Use the Force - UCLA Chemistry and Biochemistry

Use the Force!
Noncovalent Molecular Forces
Not quite the type of Force we’re talking about…
Before we talk about noncovalent molecular forces, let’s talk very briefly about covalent
The Illustrated Glossary of Organic Chemistry defines a covalent bond as a chemical
bond formed by the overlap of atomic orbitals where the electron pair between the
atoms is shared to some extent. A bond is considered to be covalent when the
electronegativity difference between the bonded atoms is 1.7 or less.
EN difference = 0
EN difference = 0.7
EN difference =1.9
Logically, it follows that noncovalent molecular forces are any attractive forces that
are not covalent! This is important because noncovalent forces control how molecules
associate with each other, which causes many different properties, such as boiling point
and solubility.
**It should be noted that there is no such thing as a perfectly covalent or perfectly ionic
bond. Ionic bonds share electron pairs highly unequally, while covalent bonds share
them approximately equally.
0.0 – 0.4
Nonpolar covalent
0.5 – 0.9
Slightly polar covalent
1.0 – 1.3
Moderately polar covalent
1.4 – 1.7
Highly polar covalent bond
1.8 – 2.2
Slightly ionic bond
2.3 – 3.3
Highly ionic bond
Increasing ΔEN causes increasingly ionic (polar) character. Increasing bond length has
the same effect.
Let’s take a look at the seven noncovalent molecular forces in
order of relative importance…
Ionic Bonding / Anion-Cation Interaction
Anion-cation interaction is caused by electrostatic attraction between an anion and
a cation.
 Electrostatic attraction is an attraction between opposite charges.
NaCl is an example of anion-cation interaction. The electronegativity of Cl is 3.0, and
the electronegativity of Na is 0.9, so the difference is 2.1.
The greater the charge difference, the stronger the attraction.
**Anion-cation interactions can be seen as the “gold standard” of noncovalent molecular
forces. All other forces are trying to mimic it and its strength.**
Dipole-Dipole Interaction
Dipole-dipole interaction is caused by the alignment of bond dipoles with opposite
 A bond dipole is the partial charges assigned to bonded atoms due to differences
in electron density caused by electronegativity, inductive effects, and other
factors. It is denoted with δ+ and δ- on their respective atoms.
 Dipole-dipole interactions are also electrostatic.
Br and F are both electronegative, but F is more electronegative than Br! F has an
electronegativity of 4.0, while Br has an electronegativity of 2.8. The electronegativity
difference is 1.2, which if you refer back to the table on the first page, means that BrF is
not ionic, but polar covalent.
Greater differences in electronegativity between the atoms and high polarizability
(ability to distort electron cloud) increase the magnitude of a bond dipole.
Because dipole-dipole interactions involve only partial charges, it is not as strong as
anion-cation interactions that have full charges. BrF’s lower boiling point (22°C)
compared to NaCl’s (1413°C) confirms this.
Hydrogen Bonding
Hydrogen bonding is caused by electrostatic attraction of a hydrogen atom with a
lone pair of another atom.
 The hydrogen bond donor must have a sufficiently large δ+ charge caused by
bonding to a highly electronegative atom, most commonly O or N.
 The hydrogen bond acceptor must have a lone pair and a sufficiently high
electron density. The accepting atom must have a negative formal charge; if it is
neutral, it must be O or N.
Hydrogen bonding is present in water, for example.
**Hydrogen bonding does not always involve dipole-dipole interactions!
Case in point: hydrogen bonding between fluoride ion and ethanol
Hydrogen bonding also has its biological uses.
Case in point: DNA base pairs
**Hydrogen bonds are strongest when linear!**
London Forces
In the case of argon…
 It is not ionic because there is no electronegativity difference nor are there ions.
 It does not have any bond dipoles so there are no dipole-dipole interactions.
 There are no hydrogens, so of course there is no hydrogen bonding.
Does this mean there are no attractive forces at work (and therefore no energy necessary
for vaporization)?
Argon’s boiling point (-186°C) is above absolute zero (-273°C), so there must be some
force at work!
London forces (dispersion forces) are caused by attraction of polarized electron
 The electron cloud polarization is induced, caused when the electron clouds repel
one another, creating adjacent regions of electron deficiency (δ+) and electron
excess (δ-).
 Because the charge is induced and the electrostatic attraction is only momentary,
this force is very weak.
 All molecules have electrons so all molecules are influenced by London forces!
A higher surface area means a strong attraction which leads to stronger London forces.
Like dipole-dipole interactions, a higher polarizability leads to stronger London forces.
**Until recently, the terms van der Waals force and London force were used
interchangeably, but now van der Waals force is a broad term which refers collectively to
all noncovalent forces that operate in a given circumstance.
Ion-Dipole Interaction
Ion-dipole interaction is a noncovalent attraction between one pole of a bond dipole
and an oppositely charged ion.
 The δ+ end of a bond dipole is attracted to a anion, and the δ- end of a bond dipole
is attracted to an cation.
This explains the water solubility of NaCl.
Occurrences of ion-dipole interactions are not that common.
Cation-Pi Interaction
Cation-pi interaction is a noncovalent attractive force between a cation and a pi
electron cloud, usually the pi electron cloud of an aromatic ring.
Attraction between potassium cation and the pi electron cloud of a benzene ring
This is important in some enzyme-substrate binding.
Occurrences of cation-pi interactions are not that common.
Pi Stacking
Pi stacking (aromatic stacking) is a noncovalent attractive force between two aromatic
Pi stacking is important in DNA.
Not this kind of pie stacking!
This kind of pi stacking!
Occurrences of pi stacking are not that common.
Relative Strength of Noncovalent Forces
Cation-anion (ionic bonds)
Covalent bonds
Hydrogen bonding
Pi stacking
**When more than one force operates, the strongest force dominates.
London forces
Noncovalent Forces and Boiling Point
A stronger attraction means…
 more energy required to disrupt the attraction
 more energy needed for evaporation
 a high boiling point
Boiling point is a useful qualitative measure of attractive forces. The higher the boiling
point, the stronger the noncovalent forces.
Noncovalent Forces and Solubility
Two layers
A and B immiscible
A and B dissolve
If a substance dissolves, A interrupts the attractive forces in B.
This means that it is soluble and that A/B attractions are better than A/A and B/B
attractions. All A/B attractions outweigh all A/A and B/B attractions.
If it is insoluble, then A/B attractions are not better than A/A and B/B attractions. All
A/A and B/B attractions outweigh all A/B attractions.
By “better”, we are referring to the quantity and quality of the attractions. When there
are more and/or greater attractions, then something is more soluble.
There is a poor attraction between oil and water and a strong attraction between water
and water. Like dissolves like. Polar molecules dissolve in polar substances, and
nonpolar molecules dissolve in nonpolar substances.
Practice Problems
1. Rank the following compounds in order of increasing boiling point and very
briefly explain your reasoning: LiBr, IBr, and Br2.
(Practice Problem 14)
2. Is squaric acid soluble in water? Explain your reasoning.
(Practice Problem 18)
3. Consider the molecular structures of propionic acid (CH3CH2COOH) and methyl
acetate (CH3COOCH3).
a) Which substance has a high boiling point? Briefly explain.
b) Which substance is more soluble in water? Briefly explain.
(Practice Problem 2)
Practice Problems
1. Br2 < IBr < LiBr
Br2 has the lowest boiling point, and LiBr has the highest boiling point. Stronger
intermolecular attractions require more energy to disrupt them, which means
substances with stronger intermolecular attractions have higher boiling points.
LiBr is an ionic compound, whereas IBr2 are not, so it has the strongest
attractions and highest boiling point. IBr has a permanent bond dipole, while Br2
does not because it is symmetric. Therefore IBr has stronger attractions than Br2
and has a higher boiling point.
LiBr’s boiling point is 1265°C, IBr 116°C, and Br2 59°C.
2. Solubility occurs when the number and/or magnitude of attractive solute-solvent
interactions is high. Squaric acid is both a hydrogen bond donor and acceptor, so
it can form many hydrogen bonds with water. Many dipole-dipole interactions
are possible as well. All parts of squaric acid are attracted to water, so it should be
highly soluble in water.
H-bond acceptors
H-bond donors
3. propionic acid (CH3CH2COOH) vs. methyl acetate (CH3COOCH3)
a) Stronger molecular associations require more energy to break them, which
leads to a higher boiling point. Neither propionic acid nor methyl acetate
is ionic, so that does not need to be considered. Both have polar bonds (CO, O-H), so they both have dipole-dipole interactions. A propionic acid
molecule can act as both a hydrogen bond acceptor and donor, so it is
involved in hydrogen bonding, while methyl acetate is not. Because
propionic acid has more attractive forces at work, it should have a higher
boiling point.
The boiling point of propionic acid is 141°C and the boiling point of methyl
acetate is 58°C.
b) Solubility can be seen as a strong association of the solute (propionic
acid/methyl acetate) with the solvent (water in this case). A propionic acid
molecule can act as both a hydrogen bond acceptor and donor, so it can
both accept and give H-bonds to water, while methyl acetate can only
accept H-bonds. Propionic acid should be more soluble in water.
Works Cited
Definitions and some images are from The Illustrated Glossary of Organic Chemistry.
Some images are also taken from Dr. Hardinger’s lecture PowerPoints.
Much of the information and explanations are from Dr. Hardinger’s Chem 14C Lecture
Supplement, 5th Edition and Dr. Hardinger’s Chem 14C ThinkBook, 9th Edition.
(Luke Skywalker/Darth Vader and pie picture found off of Google Images; picture of
squaric acid taken from Wikipedia)