# Electrochemical Series and Activity Series The **electrochemical series** is an arrangement of elements in order of their standard electrode potential ($E^{\circ}$) with reference to the Standard Hydrogen Electrode (SHE). This arrangement helps predict the direction of redox reactions and the relative reactivity of elements. For related concepts on electrode potential, refer to . ## IUPAC Convention Under the recommendation of the International Union of Pure and Applied Chemistry (IUPAC), half-cell reactions are typically presented as **reduction reactions**. Therefore, the $E^{\circ}$ values are known as *reduction potentials*. This convention provides a standardized way to compare the tendency of species to gain electrons. ## Oxidation Potential The value of the oxidation potential for an electrode can be obtained by reversing the sign of its reduction potential. Mathematically: $$\text{Oxidation Potential} = -\text{Reduction Potential}$$ This relationship allows easy conversion between the two potential values depending on the context of the reaction being analyzed. ## Standard Conditions The given reduction potential values relate to specific standard conditions only. These conditions ensure consistency and allow for meaningful comparisons between different elements and their potentials. - $1\,M$ solution of ions - $25^{\circ}C$ ($298\,K$) - $1\,atm$ pressure **Note:** Changes in these conditions will alter the $E^{\circ}$ values. For calculations under non-standard conditions, the is used. ## Activity Series of Metals ### Displacement Reactions A *displacement* or *replacement reaction* occurs when a more reactive element displaces another less reactive element from a compound. The general equation for such a reaction is: $$A + XY \rightarrow X + AY$$ In this reaction, atom $A$ replaces atom $X$ in the compound $XY$. For oxidation-reduction concepts, see . **Example: Zinc and Copper(II) Sulphate** When zinc metal is placed in a blue solution of copper(II) sulphate, the blue color slowly fades, and red-orange copper metal is deposited. In this process, $\mathrm{Cu}^{2+}$ ions in the solution are reduced to copper metal, and zinc atoms are oxidized to $\mathrm{Zn}^{2+}$ ions. $$\mathrm{Zn}_{(\mathrm{s})}+\mathrm{Cu}_{(\mathrm{aq})}^{2+} \rightarrow \mathrm{Zn}_{(\mathrm{aq})}^{2+}+\mathrm{Cu}_{(\mathrm{s})}$$ Conversely, when copper metal is placed in a zinc sulphate solution, *no replacement reaction occurs*. This is because: - The standard reduction potential of copper ($E^{\circ}_{\mathrm{Cu}^{2+}/\mathrm{Cu}}$) is greater than that of zinc ($E^{\circ}_{\mathrm{Zn}^{2+}/\mathrm{Zn}}$). - This indicates that it is easier to oxidize Zn to its ions ($\mathrm{Zn} \rightarrow \mathrm{Zn}^{2+}$) and reduce $\mathrm{Cu}^{2+}$ ions to its atoms ($\mathrm{Cu}^{2+} \rightarrow \mathrm{Cu}$). - Thus, Zn, being more easily oxidized, can displace $\mathrm{Cu}^{2+}$ ions from its solution. - Similarly, metals like Mg and Al can also displace $\mathrm{Cu}^{2+}$ ions, but Ag cannot, as Ag is less reactive than Cu. - It is also observed that highly reactive metals like Na and K can displace hydrogen ($\mathrm{H}_2$) from water, while less reactive metals like Cu and Ag cannot. ### Definition of Activity Series Based on their ability to displace other metals and hydrogen from their compounds, metals are ranked. This ranking, where metals and hydrogen are arranged in order of *decreasing ease of oxidation* to their respective ions in aqueous solution, is called the **activity series**. The activity series lists metals from most reactive to least reactive. Metals higher on the list can displace metals (and hydrogen) lower on the list from their compounds. Only metals above hydrogen can displace hydrogen from acids, and only metals above zinc in the series can displace zinc from zinc sulphate solution. ### Rules for Displacement Reactions Using Activity Series 1. **Electron Transfer:** Metals higher on the activity series list (more reactive) will transfer electrons to (and thus displace) metal cations lower on the list (less reactive). 2. **Reaction Vigor:** The greater the separation between the species in the activity series, the more vigorous the reaction will be. **Example 1: Barium and Lead(II) Oxide** When powdered barium is heated with lead(II) oxide, a replacement reaction occurs because Ba is more active (higher on the series) than Pb. Barium is oxidized to form barium oxide, and lead is reduced to elemental lead. $$\mathrm{Ba}_{(\mathrm{s})}+\mathrm{PbO}_{(\mathrm{s})} \rightarrow \mathrm{BaO}_{(\mathrm{s})}+\mathrm{Pb}_{(\mathrm{s})}$$ **Example 2: Iron and Magnesium Chloride** When iron pellets are added to a solution of $\mathrm{MgCl}_2$, no reaction will occur. This is because Fe is below Mg in the activity series, meaning Fe is less reactive than Mg and cannot displace $\mathrm{Mg}^{2+}$ ions. ## Concept Assessment Exercise Predict whether a replacement reaction will occur in the following instances. Defend your conclusion. **a) Magnesium ribbon is in a solution of silver nitrate.** A reaction will occur. Magnesium (Mg) is significantly more reactive than silver (Ag) and will displace $\mathrm{Ag}^{+}$ ions from the solution. $$\mathrm{Mg}_{(\mathrm{s})} + 2\mathrm{AgNO}_{3(\mathrm{aq})} \rightarrow \mathrm{Mg}(\mathrm{NO}_{3})_{2(\mathrm{aq})} + 2\mathrm{Ag}_{(\mathrm{s})}$$ **b) A copper plate is dipped in iron(II) sulphate ($\mathrm{FeSO}_{4}$) solution.** No reaction will occur. Copper (Cu) is less reactive than iron (Fe) and therefore cannot displace $\mathrm{Fe}^{2+}$ ions from the solution. **c) Lead(II) oxide is heated with powdered zinc.** A reaction will occur. Zinc (Zn) is more reactive than lead (Pb) and will displace Pb from lead(II) oxide. $$\mathrm{Zn}_{(\mathrm{s})}+\mathrm{PbO}_{(\mathrm{s})} \rightarrow \mathrm{ZnO}_{(\mathrm{s})}+\mathrm{Pb}_{(\mathrm{s})}$$ **d) Nickel wire is placed into a solution of silver nitrate.** A reaction will occur. Nickel (Ni) is more reactive than silver (Ag) and will displace $\mathrm{Ag}^{+}$ ions from the solution. $$\mathrm{Ni}_{(\mathrm{s})}+2\mathrm{AgNO}_{3(\mathrm{aq})} \rightarrow \mathrm{Ni}(\mathrm{NO}_{3})_{2(\mathrm{aq})}+2\mathrm{Ag}_{(\mathrm{s})}$$