Metal Reactivity Lab: Redox Reactions and Reactivity Series

Introduction

This laboratory explores the concept of reactivity in metals. A metal's reactivity is a measure of the ease with which it reacts with other substances (Hein et al. 2023). Some metals, such as sodium, are highly reactive, while others, such as platinum and gold, are inert (Hein et al. 2023). Generally, how reactive a metal is depends on the ease with which it loses electrons in its outer shell to form cations (Hein et al. 2023). Metals that lose electrons easily are more reactive than those that do not. The reactivity series is a list that arranges metals by reactivity, with the least reactive at the bottom and the most reactive at the top.

It is important to study metal reactivity because such knowledge helps in predicting how different metals would behave in a chemical reaction, and therefore, what metal to choose for a certain function (Hein et al. 2023). For instance, very reactive metals such as sodium and potassium react vigorously with oxygen and water and could be used in developing explosives. At the same time, inert metals such as gold and silver, which do not easily corrode, are well suited for making coins and jewelry.

A reactive metal has the capacity to displace another that is less reactive from its aqueous salt solution (Hein et al. 2023). Taking the example of zinc immersed in copper(II) sulfate, zinc displaces the copper in the aqueous copper(II) sulfate solution, leading to the formation of zinc sulfate and copper, as shown in Figure 1 (Hein et al. 2023):

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) ……………………………… Figure 1

Zinc, the free metal, is more reactive than copper. It loses electrons and dissolves into the salt solution, while the less reactive copper(II) ions accept the electrons donated by zinc to form free copper metal.

The net ionic reaction between the elements can be represented as:

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) ………………………………… Figure 2

This can be further simplified to a general reaction of the form:

A(s) + B⁺(aq) → A⁺(aq) + B(s) ………………………………………… Figure 3

In Figure 3, the free metal A donates an electron and dissolves into the cation A⁺(aq), which forms the salt solution, while cation B⁺(aq) accepts an electron and converts to the free metal B(s). However, free metal A must be more reactive than ion B⁺ for the reaction to occur.

Relative reactivity is the comparison of different metals by reactivity using a standard medium, such as their behavior when placed in water or acid. The indicators to observe in relative reactivity include:

(i) Change in temperature
(ii) Formation of a solid precipitate
(iii) Evolution of gas, indicated by the presence of bubbles
(iv) Dramatic change in color

This lab involves performing a series of redox reactions to analyze the relative reactivity of four metals: lead, zinc, magnesium, and copper. Its objective is to enhance understanding of the concept of relative reactivity of metals. The general hypothesis is that the alkaline earth metals (magnesium and zinc) are more reactive than the transition metals (lead and copper), because they lose electrons more easily.

Materials and Methods

(i) Click on the link labeled "Virtual Lab" on the course home page to load the lab environment.
(ii) Load the assignment via the "File" option and choose the "Redox" category.
(iii) Select the assignment titled "Redox Reaction Series." This sets the stage for the first experiment.

(iv) In the "Stockroom," select the "Glassware" tab and choose a 5 mL pipette and a 250 mL Erlenmeyer flask.
(v) Choose 0.1 M magnesium nitrate solution from the "Solutions" tab and move the flask containing the solution to the workbench.
(vi) Choose the container with Cu metal from the "Solids" tab and move it to the workbench.
(vii) Using the pipette, transfer 5 mL of 0.1 M magnesium nitrate to the 250 mL Erlenmeyer flask and record the solution's initial temperature.
(viii) Transfer 1 g of Cu metal to the flask containing the magnesium nitrate solution and record the temperature as well as any other changes that occur.
(ix) Clear the workbench and repeat with another solution/metal combination as shown in Table 1.1.

(x) Record the observations from each solution/metal combination in Table 1.1.
(xi) Carefully analyze the data collected in Table 1.1.
(xii) List the four metals in order of relative reactivity, starting with the least reactive to the most reactive.
(xiii) Justify the selection in step (xii).

Data Table

The table below presents the observations recorded for each metal and solution combination tested during the experiment.

Table 1.1 — Experimental Observations

Solution: Mg²⁺ | Temperature: 25.00°C | Metal: Cu | Signs of Reaction: None
Net ionic equation: Cu(s) + Mg²⁺(aq) → Mg²⁺(aq) + Cu(s)

Solution: Zn²⁺ | Temperature: 25.00°C | Metal: Cu | Signs of Reaction: None
Net ionic equation: Cu(s) + Zn²⁺(aq) → Zn²⁺(aq) + Cu(s)

Solution: Pb²⁺ | Temperature: 25.00°C | Metal: Cu | Signs of Reaction: None
Net ionic equation: Cu(s) + Pb²⁺(aq) → Pb²⁺(aq) + Cu(s)

Solution: Cu²⁺ | Temperature: 31.00°C → 26.74°C | Metal: Mg | Signs of Reaction: Temperature change, color change from blue to colorless, formation of brown solid (copper)
Net ionic equation: Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s)

Solution: Zn²⁺ | Temperature: 30.00°C → 25.00°C | Metal: Mg | Signs of Reaction: Temperature change
Net ionic equation: Mg(s) + Zn²⁺(aq) → Mg²⁺(aq) + Zn(s)

Solution: Pb²⁺ | Temperature: 35.00°C → 25.00°C | Metal: Mg | Signs of Reaction: Temperature change
Net ionic equation: Mg(s) + Pb²⁺(aq) → Mg²⁺(aq) + Pb(s)

Solution: Cu²⁺ | Temperature: 27.00°C → 25.30°C | Metal: Zn | Signs of Reaction: Temperature change, color change from blue to colorless, brown solid forms (copper)
Net ionic equation: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Solution: Mg²⁺ | Temperature: 25.00°C | Metal: Zn | Signs of Reaction: None
Net ionic equation: Zn(s) + Mg²⁺(aq) → Mg²⁺(aq) + Zn(s)

Solution: Pb²⁺ | Temperature: 27.00°C → 25.00°C | Metal: Zn | Signs of Reaction: Temperature change
Net ionic equation: Zn(s) + Pb²⁺(aq) → Zn²⁺(aq) + Pb(s)

Solution: Cu²⁺ | Temperature: 26.00°C → 25.00°C | Metal: Pb | Signs of Reaction: Temperature change, color change from blue to colorless, brown solid forms (copper)
Net ionic equation: Pb(s) + Cu²⁺(aq) → Pb²⁺(aq) + Cu(s)

Solution: Mg²⁺ | Temperature: 25.00°C | Metal: Pb | Signs of Reaction: None
Net ionic equation: Pb(s) + Mg²⁺(aq) → Mg²⁺(aq) + Pb(s)

Solution: Zn²⁺ | Temperature: 25.00°C | Metal: Pb | Signs of Reaction: None
Net ionic equation: Pb(s) + Zn²⁺(aq) → Zn²⁺(aq) + Pb(s)

Conclusion

This laboratory sought to enhance understanding of the relative reactivity of metals. A reactive metal has the capacity to displace another that is less reactive from its aqueous salt solution. A metal's reactivity depends on the ease with which it loses electrons in its outer shell to form cations; the higher the ability to lose electrons, the more reactive the metal.

This lab involved conducting a series of redox reactions to analyze the relative reactivity of four metals: lead, zinc, magnesium, and copper. The general hypothesis at the start of the lab was that alkaline earth metals (magnesium and zinc) are more reactive than transition metals (lead and copper). One gram of solid magnesium metal was placed in a flask containing aqueous solutions of copper(II) nitrate, zinc(II) nitrate, and lead(II) nitrate at different times. The temperature and color of each aqueous solution were observed before and after the reaction. The procedure was repeated for the other three metals — zinc, lead, and copper.

The data showed that magnesium metal displaced the other three metals from their aqueous salt solutions, while zinc displaced both lead and copper. Lead displaced copper ions from the aqueous solution of copper(II) nitrate. This confirmed that magnesium was the most reactive of the four metals, followed by zinc, then lead, with copper being the least reactive. The lab results therefore supported the hypothesis that alkaline earth metals are more reactive than transition metals.

References

Hein, M., Arena, S., & Willard, C. (2023). Foundations of college chemistry, international adaptation (16th ed.). John Wiley & Sons.