Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 110	General Chemistry	Fall 2003
Lecture Notes::Lec 16_1 October	© R. Paselk 2003
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Guest lecturer, Dr. Robert Zoellner

Electrochemical Corrosion

Corrosion is often a very complex electrochemical process. Frequently it is not what we might at first expect. We will look at three situations illustrating some of these processes.

Iron corrosion: From the E° values for iron and oxygen, we would expect ready corrosion of iron.

In fact iron is not easily corroded in a dry atmosphere, and it is protected from corrosion by polishing. Under these conditions iron corrosion is kinetically unfavorable so that it doesn't occur at all rapidly.

Iron corrosion normally occurs in the presence of moisture, since as we'll see below water acts catalytically to enhance the corrosion reactions:

The overall balanced equation can be represented as two half cells:

Of course ferrous hydroxide is not stable in an oxidizing atmosphere, so additional redox reactions and hydration occur to give a final product (rust) with the formula Fe₂O₃*nH₂O, the net reaction at the cathodic site is then (overhead, figure in text):

We can prevent iron corrosion by galvanizing (zinc plating) where the zinc coating protects the iron from exposure to oxygen while also oxidizing preferentially when oxygen does get through (sacrificial corrosion).

Iron can also be protected from corrosion by attaching it to an active metal such as magnesium, aluminum or titanium. In this case the magnesium etc. is in the same environment as the iron, attached directly or by a wire. The magnesium then reacts preferentially keeping the iron reduced. When the magnesium is gone, just replace it. Used with underground iron pipes, ships etc.

Bonding Review

Chemical bonds are the strongest forces that exist between atoms. They are the forces that hold atoms together in molecules and atoms or ions together in solids. We will look at other weak bonds and forces later.

The two most important and common strong bond types in chemistry are ionic bonds and covalent bonds, a third bond type, found in metallic solids, will be discussed later.

Ionic Bonds. An ionic bond is the result of the electrostatic force of attraction between ions that carry opposite electrical charges, as described by Coulomb's Law:

where r is the distance between ion centers in nm.

Covalent Bonds. Covalent Bonds occur with the sharing of electrons by two atoms with similar tendencies to gain and loose electrons. Let's look at the formation of HCl as an example of the creation of a covalent bond:

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 Æ 2 H^.
• Cl₂ Æ 2 Cl^. These radical then combine to form a bond with these two electrons shared between the two atoms.
• H^. + Cl^. Æ H:Cl (This is not a proper Lewis structure, I have only shown the bonding pair of electrons.)
• Lewis Structure:
• Covalent Compound Lewis Structure Examples:
  • Ammonium ion
  • Methane
  • Carbon dioxide
  • Carbon monoxide

Electronegativity. Electronegativity is a periodic measure of how electrons are shared by atoms. It enables us to guess the degree of polarity of a bond between two atoms (i.e. how the bonding electrons are shared), from non-polar covalent (equal sharing) to fully ionic bonds (no sharing). Recall that F has the highest electronegativity value for and Cs has the lowest. We have used two common ways of determining EN's:

• Look up values on table. Recall that a difference of 1.7 is a reasonable way to distinguish likely covalent (D EN< 1.7) from ionic (D EN> 1.7) bonds.
  • You should know 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. Recall that a combination of hi and lo gives ionic bonding, whereas other combinations are covalent.
  • all metals are low
  • the most electronegative of the non-metals are high (N, O, F, S, Cl, Br, I)
  • other elements are intermediate.

Lewis structures are of course quite limited - they work well only for the representative elements, and even then we have to stretch the concept to accommodate all covalent structures., Thus to follow the "Octet rule" we invented resonance for molecules which don't have enough electrons to give octets even with multiple bonding. Clark's rules can help you determine octet violations:

Clark's Method (abbreviated) for determining bonding in covalent Lewis Structures: Add up all of the valence electrons in the structure (remember to add one electron for each negative charge, or subtract one for each positive charge) If S e- = 6y + 2 where y = # atoms other than H, then octet rule is followed with single bonds only. If S e- < 6y + 2 then probably have multiple bonding with the number of multiple bonds = D/2 (remember a triple bond is 2 multiple bonds!). However, note the exceptions with small atoms (H, Li, Be, and B). If S e- > 6y + 2 then have an expanded valence shell. Note that if D = 2, then pentavalent (10 electrons in the valence shell) , and if D = 4, then hexavalent (12 electrons in the valence shell). If you can draw more than one structure, then chose the most symmetrical. If two or more structures are equally symmetrical, then you probably have resonance and should show all structures connected by double arrows.

© R A Paselk

Last modified 26 September 2003