Electrochemistry & Corrosion Worksheet for 15-Year-Olds (Mel Science Kits)

Instructions

Read each question carefully. Use the knowledge you gained from your Chemistry & Electricity and Corrosion experiments to answer the questions below. Think about the underlying scientific principles, not just the steps you followed.

Part 1: Generating Electricity

In your first set of experiments, you created galvanic cells, which are commonly known as batteries. These devices convert chemical energy into electrical energy through redox reactions.

1. The Lemon Battery: In the lemon battery experiment, you used a lemon and two different metal strips (like copper and zinc or magnesium) to power an LED.
   a) What is the role of the lemon juice in this setup?
   b) Why is it essential to use two different metals for the electrodes?


2. The Daniell Cell: You also constructed a more formal galvanic cell using a zinc strip in a zinc sulfate solution and a copper strip in a copper(II) sulfate solution.
   a) Between zinc and copper, which metal is more easily oxidized (loses electrons)? This electrode is called the anode.
   b) Write the oxidation half-reaction that occurs at the anode.
   c) At the other electrode (the cathode), copper(II) ions (Cu²⁺) from the solution are reduced. Write the reduction half-reaction that occurs at the cathode.
   d) In which direction do the electrons flow through the external wire connecting the two metals? (e.g., from zinc to copper, or from copper to zinc?)


3. Critical Thinking: The reactivity of a metal determines its tendency to lose electrons. The general reactivity series for the metals you used is: Magnesium > Zinc > Iron > Copper. If you built a galvanic cell using a magnesium strip and a copper strip, how would you expect its voltage to compare to the zinc/copper cell? Explain your answer.

Part 2: The Chemistry of Corrosion

Corrosion, such as the rusting of iron, is an unwanted electrochemical process. Your experiments explored how rust forms and how it can be prevented.

1. Anatomy of Rust: For an iron nail to rust, it needs to be exposed to two key substances from the environment. What are they?


2. Visualizing Corrosion: In one experiment, you used chemical indicators to observe the electrochemical process of rusting in a Petri dish.
   a) The indicator potassium hexacyanoferrate(III) turns a deep blue color (Prussian blue) when it reacts with the Fe²⁺ ions that form as iron oxidizes. Is the blue region that appears the anode or the cathode?
   b) The indicator phenol red turns pink in the presence of hydroxide ions (OH⁻). These ions are produced where reduction occurs. Is the pink region that appears the anode or the cathode?


3. Sacrificial Protection: One way to protect iron is to attach a more reactive metal to it. This is called cathodic protection or using a sacrificial anode. Explain how attaching a strip of magnesium to an iron nail prevents the iron from rusting, even when both are in a corrosive environment.

Part 3: Tying It All Together - Redox Reactions

Both generating electricity in a battery and the process of corrosion are governed by oxidation-reduction (redox) reactions.

1. Defining the Terms: A common mnemonic for remembering the definitions of oxidation and reduction is LEO the lion says GER. What does this stand for?


2. Identifying Agents: Look again at the overall balanced equation for the Daniell cell:
   Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
   a) Which substance is the reducing agent (it causes the other substance to be reduced)?
   b) Which substance is the oxidizing agent (it causes the other substance to be oxidized)?

Answer Key

Part 1: Generating Electricity

1. a) The lemon juice (which contains citric acid and water) acts as the electrolyte. It provides ions that can move between the electrodes to complete the electrical circuit.
   b) It is essential to use two different metals because they have different tendencies to lose electrons (different electrode potentials). This difference in potential is what "pushes" the electrons to flow, creating an electric current. If you used two identical metals, there would be no potential difference and no current.
2. a) Zinc (Zn) is more easily oxidized. It is more reactive than copper.
   b) Anode (Oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻
   c) Cathode (Reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)
   d) Electrons flow from the site of oxidation (the anode) to the site of reduction (the cathode). Therefore, they flow from zinc to copper.
3. The voltage would increase. The voltage of a galvanic cell depends on the difference in reactivity between the two metals. Magnesium is much more reactive than zinc (it loses its electrons even more readily). Therefore, the potential difference between magnesium and copper is greater than the potential difference between zinc and copper, resulting in a higher voltage.

Part 2: The Chemistry of Corrosion

1. For iron to rust, it requires both oxygen (O₂) and water (H₂O).
2. a) The blue region indicates the presence of Fe²⁺, which means the iron is losing electrons (oxidizing). Oxidation occurs at the anode.
   b) The pink region indicates the formation of OH⁻ ions, which is the result of the reduction of oxygen. Reduction occurs at the cathode.
3. Magnesium is significantly more reactive than iron. When both metals are present in a corrosive environment, the more reactive magnesium will be preferentially oxidized (it will "sacrifice" itself). It becomes the anode, losing electrons, while the iron is forced to become the cathode. This prevents the iron from losing its own electrons and thus prevents it from rusting.

Part 3: Tying It All Together - Redox Reactions

1. LEO says GER stands for: Lose Electrons Oxidation, Gain Electrons Reduction.
2. a) The reducing agent is the substance that gets oxidized. In this reaction, Zinc (Zn) is oxidized, so it is the reducing agent.
   b) The oxidizing agent is the substance that gets reduced. In this reaction, the Copper(II) ion (Cu²⁺) is reduced, so it is the oxidizing agent.