Topic 11: Corrosion

the deterioration of the material resulting from chemical attack by its environment

what is corrosion

• electro → movement of electrons, chemical → changes in chemical composition

• consider sodium chloride added to water:
  → NaCl granules dissolve
  → the sodium and chlorine atoms separate or “dissociate” to form the ion species Na+ and Cl-
  → the positively charged ion (Na+) = cation
  → the negatively charged ion (Cl-) = anion
  → a solution containing dissolved ions is called an electrolyte

describe the electrochenical process of corrosion

• reaction by which metals form cations that go into aqueous solution (water based), i.e. reaction by which metal corrodes

• called the anodic reaction

• the location where this reaction takes place is called the anode

• the electrons generated in this reaction remain in the metal

what is an oxidation reaction

• reaction where a metal or non-metal cation accepts electrons

• called the cathodic reaction

• the location where this reaction takes place is called the cathode

• this reaction consumes electrons

what is a reduction reaction

• at the same time

• at the same rate

what time and rate must oxidation and reduction reaction occurs at

2H+ (aq) + 2e- (aq) → H2(g)

what is the reaction for corrosion occurring in acidic solutions (containing H+ ions) and no metal ions

O2(g) + 4H+(aq) + 4e-(aq) → 2H2O(l)

what is the reaction for corrosion occurring in an oxidising acidic solution with no metal ions

O2(g) + 2H2O(l) + 4e- (aq) → 4OH-(aq)

what is the reaction for corrosion occurring in a neutral or alkaline solution with no metal ions

• if a piece of iron is left outside it will undergo uniform corrosion and form ‘rust’

• in this case, water from rain or condensation, which contains some dissolved oxygen from the air, sits on the surface

• the anodic corrosion reaction that occurs at microscopic local anodes is:
  → 2Fe(s) → 2Fe2+ (aq) + 4e-

• the cathodic reaction that occurs at a microscopic local cathode is:
  → O2(g) + H2O(l) + 4e- (aq) → 4OH- (aq)

• the overall reaction in this case is the sum of these:
  → 2Fe(s) + O2(g) + 2H2O(l) → 2Fe2+(aq) + 4OH-(aq)

• if the surface is allowed to dry, the ferrous hydroxide dehydrates to form ferric oxide, Fe2O3:
  → 2Fe(OH)2 (aq) → Fe2O3 . 3H2O — rust

describe the rusting of iron

• Every metal has a different tendency to corrode in a particular environment.

• For example: Zinc is chemically attacked or corroded by dilute hydrochloric acid, but gold is not.

• One method for comparing the tendency for metals to form ions in aqueous solution is to compare their halfcell oxidation or reduction potentials to that of a known reference. The reference typically used is a standard hydrogen reference electrode.

• Platinum is the standard hydrogen reference electrode

draw the electrodes used in electrolysis

• Ecell = Ereduction - Eoxidation

  • the potential of the hydrogen reference electrode is by convention assigned a zero voltage (0V

• metals which are more reactive than hydrogen are assigned negative potentials

  • if we connect the metal to the hydrogen reference electrode without a voltmeter, these metals are oxidised to ions and hydrogen ions are reduced to form hydrogen gas

• the reaction with the more negative electropotential → will be oxidised

• the reaction with the most positive electrode potential → will be reduced

how is the half-cell potential calculated

• the galvanic series represents the relative reactivities of the listed metals and alloys

  • metals/alloys near the top are cathodic/noble/inert/unreactive

    • i.e. there is a decreasing tendency to corrode in moving up the table

  • metals/alloys near the bottom are increasingly anodic/reactive

    • i.e. there is an increasing tendency to corrode in moving down the table

• no voltages are provided, but the table works in the same way as for the standard electrode potentials — materials lower down the table have the greater tendency to corrode

• this galvanic series differs from the standard electro potential table in that it includes common alloys as well as elemental metals

• some of the alloys are listed twice, in both active and passive conditions

  • active: conditions where alloy cannot form a stable oxide layer

  • passive: conditions where alloy can form a stable oxide layer

• the presence of a stable oxygen containing film (oxide layer) on the surface shifts these alloys towards the cathodic (less reactive) end of the series

galvanic series

• since most metallic corrosion involves electrochemical reactions, it is importat to under the principles of the operation of an electrochemical galvanic cell, this is a useful concept to describe several different corrosion mechanics as well as methods to prevent corrosion

• there are four components of a galvanic cell

  • 1. anode

    • this gives up electrons to the circuit, i.e. where the anodic/oxidation reaction occurs

    • undergoes corrosion

  • 2. cathode

    • receives electrons from the circuit, i.e. where the cathode/reduction reaction occurs

    • ions that combine with the electrons produce a by-product at the cathode

    • the cathodic reaction depends on the environmental conditions as previously described

  • 3. physical electrical connection

    • the anode and cathode must be electrically connected, usually by a physical connection to permit the electrons to flow from the anode to the cathode to continue the reaction

  • 4. a liquid electrolyte

    • this liquid electrolyte must be:

      • in contact with both the anode and cathode

      • electrically conducting in order to complete the circuit (note: typically this criteria is reached with impurities or species dissolved into the solution)

• galvanic corrosion cells may be characterised into two different groups:

  • composition cells

  • concentration cells

galvanic cells

• a composition cell may be established between and two dissimilar metals

• in each case the metal lower on the standard electrode potential (or galvanic series) table acts as the anode and undergoes corrosion

• the further apart the two metals are, the greater the potential/tendency for corrosion to occur

• during corrosion the metal atoms at the anode go into solution and release electrons

• the cathode is protected from corrosion in this process

• typically, if either metal were exposed to the electrolyte by itself, it would corrode, but at a comparatively slow rate — they would undergo general corrosion

  • the danger comes when they are coupled together i a galvanic cell — the corrosion rate of the anode metal in this case is much greater than its rate of general corrosion

  • in addition, the cathode material is protected from undergoing corrosion by this process

what are composition cells

• the other factor that controls galvanic corrosion is the relative size of the anode and cathode

• if the anode is small relative to the cathode:

  • attack on anode will be large

• if the anode is large relative to the cathode:

  • attack on anode will be small

• why does this occur:

  • during galvanic corrosion both the anodic/oxidising and cathodic/reducing reactions must occur at the same time and at the same rate

  • electrons flow from the anode to the cathode

  • for a given amount of current flow in a galvanic cell, the current density is much greater for the smaller electrode than the larger one

  • therefore, the smaller the anode, the faster the rate of corrosion

• current density equation: i = I/A (current/area), measured in A/m²

how does area effect corrosion

• the amount of metal removed from the anode by corrosion can be determined by Faraday’s equation

  • m = (I x t x M)/(n x F)

    • m = mass in grams of metal lost

    • I = corrosion current in amps

    • t = time in seconds

    • M = atomic mass of the metal (g/mol)

    • n = number of electrons exchanged in the anodic reaction

    • F = Faraday's number (96,500 C/mol or amp.sec/mol)

what is the rate of metal loss to corrosion (for uniform corrosion)

• concentration cells develop due to differences in the concentration of the electrolyte, in this situation:

  • the metal in contact with the most concentrated solution will be the cathode (protected)

  • the metal in contact with the dilute solution will be the anode (will corrode)

• one of the most common concentration cells is the oxygen concentration cell, this occurs in situations where some parts of the electrolyte have a very low oxygen concentration (the anode) and other regions have a high oxygen concentration (the cathode) — this can be represented in the follow way using a schematic galvanic cell

concentration cells

• crevice corrosion is an example of a concentration cell 

• occurs in crevices or under deposits of dirt/corrosion products, in these situations the electrolyte is stagnant (not moving) and becomes depleted in dissolved oxygen

what is crevice corrosion

• pitting corrosion is another form of localised corrosion attack where small pits or holes form by the same mechanism as for crevice corrosion

• once formed, it develops vertically downwards

• important in alloys that rely on an oxide layer for protection

• once pits form, the low oxygen concentration at the bottom makes it difficult for an oxygen layer to form, i.e. no protection from corrosion

what is pitting corrosion

• droplet corrosion → leads to pitting corrosion

• water line corrosion

• buried steel structure

• buried steel pipe

what are other examples of concentration cells/corrosion

1. remove the electrolyte

   • design components so water does not sit for extended periods

   • clean surface deposits which may trap moisture near the surface

     • stagnant water forms O2 concentration cells

2. change materials so that the anodic material becomes cathodic

   • use of galvanising and sacrificial anodes

     • e.g. Zn-coating on steel

     • cathodic protection

3. cover the cathode to isolate it from the system

   • apply protective coatings

     • e.g. paint

     • passive layer (e.g. stainless steel)

     • electroplating (e.g. Sn-coated steel)

4. remove the electrical connection between the anode and cathode

   • use insulating materials between dissimilar materials

how do we prevent corrosion

• corrosion of a metal, M, occurs by the generalised reaction

  • M(s) → Mⁿ⁺ (aq) + ne-

• cathodic protection works by continuously supplying electrons to the metal to be protected, forcing it to be the cathode

• there are two methods of supplying electrons to the metal:

  • sacrificial anode: a metal that is more reactive on the galvanic series is attached to the metal to be protected and acts as a sacrificial anode

    • as it corrodes it provides electrons to the metal to be protected

    • it must eventually be replaced

    • typically use magnesium or zinc as sacrificial anodes

    • examples: protection of buried pipes, ships, offshore drilling platform

  • impressed voltage: a power supply/battery is used to supply electrons to the part to be protected

cathodic protection