Experiment: Observing Metal Displacement Reactions - Tenth Grade Science
The chemical reactivity of metallic elements varies across a predictable thermodynamic hierarchy known as the reactivity series. By immersing elemental metal strips into aqueous solutions of other metal salts, students observe single displacement reactions. These observable reactions—marked by color changes, metallic precipitation, and temperature shifts—provide the empirical basis for ranking metals by their relative reducing power.
Experimental Aim and Objectives
The primary objectives of this practical investigation are:
1. To observe the single displacement reactions that occur when strips of clean zinc, iron, copper, and aluminum are placed into aqueous solutions of copper(II) sulfate, iron(II) sulfate, zinc sulfate, and aluminum sulfate.
2. To record macroscopic reaction evidence, including the deposition of precipitated metals, solution color changes, and reaction speeds.
3. To deduce the relative positions of zinc, iron, copper, and aluminum in the electrochemical reactivity series from experimental displacement trends.
Theoretical Background and Chemical Principles
A single displacement reaction occurs when a more chemically reactive metal displaces a less reactive metal cation from its aqueous salt solution.
THE CHEMICAL DISPLACEMENT MECHANISM
More Reactive Metal (A) + Salt of Less Reactive Metal (B)
| |
v v
Readily Loses Electrons Accepts Electrons (Reduced)
(Oxidized to Cation) (Precipitates as Neutral Metal)
\ /
v v
Salt Solution of A + Precipitated Metal (B)
Redox Principles
From an electrochemical perspective, metallic reactivity reflects a metal's tendency to lose valence electrons and oxidize into positive cations:
$$M \longrightarrow M^{n+} + n e^-$$
When an elemental metal with high oxidation potential (such as zinc) is immersed in a solution containing cations of a metal with lower oxidation potential (such as copper), electrons transfer spontaneously from the more electropositive metal to the less electropositive cation:
$$\text{Zn}(s) + \text{Cu}^{2+}(aq) \longrightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$$
In this reaction: Zinc atoms lose two electrons, undergoing oxidation to colorless zinc ions: $\text{Zn} \to \text{Zn}^{2+} + 2e^-$. Copper(II) ions in solution gain two electrons, undergoing reduction to elemental copper: $\text{Cu}^{2+} + 2e^- \to \text{Cu}$. Macroscopically, the blue color of aqueous copper(II) sulfate fades as copper ions are consumed, while a reddish-brown deposit of metallic copper precipitates onto the zinc strip.
Conversely, a less reactive metal cannot transfer electrons to the cation of a more reactive metal. Placing metallic copper into a zinc sulfate solution produces no reaction:
$$\text{Cu}(s) + \text{ZnSO}_4(aq) \longrightarrow \text{No Reaction}$$
THE LABORATORY REACTIVITY SERIES
Most Reactive: Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminum (Al) <-- Tested Metal
Zinc (Zn) <-- Tested Metal
Iron (Fe) <-- Tested Metal
Lead (Pb)
[ Hydrogen (H) ]
Least Reactive: Copper (Cu) <-- Tested Metal
Silver (Ag)
Gold (Au)
Apparatus and Reagents Required
LABORATORY EXPERIMENTAL ARRAY
Test-Tube Rack containing 16 Pyrex Test Tubes (4x4 Grid):
CuSO4 Solution FeSO4 Solution ZnSO4 Solution Al2(SO4)3
+---------------+ +---------------+ +---------------+ +---------------+
Zn | (1) Blue->Clr | | (2) Pale Grn | | (3) No React | | (4) No React |
Strips| Red Cu Deposit| | Black Fe Dep | | | | |
+---------------+ +---------------+ +---------------+ +---------------+
Fe | (5) Blue->Grn | | (6) No React | | (7) No React | | (8) No React |
Nails | Red Cu Deposit| | | | | | |
+---------------+ +---------------+ +---------------+ +---------------+
Cu | (9) No React | | (10)No React | | (11)No React | | (12)No React |
Wire | | | | | | | |
+---------------+ +---------------+ +---------------+ +---------------+
Al | (13)Blue->Clr | | (14)Grn->Clr | | (15)Clr Soln | | (16)No React |
Foil | Vigor Red Cu | | Iron Deposit | | Gray Zn Dep | | |
+---------------+ +---------------+ +---------------+ +---------------+
Chemical Reagents
Aqueous Copper(II) sulfate ($1.0\text{ M } CuSO_4\cdot 5H_2O$ - bright blue). Aqueous Iron(II) sulfate ($1.0\text{ M } FeSO_4\cdot 7H_2O$ - fresh pale green). Aqueous Zinc sulfate ($1.0\text{ M } ZnSO_4\cdot 7H_2O$ - colorless). Aqueous Aluminum sulfate ($1.0\text{ M } Al_2(SO_4)_3\cdot 18H_2O$ - colorless). Strips of zinc metal (granules or polished zinc sheet, $1 \text{ cm} \times 4 \text{ cm}$). Clean iron nails (free of surface rust, $5 \text{ cm}$ length). Clean copper wire coils or polished copper foil strips. Strips of aluminum sheet or thick aluminum foil. Fine-grit emery paper or sandpaper for surface preparation. Dilute hydrochloric acid ($1.0\text{ M } HCl$, used briefly for oxide removal).
Laboratory Apparatus
Sixteen standard borosilicate glass test tubes ($15 \text{ mm} \times 150 \text{ mm}$). Two test-tube racks. Fine tweezers and forceps. Safety goggles, nitrile laboratory gloves, and test-tube brushes.
Step-by-Step Experimental Procedure
- Mechanical Surface Preparation: Using emery paper, rub the surfaces of all iron nails, zinc strips, copper wires, and aluminum strips until their outer oxide coatings are removed and the metals display metallic luster.
- Arrange sixteen test tubes in a test-tube rack in a $4 \times 4$ grid, labeling the rows with the four metals and the columns with the four salt solutions.
- Dispense roughly $10\text{ mL}$ of copper(II) sulfate solution into each of the four test tubes in the first column.
- Dispense roughly $10\text{ mL}$ of iron(II) sulfate solution into each tube in the second column.
- Dispense roughly $10\text{ mL}$ of zinc sulfate solution into each tube in the third column.
- Dispense roughly $10\text{ mL}$ of aluminum sulfate solution into each tube in the fourth column.
- Using tweezers, immerse a polished zinc strip into each tube in the first row ($CuSO_4, FeSO_4, ZnSO_4, Al_2(SO_4)_3$).
- Immerse a clean iron nail into each tube in the second row.
- Immerse a clean copper strip into each tube in the third row.
- Immerse an oxide-free aluminum strip into each tube in the fourth row.
- Observe the tubes over 30 minutes, noting color changes in the solutions, metallic deposits on the metal strips, and gas evolution. Record all observations systematically.
DISPLACEMENT REACTION IN A TEST TUBE
+---+
| |
| |
| |
|===| <--- Pale green FeSO4 solution formed
| . |
| | | <--- Iron nail coated with reddish-brown
|===| metallic copper deposit
+---+
Observations and Experimental Results
OBSERVATION AND DISPLACEMENT MATRIX
Metal Strip CuSO4 Solution FeSO4 Solution ZnSO4 Solution Al2(SO4)3
-----------------------------------------------------------------------------
Zinc (Zn) Displacement Displacement No Reaction No Reaction
Blue fades Green fades No change No change
Red Cu deposit Dark Fe deposit
-----------------------------------------------------------------------------
Iron (Fe) Displacement No Reaction No Reaction No Reaction
Blue turns green No change No change No change
Red Cu deposit
-----------------------------------------------------------------------------
Copper (Cu) No Reaction No Reaction No Reaction No Reaction
No change No change No change No change
-----------------------------------------------------------------------------
Aluminum (Al) Displacement Displacement Displacement No Reaction
Blue fades fast Green fades Colorless No change
Dense Cu deposit Dark Fe deposit Gray Zn deposit
Chemical Equations and Reaction Analysis
- Aluminum with Copper(II) Sulfate: Aluminum displaces copper vigorously once its protective oxide layer is breached:
$$2 \text{Al}(s) + 3 \text{CuSO}_4(aq) \longrightarrow \text{Al}_2(\text{SO}_4)_3(aq) + 3 \text{Cu}(s)$$
The blue solution turns colorless, and metallic copper precipitates rapidly.
2. Zinc with Copper(II) Sulfate: Zinc displaces copper ions:
$$\text{Zn}(s) + \text{CuSO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{Cu}(s)$$
The blue solution fades to colorless, and reddish-brown copper deposits on the zinc.
3. Zinc with Iron(II) Sulfate: Zinc displaces iron ions:
$$\text{Zn}(s) + \text{FeSO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{Fe}(s)$$
The pale green solution fades, and grayish-black iron deposits on the zinc.
4. Iron with Copper(II) Sulfate: Iron displaces copper:
$$\text{Fe}(s) + \text{CuSO}_4(aq) \longrightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$$
The blue solution turns pale green (forming $FeSO_4$), and a reddish-brown layer of copper coats the nail.
5. Copper in Any Solution: Copper fails to displace zinc, iron, or aluminum, confirming it is the least reactive metal tested.
From these experimental results, the metals can be ranked in descending order of chemical reactivity:
$$\text{Aluminum (Al)} > \text{Zinc (Zn)} > \text{Iron (Fe)} > \text{Copper (Cu)}$$
Laboratory Precautions and Sources of Error
- Removal of Oxide Passivation: Aluminum forms an impervious surface oxide film ($Al_2O_3$) upon contact with air. If aluminum strips are not thoroughly abraded with emery paper or dipped briefly in dilute $HCl$ before testing, this oxide layer blocks contact with the solution, delaying or preventing displacement reactions.
- Use Fresh Iron(II) Sulfate: Aqueous iron(II) sulfate oxidizes in the presence of dissolved atmospheric oxygen, converting pale green $Fe^{2+}$ into yellow-brown iron(III) compounds ($Fe_2(SO_4)_3$ and basic ferric sulfates). Always prepare $FeSO_4$ solutions fresh before the laboratory session, and store them with a clean iron nail to maintain the iron in the $+2$ oxidation state.
- Allow Sufficient Reaction Time: Moderate displacement reactions—such as iron in copper sulfate or zinc in iron sulfate—require 15 to 20 minutes to develop visible metallic deposits and noticeable solution color shifts. Avoid rushing experimental conclusions.
Chapter Exercises and Solutions
Exercise 1: An iron nail is suspended in a beaker containing blue copper(II) sulfate solution for 45 minutes. Describe the two macroscopic changes observed in the beaker, write the balanced chemical equation with state symbols, and identify which species acts as the reducing agent.
Solution 1: After 45 minutes, two distinct macroscopic changes are visible: the bright blue color of the copper(II) sulfate solution fades, transitioning into a pale light green color due to the formation of aqueous iron(II) sulfate ($FeSO_4$), and a thick, reddish-brown deposit of metallic copper coats the immersed portion of the iron nail. The balanced chemical equation is $\text{Fe}(s) + \text{CuSO}_4(aq) \to \text{FeSO}_4(aq) + \text{Cu}(s)$. In this single displacement reaction, elemental iron ($\text{Fe}$) donates two valence electrons to reduce copper(II) ions ($\text{Cu}^{2+} + 2e^- \to \text{Cu}$), while itself being oxidized into iron(II) ions ($\text{Fe} \to \text{Fe}^{2+} + 2e^-$). Because elemental iron donates electrons, it acts as the reducing agent.
Exercise 2: A student immerses a clean strip of metallic zinc into an aqueous solution of aluminum sulfate. Predict whether a single displacement reaction will occur, providing an explanation based on electrochemical principles and the reactivity series.
Solution 2: No single displacement reaction will occur when a zinc strip is immersed in aluminum sulfate solution ($\text{Zn}(s) + \text{Al}_2(\text{SO}_4)_3(aq) \to \text{No Reaction}$). In the electrochemical reactivity series, aluminum lies higher than zinc, meaning aluminum has a higher oxidation potential, lower ionization energy, and greater tendency to lose valence electrons than zinc. Because aluminum is more electropositive than zinc, zinc cannot reduce aluminum cations ($Al^{3+}$) to elemental aluminum metal.
Exercise 3: Explain why a freshly cut strip of aluminum metal often displays an initial induction delay of several minutes before it begins displacing copper from a blue copper(II) sulfate solution, whereas a polished zinc strip reacts almost immediately.
Solution 3: When exposed to atmospheric air, aluminum reacts spontaneously with oxygen to form a tough, continuous, chemically resistant layer of aluminum oxide ($Al_2O_3$) on its surface. This oxide film passivates the metal, acting as a physical barrier that prevents aqueous copper(II) ions from contacting the underlying metallic aluminum. The reaction cannot proceed until the aqueous solution penetrates or dissolves this protective film (often aided by sulfate or trace chloride ions). Once this protective barrier is breached, the displacement reaction proceeds vigorously. In contrast, the natural oxide film that forms on zinc is more porous and loosely bound. The zinc oxide dissolves quickly in aqueous solution, allowing copper ions to reach the metallic zinc surface and react almost immediately.