Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2011

Lecture Notes 30: 8 April


Bond Energies and Enthalpies of Reaction

Bond energies, as tabulated in Table 8.4 of your text (p 351) can be used much like heats of formation to calculate the heat (energy) involved in a reaction. Note that in the table all of the bond energies are positive values, so we have to think and assign the appropriate sign depending on what's occuring. Thus, it takes energy to break a bond (in a sense a bond is a situation where the energy is lower, or it wouldn't be a bond) - the bond energy is positive, but energy will be released when a bond is made - the bond energy is negative.

Let's try an example: How much energy is released in the complete combustion of methane?

Writing a balanced equation:

CH4 + 2 O2 right arrow CO2 + 2 H2O

From the table the bond energies are:

Combining the bond energies (reactants - products):

4 (413 kJ/mol) + 2 (495 kJ/mol) - 2 (799 kJ/mol) - 4 (467 kJ/mol)

2642 kJ/mol - 3466 kJ/mol = -824 kJ/mol

Liquids & Solids

Gases vs. Liquids vs. Solids:

  • Gases are either monatomic (i.e. inert gases) or covalently bonded molecules.
  • Pure liquids at "comfortable" temperatures (20 - 40 °C) all consist of molecules (uncharged) or are metals (Hg, Fr, Cs, Ga, and Rb).
  • At higher temperatures (around 300 - 1200 °C) ionic compounds become liquids.
  • Melting points for pure metals range up to 3680 K, while the highest melting non-metal is carbon at 3820 K.
    • These melting points are detemined by the types of forces involved (van der Waals, ionic, metallic, or covalent), and, to a lesser extent by the sizes of the particles.

Liquids: 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 under standard conditions).

  • Evaporation
    • cooling by evaporation
      • number of molecules vs. KE diagram (see Fig 10.41, p 461 of Zumdahl 5th ed)
    • vapor pressure (equilibrium)
    • boiling (occurs when the v.p. = atmospheric pressure)
    • superheating (occurs because its hard to initiate bubbles)
  • Freezing
    • freezing/melting point: temperature at which solid and liquid are in equilibrium.
    • heat of fusion/crystallization
    • supercooling (occurs because its hard to "seed" crystals - will see later when solids are discussed)
      • Glasses: supercooled "solid" liquids.
  • Heating/cooling curves, below.
  • Viscosity
    • due to van der Waals forces and long molecules
    • due to strong H-bonding (ethylene glycol, HOCH2CH2OH has a viscosity much like high MW syrup)
  • Surface Tension: due to attraction of molecules of liquid for each other (diagram).
    • meniscus
  • detergents/surfactants and wetting (water striders, ducks etc.)

Weak Bonds

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:

 Interaction Type
Example Average Strength, kcal/mol (kJ/mol) Range**
Charge-dipole -NH3+ Cl-   1/r2
Dipole-dipole ClCH3 ClCH3    1/r3
Dipole-induced dipole* CH4 ClCH3 0.1-0.2 (0.4-4)  1/r6
(induced dipole-induced dipole)

0.1-0.2 (0.4-4)
Hydrogen bond  Hydrogen bonded water dimer 3-8 (12-30)  

 van der Waals repulsion
*van der Waals interactions, **from Zubay Biochemistry 3rd. Table 4.3, pg. 89.
The van der Waals bonds are strictly the dipole-induced dipole and dispersion types, but 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 a weak tape. The corollary is that they increase in importance with increases in molecular size (and thus contact surface).

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 forces holding the substance together have become significant due to the very large contact areas.

Weak Bonds

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. 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. (Grp IVA-VIIA bp plot, text Fig 10.4, p 427:

Plot of covalent hydrides in groups 4A-7A

Hydrogen bonding accounts for much of the special properties of water, such as its very high boiling point (261°C higher than methane with only a 10% increase in MW), high viscosity, high heat capacity etc. which in turn are due to the strong bonds between the individual molecules so they stick together.

Examples of water excluding non-polar substances to force the formation of biomembranes, separate out oils etc.


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© R A Paselk

Last modified 8 April 2011