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 bonds. 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 COVALENT BOND POLAR COVALENT BOND IONIC BOND 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. ΔEN Bonding 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 charges. 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 ` In the case of argon **Hydrogen bonds are strongest when linear!** London Forces 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 clouds. 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 rings. 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 STRONGEST > Dipole-dipole Hydrogen bonding Ion-dipole Cation-pi Pi stacking > London forces WEAKEST **When more than one force operates, the strongest force dominates.
A stronger attraction means Noncovalent Forces and Boiling Point 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 ABABABABABABABA ABABABABABABABA ABABABABABABABA Homogeneous 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 (C- O, 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, 5 th Edition and Dr. Hardinger s Chem 14C ThinkBook, 9 th Edition. (Luke Skywalker/Darth Vader and pie picture found off of Google Images; picture of squaric acid taken from Wikipedia)