Topic Report of the EEI

My topic report includes all background research gathered.

1.1  Introduction

Tin cans are Iron or steel (iron and carbon mixture) cylinders of metal that are electroplated in a thin Sn layer and used primarily in the hospitality industry when preserving food and in other cases polyesters or other fabrics. When a tin can is scratched, its Sn layer is damaged leaving the Fe exposed to the air around itself which contains Oxygen. Thus begins the oxidation of the Fe and Sn until the both the Fe and Sn metals have corroded away. However when left on its own without a scratch the tin can will not rust for an extended period of time. The layer of Sn will protect the Fe from corrosion much like the role of paint used on automobiles and various forms of machinery. Several factors are involved in tin can corrosion such as the size of the inflection on the can or even the pH of the organic matter present within the tin can.  It is because of these factors tin cans can provide an excellent basis on different areas of electrochemistry which are; electrolysis, galvanic cells, galvanic when transforming them into a galvanic cell and altering the electrolyte or surface area of electrode. This EEI on tin can corrosion will provide an in depth analysis on the factors associated and affecting tin can corrosion.

1.2  Discussion

Research Question

What effect will the pH of the solution within a galvanic cell and surface area of metals have on the rate of corrosion?

The structure of the tin can consists of various layers when containing organic matter and when not. As seen in figure 1 a tin can containing organic matter consists of Fe metal, steel-tin alloy, steel, oxide and oil; refer to Figure 1.

Figure 1

A tin can must consist of the metals Sn and Fe (as itself or in Steel) or else it is not considered a tin can. These two metals are both on the activity series of metals to the left of hydrogen. The activity series of metals is a listing of metals in order from left to right of decreasing reactivity (see below). This means that metals to the left are more like to oxidise or decrease and give up electrons to the metals on the right. Metals to the left of hydrogen prefer to oxidise and create metal ions where metals to the right of hydrogen prefer to convert from metal ions to metal.

K    Na    Li    Ba    Ca    Mg    Al    Zn    Fe    Sn    Pb    (H)    Cu    Ag    Pt   Au

On the reactivity series of metals Fe is to the left of Sn. Therefore according to the law regarding reactivity Fe metal should corrode in preference to Sn. However this is not the case within tin cans partly due to the Fe not being exposed to O, a key component in oxidation an ultimately rusting. In this situation the tin is acting as the anode producing the half reaction of Sn → Sn²⁺ + 2e^- (oxidation of Sn metal to Sn²⁺ metal ions). The Fe protected by the anode acts as the cathode presenting the following half reaction 2e^- + 2H⁺ → H₂ when in an acid solution(reduction of the H⁺ ions to H₂ gas).  The full redox reaction occurring here is Sn +2H⁺ → Sn²⁺ + H₂. The Fe (steel) does not corrode as it is not exposed to O being shielded by the Sn anode; however there are still tiny holes within the Sn layer allowing for reaction at the cathode. The reason there is no exponential change is due to this situation being classified as cathodic control where reaction is occurring, but at a slow pace. (Seniorchem.com, 2000)

When Fe is exposed to the natural gases in the atmosphere or saline properties of water will eventually corrode. Corrosion can be defined as the degradation of metal, the reaction of losing electrons until the metal solid has rusted. The most common form of rust encountered is Fe rust. This is the result of Fe metal oxidising as seen in the following reaction for both Iron 1 and Iron 2.

Fe_(s) → Fe²⁺ _(aq) + 2e-

Fe²⁺→ Fe³⁺ + e-

Figure 2

The second half reaction containing Iron 2 is because the Fe³⁺ ion is the ion in Fe₂O₃ and Fe₂O₃ is the hydrated iron oxide with H₂O.

When dealing with surface area as a variable in galvanic corrosion the area ratio of the anode: cathode is an important variable that affects the corrosion rate pertaining to the anode. This area ratio is also important when considering the amount of current available that arises from the cathodic reaction.

 Figure 2 suggests that the lower the ratio of anode: cathode the higher the corrosion rate (degree of reaction) at the anode is likely to occur. When the same situations ratio is reversed the higher the ratio of anode: cathode results in little influence to help corrode.  If the cathode is larger than the anode, this allows for more oxygen reduction and other cathodic reaction which as whole results in a greater galvanic current.

When putting this information into the context of tin can corrosion the Sn is considered the anode as it is the metal oxidising. Therefore if the ratio of Fe is higher than Sn then the corrosion rate should increase.

Figure 3

Normal galvanic corrosion occurs more rapidly in acidic solutions than in alkaline ones (Chemistry in Use Book 2 p375).

In galvanic corrosion the half reaction of O_{2 (g)} + 2H₂O _(l) + 4e^-­ → 4OH^-_(aq) has a standard electrode potential of 0.40V (refer to Figure 3) where the value is [OH^-] = 1.00mol/L. If this half reaction occurs in a neutral solution of pH of 7 ([OH^-] = 1.0 x 10^-7 mol/L) the non-standard electrode potential changes to 0.81V and at pH = 4.00 a further 0.99V. This information suggests that electrode potential increases as acidity increases. If this is the situation then this half reaction has a greater tendency to occur if the pH decreases (becomes more acidic).

If we look at the following reduction half reaction for corrosion of Fe in respect to pH they change as pH lowers.

At standard pH [OH^-] = 1.0 x 10^-7 mol/L the reduction half reaction is written as:

O₂ + 2H₂O + 4e^- → 4OH^-

For writing the half reaction for acidic solutions H⁺ is the subject and the reaction must be corrected by adding 4H⁺ to each side of the equation, therefore cancelling out H₂O on both sides.

At pH [OH^-] = 1.0 x 10^-10 to 1.0 x 10^-8 mol/L the reduction half reaction is written as:

O₂ + 4H⁺ + 4e^- → 2H­₂O

At pH [OH^-] =1.0 x 10­^-11 or below the oxidation is now a result of hydrogen ions and not oxygen gas therefore the reduction half reaction is written as:

2H⁺ + 2e^- → H₂

The corrosion on Fe under all of these conditions should follow as:

2Fe_(s) + O_{2 (g)} + 2H₂O_(l) → 2Fe(OH)_2(s)

2Fe_(s) + O_{2 (g)} + 4H⁺_(aq) → 2Fe²⁺_(aq) + 2H₂O_(l)

Fe_(s) + 2H⁺_(aq) → Fe²⁺_(aq)^­ + H­­_2(g)

This information demonstrates that as pH decrease (becomes more acidic) the half reaction for the corrosion on Fe changed from one that is initially driven by O ions to one that is driven by H ions. This is all within context as within tin can corrosion the Fe is under cathodic control undergoing the half reaction above while the Sn metal converts to Sn ions.

This information provides strong background theory into the corrosion of Sn within a tin can; however it all derives the following results and chemical half reactions as pH decreases as most organic matter in tin cans at a point is subject to anaerobic respiration where it forms an electrolyte of a low pH. If the pH was to alter and the reaction above take place within a basic electrolyte (pH <7) it would this affect the rate of corrosion and possibly the chemical half reactions taking place. At the same time it is not uncommon for electrode surface area (contrasting ratios on tin to iron) to vary within tin can corrosion and this poses yet another question that is viable for testing under these conditions.

1.3  Conclusion

In tin can corrosion Sn acts as an anode and Fe acts as a cathode. The Sn layer protects the layer of Fe from oxygen and prevents corrosion; however there are still small pockets in the layer of tin exposing the layer Fe to oxygen. This allows for a reaction to take place, although it is not great and is considered to be under cathodic control. When altering the surface area of the anode and cathode high corrosion is achieved when the anode surface area is less than the surface area of the cathode. When taking pH of electrolyte of a tin can galvanic cell into account  the lower the pH is (more acidic) the greater the tendency for a reaction to occur resulting in a higher corrosion rate. The pH present also has effects on what half reaction will take place as Fe when under corrosion in solution of a pH < 3 half reaction alters to the production of H₂ and not O₂.