phase regardless of temperature or pressure.

Intermolecular forces are generally much weaker than covalent bonds. Covalent bonds involve the sharing of electrons, intermolecular forces are electrostatic in origin (opposite charges attract), but do not involve the sharing of electrons.

• Only 16 kJ/mol of energy is required to overcome the intermolecular attraction between HCl molecules in the liquid state (i.e. the energy required to vaporize the sample)
• However, 431 kJ/mol of energy is required to break the covalent bond between the H and Cl atoms in the HCl molecule

The temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive intermolecular forces (likewise, the temperature at which a solid melts). Thus, the temperature required to undergo a phase transition is a consequence of the strength of the intermolecular (noncovalent) attractive force.

Neutral molecules (unlike ions) have no net charge, however, attractive forces exist between neutral molecules. The different types of electrostatic forces include:

• Dipole-dipole forces
• London dispersion forces
• Hydrogen bonding forces

Typically, dipole-dipole and dispersion forces are grouped together and termed van der Waals forces (sometimes the hydrogen bonding forces are also included with this group). The term van der Waals forces refer to the fact that these forces produce non-ideal behavior in gases (the effects of which were formalized in the van der Waals equation of non-ideal gases (click here for more information).

Ions have permanent electrostatic charges, and a common electrostatic interaction that has been previously discussed is the attraction between oppositely charged ions, known as ionic bonding (click here for more information).

Q₁ = charge on first particle

Q₂ = charge on second particle

d = distance between centers of particles

k = 8.99 x 10⁹ J m/C²

Thus, the interaction increases:

• As the charges increase
• As the two charges are brought closer together

The minimum distance between oppositely charged ions is the sum of the atomic (ionic) radii. Although atomic radii do vary, it is not over a considerable range, thus, the attraction between two ions is determined primarily by the charge of the ions.

Attractive electrostatic forces can also exist between neutral and charged (ionic) molecules, and these are termed ion-dipole forces

• Involves an interaction between a charged ion and a polar molecule (i.e. a molecule with a dipole)
• Cations are attracted to the negative end of a dipole
• Anions are attracted to the positive end of a dipole
• The magnitude of the interaction energy depends upon the charge of the ion (Q), the dipole moment of the molecule (m) and the distance (d) from the center of the ion to the midpoint of the dipole

The degree of polarity of a molecule is described by its dipole moment, m = Q * r

where

• Q equals the charge on either end of the dipole
• r is the distance between the charges

Dipole moments are generally reported in Debye units

• Ion-dipole forces are important in solutions of ionic substances in polar solvents (e.g. a salt in aqueous solvent)

A dipole-dipole force exists between neutral polar molecules

• Polar molecules attract one another when the partial positive charge on one molecule is near the partial negative charge on the other molecule
• The polar molecules must be in close proximity for the dipole-dipole forces to be significant
• Dipole-dipole forces are characteristically weaker than ion-dipole forces
• Dipole-dipole forces increase with an increase in the polarity of the molecule

Boiling points increase for polar molecules of similar mass, but increasing dipole:

Nonpolar molecules would not seem to have any basis for attractive interactions.

• However, gases of nonpolar molecules, and even the noble gases, can be liquefied indicating that if the kinetic energy is reduced, some type of attractive force can predominate.
• Fritz London (1930) suggested that the motion of electrons within an atom or non-polar molecule can result in a transient dipole moment

Helium atoms (2 electrons)

• Consider the particle nature of electrons
• The average distribution of electrons around each nucleus is spherically symmetrical
• The atoms are non-polar and posses no dipole moment
• The distribution of electrons around an individual atom, at a given instant in time, may not be perfectly symmetrical
• Both electrons may be on one side of the nucleus
• The atom would have an apparent dipole moment at that instant in time (i.e. a transient dipole)
• A close neighboring atom would be influenced by this apparent dipole - the electrons of the neighboring atom would move away from the negative region of the dipole

• This will cause the neighboring atoms to be attracted to one another
• This is called the London dispersion force (or just dispersion force)
• It is significant only when the atoms are close together

The ease with which an external electric field can induce a dipole (alter the electron distribution) with a molecule is referred to as the "polarizability" of that molecule

• The greater the polarizability of a molecule the easier it is to induce a momentary dipole and the stronger the dispersion forces
• Larger atoms tend to have greater polarizability
• Their electrons are further away from the nucleus (any asymmetric distribution produces a larger dipole due to larger charge separation)
• The number of electrons is greater (higher probability of asymmetric distribution)
• Large molecules also tend to have greater polarizability, due to the large number of electrons present

• Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e. between all molecules)

Dispersion forces can add up, so the larger the molecule, the more extensive the dispersion forces. However, since dispersion forces are only strong when neighboring atoms are held quite close, there must be structural complementarity between the molecules for extensive dispersion forces to exist. Thus, molecules that pack well together (i.e. are structurally complementary) exhibit stronger overall dispersion forces. Thus, long unbranched aliphatic molecules can pack well together over their entire length, but branched molecules will disrupt this interface and will exhibit weaker forces. Structural complementarity is the basis of molecular recognition in biological systems:

Hydrogen bonds are considered to be a special type of dipole-dipole interaction

• A bond between hydrogen and an electronegative atom such as F, O or N is quite polar:

• The hydrogen atom has no inner core of electrons, so the side of the atom facing away from the bond represents a virtually naked proton
• This positive charge is attracted to the negative charge of an electronegative atom in a nearby molecule
• Because the hydrogen atom in a polar bond is electron-deficient on one side (i.e. the side opposite from the covalent polar bond) this side of the hydrogen atom can get quite close to a neighboring electronegative atom (with a partial negative charge) and interact strongly with it (remember, the closer it can get, the stronger the electrostatic attraction)
• Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol (so they are still weaker than typical covalent bonds.
• But they are stronger than dipole-dipole and or dispersion forces.
• They are very important in the organization of biological molecules, especially in influencing the structure of proteins

Water is unusual in its ability to form an extensive hydrogen bonding network

• As a liquid the kinetic energy of the molecules prevents an extensive ordered network of hydrogen bonds
• When cooled to a solid the water molecules organize into an arrangement which maximizes the attractive interactions of the hydrogen bonds
• Each water molecule can participate in four hydrogen bonds
• One with each non-bonding pair of electrons
• One with each H atom

• This arrangement of molecules has greater volume (is less dense) than liquid water, thus water expands when frozen
• The arrangement has a hexagonal geometry (involving six molecules in a ring structure) which is the structural basis of the six-sidedness seen in snow flakes

1996 Michael Blaber