Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 107	Fundamentals of Chemistry	Fall 2008
Lecture Notes: 19 September	© R. Paselk 2005
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Periodic Table and Periodicity Review

Read Chapter 7 on the Periodic Properties of Elements

Let's look again at trends for: (use Clickers to discuss vertical, horizontal, and diagonal trends)

• Atomic size = Atomic radii (hi - lo goes from Cs - F).

  • The pattern shown in the plot below:

Can be represented by the trend shown on the Periodic Table:

• First ionization energy (lo - hi goes from Cs - He).

  • The pattern shown in the plot below:

Can be represented by the trend shown on the Periodic Table:

• Electronegativity (Cs, lowest to F, highest).

  • Electronegativity is a measure of how electrons are shared between two associated atoms. The pattern shown in the plot below:

Can be represented by the trend shown on the Periodic Table:

Two additional properties:

• Density (bump in middle of transition elements).

  • The elements with the highest densities under standard laboratory conditions are shown on the Periodic Table below:

This pattern can be explained by a combination of the atomic sizes (shown on the plot below),



the number of nucleons (increases left - right), and the physical states of the elements. Thus densities of elements on the extreme right are low since they are gases, etc.

• Melting points

  • Highest melting points for elements include C and a bump near the middle of the transition elements as shown on the Periodic Table below:

  

Plots of periodic properties ©1994 Hanson, Harper, Paselk, & Russell

Chemical Bonds

Atoms and molecules can be held together by Strong bonds or Weak bonds. We are first going to look at strong bonds.

There are two kinds of strong bonds: ionic bonds and covalent bonds. We will begin our discussion with ionic bonds.

With the representative elements bond formation generally results in the formation of "octets" of electrons in the outermost shell as we saw earlier.

Ionic bonds and Ionic Compounds:

• Formation: 2 Na + Cl₂ NaCl. Can think of as composed of two equations, the oxidation of Na and the reduction of Cl:
  • Na + energy Na⁺ + e^-
  • Cl + e^- Cl^- + energy
• Lewis Structure
• Crystal Structure (model)
• Ionic Compound Lewis Structure Examples:
  • Potassium bromide
  • Aluminum chloride

Covalent Compounds and Covalent Bonds:

• Formation: H₂ + Cl₂ 2 HCl. In this case can consider that we get two equations each involving a homo dissociation to give radicals, that is atoms with unpaired electrons:
  • H₂ 2 H^.
  • Cl₂ 2 Cl^. These radical then combine to form a bond with these two electrons shared between the two atoms.
• Lewis Structure:
• Covalent Compound Lewis Structure Examples:
  • Water
  • Ammonia
  • Ammonium ion
  • Methane
  • Hydrogen sulfide

Electronegativity

So how do we determine whether two atoms will form an ionic or a covalent bond? Use a special property - electronegativity (EN). Electronegativity is a periodic measure of how electrons are shared by atoms with the highest value for F and the lowest for Cs. There are a couple of ways of determining EN's:

• Look up values on table.
  • You should memorize values for Period 2, Li (1.0) to F (4.0) in steps of 0.5 and Hydrogen (2.1)
• Use the high, low, intermediate approximation:
  • all metals are low
  • the most electronegative of the non-metals are high (N, O, F, S, Cl, Br, I)
  • other elements are intermediate

Bond Type

So how do we use this to predict whether a bond is covalent or ionic?

• For numbers, use a difference of 1.7 to distinguish ionic (D EN> 1.7) and covalent (D EN< 1.7).
• Easier to use the hi, lo, intermediate system:
  • If combine hi + lo, then ionic
  • Otherwise, covalent (hi + hi, lo + lo, hi + inter., lo + inter., inter. + inter.)

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© R A Paselk

Last modified 2 October 2008