Volume 1 Issue 9


Michael Van De Mark, Director Missouri S&T Coatings Institute

The Missouri S&T Coatings Institute was represented at the ICE 2004 show by Michael Van De Mark, Thomas Schuman, Suzan Van De Mark and three graduate students. The picture below is of our booth with Nicole Mason, Suzan and Michael Van De Mark and Emerentiana Sianawati . Both Nicole Mason and Siana are graduates of Missouri S&T. We had five drawings for 2004 St. Patrick's Day sweat shirts. The winners this year were:

Dennis Maples of Taber's Products Inc.
Gregory Alexander of Resolution Specialty Materials
Craig Snider of Hillyard Inc.
Pravin Sitaram of Haartz Corporation
Jenny Beckett

Congratulations to each of our winners. It was great seeing all our friends and former students at the show. We intend on going to at least one of the regional shows this coming year. We will let you know where and when.


Technical Insights on Coatings Science

Corrosion 101: The Oxide Layer

By Thomas Schuman

Corrosion is an electrochemical reaction composed of two separate processes.  The anode and cathode are where oxidation and reduction of species occur, respectively.  The dissolution process of a metal occurs at the anode resulting in the metal atoms being converted from unoxidized metal, of zero oxidation state, to a positive oxidation state ion or a metal oxide by removal of electrons.  Each element has a certain number of protons and electrons; the number of protons minus the number of electrons present is called oxidation state.  Oxidation is the increase in oxidation state of an atom caused by a loss of electrons. 

At the cathode is a reaction that consumes the electrons liberated at the anode.  Common species that consume electrons and their products after consuming the electron(s) are: oxygen, which produces oxide anions in the absence of water or, hydroxide in the presence of water; acids, including water, which produce hydrogen gas; or oxidized noble, i.e., less reactive, metals.  These species are chemical reagents, as is the metal surface, having a chemical potential (energy) to react. 

The potential is often referred to as a reduction potential, derived from the energy necessary to reduce the metal from an oxidized, electron poor, state to a less oxidized state.  The free energy ( D G) to reduce the metal to a zero oxidation state, negative free energy being spontaneous, is calculated D G = – nFE , where n is the number (moles) of electrons used to reduce the metal, F is the Faraday constant (96,485 Coulombs per mole), and E is the potential electrical energy measured in volts.  Coulombs are the unit measure of electron charge and coulombs multiplied by the electron potential in volts yields the number of joules (J), a measure of energy.  For instance, an energy of 4.184 joules warms a gram of water by one degree Celsius.  Reduction of iron (Fe) from ferrous ion (Fe2+) to reduced iron (Fe0) requires a potential energy of –0.41 volts.  If you could plug ferrous ions into a light socket (direct current) having a voltage equal to or more negative than –0.41 volts, it could reduce the ferrous ions to a slab of solid, reduced iron.  The reaction equation looks like:

Fe2+ + 2e — > Fe0              E0 = –0.41V

The potential E superscript 0 indicates the potential for standard conditions, at 25°C and at concentration [Fe2+] of 1 mole per liter.  Let's look at the free energy change for the reduction above, which tells whether the reduction is likely to occur.  Like water, reactions like to flow downhill toward more negative potential energies, but for the reaction above:

D G = – nFE = –(2 mole)(96485 C/mole)(–0.41V) = +79,118 J/mole Fe, or ~79kJ/mole

There are two things you should observe: The reaction has a positive free energy and therefore the reduction is not spontaneous, meaning we would have to push the reaction like a wagon up a hill using more than 79kJ of energy to get the reaction (as written) to occur.  Second, the non-spontaneous reduction reaction in this case requires a lot of energy input to make it occur. 

Let's say we reverse the reaction direction, to make reduced iron oxidize to ferrous iron.  Reversing the reaction also changes the sign of the reaction potential to +0.41V, to now make a VERY spontaneous, heat generating reaction.  As an example, if we could oxidize just one gram of iron (about the weight of a paperclip) totally, immediately, to ferrous ion and transfer all reaction energy into heating one gram of water, we could theoretically heat the liquid water from room temperature (20°C) to 358°C!  Of course water boils at only 100°C and there are reaction rates and solubilities to consider, so the comparison is not entirely fair but serves as an example of the energy in just one paperclip weight of iron.  Aluminum is even more reactive than iron, having a reduction potential of –1.66V, or a reduction free energy of 480kJ/mole. 

Since corrosion has two concurrently required reactions, one generating and one consuming electrons, neither occurs by itself and the corrosion reaction requires a balance in the total electrons given and taken.  Therefore, the total potential energy of corrosion, the cell potential, is made up of the sum of the individual driving forces, the anode and the cathode, working together.  Anode and cathode are then each called 'half reactions'.  The cathode reactions likewise have potentials analogous to those discussed for anodes and usually further push the reactions at the anode.  Typical cathode reduction potentials range from 0V (acid proton) to +1.23V (oxygen in acid solution).  Preventing one half reaction in effect prevents or reduces the rate of both.  Anodic inhibition prevents oxidation of metal, cathodic inhibition prevents the reduction half reaction.  Mixed inhibitors, such as chromate, prevent both half reactions. 

So, why don't our aluminum Coke® cans or steel cars quickly dissolve?  Finely polished aluminum looks very shiny like a mirror, as does polished iron (or steel) surfaces.  However, note that a Coke® can appears more grayish and not terribly reflective, which is due to a thin but effective layer of aluminum oxide on its surface.  Iron does much the same except its oxides are red or brown in color while aluminum oxide is translucent white.  Iron oxide also has a flaky, plate-like structure and is less effective at preventing corrosion than the more compact, nodular aluminum oxide.  Metal oxides, as reaction products of metal with oxygen and/or water, are inert and insulating to effects of water or oxygen.  The oxide layer acts as a protective coating for the reactive metal beneath and prevents extensive contact with water or air in the environment.  The oxide layer coatings slow the corrosion reaction rate to a crawl and can prevent further reaction altogether. 

One last thing, and then we can begin to talk about coating systems and how they supplement the protection of metal surfaces.  The oxides are not perfect, just like the organic coatings we buy and/or sell.  For instance, there can be holes in the coating and the coating might be sensitive to acids, bases or salts.  The anions of common salts, such as chloride or sulfate, can react with most metal oxides to form soluble species that can be dissolved away by water to de-protect the surface, allowing fast corrosion reactions.  The action of an acid or base on the oxide can be similar.  Pourbaix diagrams, named after Marcel Pourbaix who compiled the aqueous phase stabilities of the elements in 1966, delineate the stability of a metal and its oxides to acid or base (i.e., pH) under influence of a cell potential ( Ecell = Eanode + Ecathode ).  The term 'corrosive' points out an agent or ion's ability to enable fast corrosion of metal by removing the protective oxide layer to help expose underlying metal to water and air.

Organic coatings interface with the native oxide layer to improve system properties.  Coatings provide yet another layer, an enhanced barrier, to access by water and air to the metal surface.  The ingredients of coatings can work to stabilize the oxide layer by precluding acids, bases, or other corrosive ions from damaging the oxide or, in some cases, heal an oxide layer after damage has occurred.  In our next letter, Corrosion 102, we will address the ways coating designs are used to enhance surface protection. 

Is there a topic you would like discussed? Contact us by e-mail at coatings@mst.edu.


March 14-18,/05 Basic Composition of Coatings This course provides an overview of the components of paint and how they work. Participants are also introduced to methods for testing and manufacture of paint.
May 16-20/05 Introduction to Paint Formulation This course provides techniques used in
formulating paint from raw materials. It involves formulating and making paint in the laboratory, "Hand on!"
Coatings for Engineers available on-line anytime This course is designed to educate engineers in coatings science. Coatings systems will be covered from cleaning and surface prep to pretreatment, priming and topcoats. Specification and testing sections will aid all engineers who are charged with these tasks.



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