Resources | Subject Notes | Chemistry
This section explains the concept of sacrificial protection, a method used to prevent the corrosion of metals. We will explore this concept through the lens of the reactivity series and the principles of electron loss.
The reactivity series is a list of metals arranged in order of their reactivity, from most reactive (most likely to lose electrons) to least reactive (least likely to lose electrons). A higher position in the reactivity series indicates a greater tendency to lose electrons and form positive ions.
A simplified reactivity series is shown below:
| Metal | Reactivity |
|---|---|
| Sodium (Na) | Very High |
| Magnesium (Mg) | High |
| Zinc (Zn) | Medium |
| Iron (Fe) | Low |
| Copper (Cu) | Very Low |
Metals above a given metal in the reactivity series are more reactive and will displace it from its salt solution. This is the basis of sacrificial protection.
Sacrificial protection involves attaching a more reactive metal (the 'sacrificial' metal) to the metal that needs protection (the 'protected' metal). The sacrificial metal corrodes preferentially, protecting the protected metal from corrosion.
Consider iron (Fe) used in pipelines and ships. Iron is prone to rusting (corrosion). To prevent this, a more reactive metal like zinc (Zn) is often attached to the iron. The zinc will corrode instead of the iron.
The effectiveness of sacrificial protection relies directly on the reactivity series. If the sacrificial metal is higher in the reactivity series than the protected metal, it will be more readily oxidized (lose electrons). This means the sacrificial metal will corrode before the protected metal.
For example, if iron is in contact with zinc, zinc is more reactive than iron. Therefore, the zinc will lose electrons more easily and corrode, protecting the iron from rusting. The zinc atoms lose two electrons to become zinc ions ($Zn^{2+}$), releasing energy in the process.
Corrosion is fundamentally an electrochemical process involving the loss of electrons. When a metal corrodes, it loses electrons and becomes positively charged ions. In sacrificial protection, the more reactive metal (the sacrificial metal) has a greater tendency to lose electrons than the protected metal.
When the sacrificial metal is in contact with the protected metal and an electrolyte (e.g., saltwater), a galvanic cell is formed. The sacrificial metal acts as the anode (where oxidation occurs – electron loss), and the protected metal acts as the cathode (where reduction occurs – electron gain). The electrons released by the sacrificial metal flow to the protected metal, preventing the protected metal from losing electrons and corroding.
The reaction at the anode (sacrificial metal) is:
$M \rightarrow M^{n+} + ne^-$
Where M is the sacrificial metal, $n$ is the charge of the metal ion, and $e^-$ represents the electrons released.
The reaction at the cathode (protected metal) is:
$X^n+ + ne^- \rightarrow X$
Where X is the protected metal ion.
This electron flow effectively prevents the protected metal from corroding.