Instead, the rust continually flakes off to expose a fresh metal surface vulnerable to reaction with oxygen and water. xH 2O), commonly known as rust, that does not provide a tight protective film ( Figure 19.17 "Rust, the Result of Corrosion of Metallic Iron").In contrast to these metals, when iron corrodes, it forms a red-brown hydrated metal oxide (Fe 2O 3 Stainless steels are remarkably resistant to corrosion because they usually contain a significant proportion of chromium, nickel, or both. Chromium, magnesium, and nickel also form protective oxide films. Aluminum cans also have a thin plastic layer to prevent reaction of the oxide with acid in the soft drink. For example, aluminum in soft-drink cans and airplanes is protected by a thin coating of metal oxide that forms on the surface of the metal and acts as an impenetrable barrier that prevents further destruction. Some metals, however, are resistant to corrosion for kinetic reasons. Hence it is actually somewhat surprising that any metals are useful at all in Earth’s moist, oxygen-rich atmosphere. This is in accordance with the Second Law of Thermodynamics which tells us that on the transformation of energy from one form to another form entropy always increases and free energy always decreases.Under ambient conditions, the oxidation of most metals is thermodynamically spontaneous, with the notable exception of gold and platinum. The total entropy increases by the rusting which favours spontaneity. ∆G = - 2840 kj/mole, negative ∆G makes the reaction spontaneous. ∆G = ∆H – T ∆S= , ∆G = Gibbs free energy Huge gain in entropy by surrounding makes the rusting reaction spontaneous. ∆S total = ∆S system + ∆S surrounding = ) Jmol-1K The air molecules in the surroundings move faster and make more hard hitting collisions to speed up rusting.Įntropy and free energy change in rusting reaction: Rusting reaction really takes off and becomes a self – supporting raction when the surplus energy 1648 kJ mol-1, localized in Fe and O2 becomes dispersed to the surroundings as ‘heat’ and raises the entropy of surrounding. The reason is, even though there is significant affinity of oxygen for iron, O2 at normal room temperature around 298 K, has only few atoms moving exceptionally fast which can hit the Fe just right so that Fe-Fe and an O-O bonds breake and Fe-O bond can form. Why rusting initiation is a slow process? This surplus stored energy in Fe and O2, makes the corrosion reaction exothermic and it has huge implication in making rusting a spontaneous reaction. In other words, the bond energy of Fe-Fe-Fe-Fe atoms + O = O atoms > bond energy of Fe2O3. Iron atoms (as -Fe-Fe-Fe-) plus oxygen molecules of the air (O-O) have more energy localized within their bonds than does the product of their reaction, iron rust (iron oxide).The rust formation reaction is exothermic. The answer is iron plus oxygen to form iron oxide or rust, the reactants, iron and oxygen don't have to be at a high temperature to have energy localized within them. The first question that comes to our mind who provides energy for this chemical reaction? How iron gets activation energy to cross the hill at ambient temperature? Image below, represents the activation energy barrier for rusting. : ΔrH is enthalpy change of reaction and ∆S is entropy change for the reaction. What makes rusting a slow but self-supporting spontaneous chemical reaction?ĤFe (s) + 3O2 (g) ⟶ 2Fe2O3 (s) : ΔrH = -1648 kJ mol-1, ∆S system = -549.4 JK-1mol-1 at 298K. Current flows through the metal from the anode to the cathode.įe++ and OH- ions combine to produce hydroxides which dehydrate to Fe203.nH2O which is rust. Corrosion occurs at the anode, where metal oxidizes and dissolves. Differential O2 concentration between O2 starved point on metal and other normal part of metal sets up electrolytic corrosion cells consisting of anodes/cathodes because of potential difference. The drop on metal becomes barrier for O2 on metal surface. A drop of stationary water on iron surface is sufficient for the initiation process of rusting.
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