Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2013

Lecture Notes 30: 17 April

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Energy of Formation for Ionic Compounds

It turns out that the transfer of an electron from a metal to a non-metal will not generally provide enough energy to favor the process. So how is it that these are in fact favorable reactions?

Let's look at the energy of the process by breaking it into steps and looking at the enthapies of formation starting with free atoms (the reality will be somewhat more complex since we would start with solid metal and molecules, each of which must first react to give free atomic state, but the results are similar). Of course we can get away with this because we are looking at state functions, which as we saw before are pathway independent!

Ionization Energy  Na right arrow Na+ + e-  DeltaH = +495 kJ/mol
Electron Affinity Energy  Cl + e- right arrow Cl-  DeltaH = -348 kJ/mol
Total   DeltaH = + 147 kJ/mol
  However, this value is for the free ions. If we allow them to come together by coulombic attraction into a crystal lattice a large additional amount of energy is released:   
Lattice Energy Na+(g) + Cl-(g) right arrow NaCl(s)  DeltaH = - 449 kJ/mol
  Overall   DeltaH = - 302 kJ/mol

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 472 of Zumdahl 8th ed)
    • vapor pressure (equilibrium)
    • boiling (occurs when the v.p. = atmospheric pressure)
    • superheating (occurs because its hard to initiate bubbles)

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© R A Paselk

Last modified 17 April 2013