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Ferroxyl indicator turns blue in the presence of iron ions. This shows that rusting has begun, even if there is no reddish brown rust showing on the surface of the iron. A pink colour is also produced by the ferroxyl indicator. This shows that the ions being lost by iron are being gained by the water and oxygen that are also involved in rusting. In this experiment, one iron nail is wrapped in magnesium, another in copper and one left alone. The nail wrapped in magnesium is not corroded.


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There is slight corrosion on the normal nail and massive corrosion on the copper wrapped nail. Connecting a more reactive metal from higher in the series protects from corrosion. A lower less reactive metal accepts electrons from iron and speeds up the rusting process. Since oxygen and water are needed for corrosion, the main theories behind protection are based on the prevention of any of these from contacting iron.

Tertiary Catalogue

An example is shown below. The disadvantage of this method is that it must be constantly renewed eg oiling a bike chain. This process deposits a thin layer of metal on the object being protected. An iron object becomes coated in atoms of a less reactive metal.

Prophylactic Protection

As the new metal is less reactive, it is slower to corrode. This process is common with gold, silver, nickel, copper and tin as they are low down in the reactivity series. For instance, tin cans are actually steel containing iron coated in tin. Bashed or scratched tins rust even quicker as iron is higher than tin. This is why a bashed tin in a supermarket might be sold off at a cheaper than usual price. When iron is coated in zinc, the process is called galvanising.

Chemistry Corrosion - Concept - Chemistry Video by Brightstorm

While zinc is more reactive than iron, it still offers a physical barrier but also provides chemical protection. This video shows what happens during the galvanising process. Preventing corrosion- galvanisation. An electrolyte is a substance that conducts electricity by mechanisms distinctly different from those in metals. Normal metallic conduction is accounted for by the drifting of electrons that are free to roam around away from their parent atoms. Conduction of electricity in this way is a defining feature of metals.

Electrolytes conduct electricity through a drifting of electrically charged atoms rather than electrons. Electrically charged atoms are known as 'ions'.

Corrosion and Surface Chemistry of Metals

Figure 84 contrasts conduction in metals and electrolytes. Many electrolytes are solutions in water or some other solvent. Strong electrolytes contain a high concentration of ions and conduct electricity well. These good electrolytes include strong acids and alkalis, and most 'salts'. A salt is a chemical compound which, when dissolved, liberates positive and negative ions from their ordered positions in the solid. These released ions are then free to carry electric charges between electrodes immersed in the solution.

Corrosion of Metals - The Chemistry Journey - The Fuse School

A battery contains an electrolyte in either a liquid or a paste-like solution. Liquid electrolytes are used in electrolysis, electroplating, and other chemical processes. A few solids, particularly oxides, conduct electricity ionically meaning through the movement of charged atoms at temperatures close to their melting point. A few other substances exhibit ionic conduction when they are in a molten state.

Whenever chemical rearrangements — or 'reactions' — occur, atoms are required to swap or share electrons with each other. Energy is bound up in any arrangement of atoms. Where the result of a rearrangement binds up less chemical energy than before, chemical potential energy is released. In some cases the energy is released as electrical energy, just as a rolling ball can transfer its potential energy into kinetic energy as it rolls down a hill.

The cell within a battery is configured to intercept these transactions so that the electrical energy made available by the chemical reorganisation can be gathered up. Ultimately things are arranged so there is a separating of electrical charge, driven by chemistry. You may already know that we call this subsequent pushing and pulling of charges an electromotive force or emf, which is denoted and measured in volts. In the case of batteries, the chemical rearrangements involve relocating atoms from the electrolyte onto the surface of one of the 'two dissimilar metals' that Volta prescribed.

At the surface of the other metal, atoms pass into the electrolyte Figure The electricity is inseparable from the chemistry here as the atoms leaving the electrolyte are electrically charged positive and more properly called positive ions. These atoms collect negative charges from the metal as they pass from the electrolyte onto the metal surface, so becoming neutral atoms once more.

Similarly atoms of the dissolving electrode leave negative charge behind and enter the electrolyte as positive ions. A battery will be spent when the electrode that is giving up ions to the electrolyte has been completely consumed, or rather 'dissolved'. Metals that are consumed in this way are said to be undergoing corrosion. In a similar way, I am certain that the bottom of my wheelbarrow is spent when rust has eaten a hole through which I can see the ground.

Chemically the processes in a battery and in the Corrosion of steel have a great deal in common: two dissimilar metals and an electrolyte are all in contact. Chemical reactions that occur spontaneously, such as corrosion, are useful for generating energy. Sometimes, by careful design, we can divert the energy into electricity, otherwise it usually ends up as heat, light, or sound. In some cases, we can even use electricity to drive the chemistry backwards. This is the secret of rechargeable batteries, which we will come to in the next section. It's also behind a clever strategy for inhibiting corrosion called Galvanic protection.

Steel is not the only metal that corrodes in the atmosphere, but as a major structural metal it's the one of which we are most aware. The iron in steel was extracted from an ore in which it was in a stable chemical combination with oxygen and to a much lesser extent with other elements. As it corrodes it is simply returning to that stable state as inexorably as a ball rolls downhill.


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Corrosion occurs whenever steel is in contact with air and moisture, which acts as an oxygen-bearing electrolyte. The 'dissimilar metals' requirement we mentioned earlier can be provided initially by microscopic variations in the steel composition, or in the oxygen concentration of the moisture with which it is in contact. In the corrosion process, chemical rearrangements occur because, chemically speaking, a 'more efficient' arrangement can be found that lowers chemical potential energy.

Figure 86 shows the process schematically, with the steel joined to a dissimilar metal to make things easier to follow. Of the two dissimilar metals, one, and always the same one for any given pair, will be the preferred local host for any oxygen and this is the one that corrodes. Atoms of this metal swap their own electrons for other negative charge electrons stuck to atoms of oxygen. This process, the familiar hallmark of corrosion, is marked by the red arrow in Figure The metal atoms leave the metal and enter the electrolyte 'in search of oxygen'. The other metal cooperates in two ways.

First it takes in refugee electrons from the dissolving material; a local electric current is always associated with corrosion.


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  5. Second, it takes part in the hand over of electrons to oxygen atoms, indicated by the mauve arrow. The final products of corrosion build up in the electrolyte. Corrosion of steel is one of the most familiar examples of oxidation, rust being a mixture of iron oxides. Oxygen has quite a reputation as an aggressive collector of two electrons per atom; in fact among chemists it's common practice to refer to 'oxidation' when atoms surrender electrons to other atoms even if oxygen isn't directly involved.

    You are likely to have seen corrosion if you've looked at ships in harbours. Scratches in the paintwork expose the steel. The 'dissimilar metals' in this case extends to the difference between painted and bare steel. Corrosion proceeds under the paint, causing it to blister and peel off, exposing more steel see Figure When dissimilar metals are in contact in the presence of an electrolyte, corrosion may occur. Knowing this, could we use it to our advantage? We would need to know, in the first instance, which of the two metals corrodes.

    The answer is in Volta's ordering of bad-tasting metals that was presented in Section 4.