|Lecture Notes: 3 November||
Liquids & Solids (Chapter 9)
Weak bonds range from about 10% as strong as a covalent or ionic bond to <1% as strong. Note the examples in the table below. The polar bonds (charge-dipole and dipole-dipole) are for your information. We will focus on van der Waals (* in table) and hydrogen bonds, returning to hydrogen bonding later when we look at water in depth.
Interaction Type Example Average Strength, kcal/mol (kJ/mol) Range** Dipole-dipole 1/r3 Dipole-induced dipole* 0.1-0.2 (0.4-4) 1/r6
(induced dipole-induced dipole)
0.1-0.2 (0.4-4) 1/r6 Hydrogen bond 3-8 (12-30)
van der Waals repulsion 1/r12 *van der Waals interactions, **from Zubay Biochemistry 3rd. Table 4.3, pg. 89.
van der Waals bonds are strictly the dipole-induced dipole and dispersion types, but the term is also often used to refer to other weak bonds other than hydrogen bonds. Notice that the bonds are not only very weak (about 0.1 - 0.3% as strong as a covalent bond), they also do not act at a distance. Essentially they are contact bonds - they sort of act like "velcro" or weak tape. The corollary is that they increase in importance with increases in molecular size.
Thus for hydrocarbons, which are essentially completely non-polar, we see a very low boiling point for methane (CH4) of - 161°C and a fairly regular increase in boiling point as carbons are added (ethane, C2H6 - 88°C; butane, C4H10 - 0.5°C; hexane, C6H14 69°C; octane, C8H18 126°C; etc.) until very large molecules such as paraffin (about 100 C's) and polyethylene (>1,000 C's) are essentially non-volatile. Note also though, in these very large molecules the masses are very large making them non-voltile as well as the forces holding the substance together have become significant due to the very large contact areas.
Hydrogen bonds are a special case of weak bonds. Note that they are significantly stronger (>100 fold) than the other weak bonds at about 4-10% as strong as a covalent bond. Hydrogen bonds only occur when a hydrogen bound to a small, very electronegative atom is brought close to another small, very electronegative atom, such oxygen in water. Essentially this means that we only see hydrogen bonds between hydrogens bound to N, O, or F (second Period electronegative elements) and N, O, or F. So we can have O-H O, O-H N, O-H F, N-H O, N-H N etc. hydrogen bonds. This is because hydrogen bonds involve dipole-dipole interactions, but they also have covalent character (about 10% of the sharing we see in true covalent bonds) which requires that the participating atoms be small enough to get close enough to allow such partial sharing. Hydrogen bonding accounts for much of the special properties of water, which we will see below
These melting points are determined by the types of forces involved (van der Waals, ionic, metallic, or covalent), and, to a lesser extent by the sizes of the particles.
The particles of a liquid are in continuous motion, but the distances between collisions are very short compared to those of gases. Thus liquids are largely incompressible - need to increase pressure about a million-fold to halve volume. Diffusion though liquids is much slower than in gases (hours to days vs. seconds to minutes).
So what's going on? For each phase see a steady, linear increase in temperature with added energy (heat). But there are definite breaks where energy is added (lost on cooling) when a phase change takes place. Let's look in a bit more detail.
Water: water is so ubiquitous, and has so many important and even special properties, that we will talk a bit more about it.
Water is a very unusual, even incredible substance whose amazing properties are often unappreciated because of its ubiquity. Water's special properties include:
© R A Paselk
Last modified 4 November 2009