Printable Worksheet for 15-Year-Olds — Chemistry & Electricity Labs (Mel Science Starter Kit: Electrochemistry & Corrosion)

Instructions

Read each question carefully. Use your observations from the Mel Science "Chemistry & Electricity" and "Corrosion" kits, along with your knowledge of electrochemistry, to answer the questions in the spaces provided. Think about the underlying scientific principles for each experiment.

Part 1: Generating Electricity - The Galvanic Cell

These questions relate to your experiments creating a lemon battery and a Daniell galvanic cell.

1. A galvanic cell (or voltaic cell) converts chemical energy into electrical energy. Based on your experiments, what are the three essential components that make up a simple galvanic cell?

 

2. Consider the Daniell cell you built using a zinc (Zn) strip in zinc sulfate solution and a copper (Cu) strip in copper(II) sulfate solution.
   • a) The metal that is more easily oxidized loses electrons and is called the anode. Which metal was the anode?

    

   • b) The metal where reduction occurs is called the cathode. Which metal was the cathode?

    

   • c) Electrons flow from the anode to the cathode. Describe the path of the electrons in your setup.

    

3. The lemon juice and the sulfate solutions are all examples of electrolytes. What is the primary function of an electrolyte in a working galvanic cell?

 

4. When building the lemon battery, you had to use two different metals (like copper and zinc) to make the LED light up. Explain why a "battery" made with two identical copper wires would not produce a current.

 

Part 2: Unwanted Electricity - The Science of Corrosion

These questions relate to your experiments with rusting iron nails and rust protection.

1. Corrosion, like the rusting of iron, is an electrochemical process.
   • a) What is the specific chemical process that happens to iron when it rusts (e.g., oxidation, reduction, neutralization)?

    

   • b) The familiar reddish-brown substance we call rust is a hydrated form of what iron compound? (Chemical name)

    

2. In your experiments, you used chemical indicators to visualize the electrochemical reactions. What did the formation of each color signify?
   • a) Blue color (from potassium hexacyanoferrate(III)):

    

   • b) Pink color (from phenol red):

    

3. You observed that wrapping an iron nail in a magnesium strip protected the iron from rusting. This is called "sacrificial protection." Explain why the more reactive magnesium "sacrifices" itself to protect the less reactive iron. Refer to the concepts of anode and cathode in your answer.

 

4. In the "Electricity vs. Iron" experiment, you used an external power source (AA batteries). When the iron nail was connected to the negative terminal, it was protected from corrosion. This is called cathodic protection. Why does forcing electrons onto the iron prevent it from rusting?

 

Part 3: Critical Thinking and Application

1. Real-World Chemistry: Large steel (mostly iron) ships have blocks of zinc attached to their underwater hulls to prevent them from rusting away in the saltwater. Based on what you learned, explain how these zinc blocks work. What is the scientific term for this type of protection?

 

2. Challenge Question: The tendency of a metal to be oxidized is described by its "activity." A more active metal will be the anode when paired with a less active metal. The activity series for the metals you used is: Magnesium (most active) > Zinc > Iron > Copper (least active).
   
   If you created a galvanic cell by wrapping an iron nail with a copper wire and placing it in a salt solution, which metal would corrode (rust) faster, the iron or the copper? Explain your reasoning.

 

Answer Key

Part 1: Generating Electricity - The Galvanic Cell

1. The three essential components are:
   • An anode (negative electrode)
   • A cathode (positive electrode)
   • An electrolyte (a substance containing ions that can move)
2. For the Daniell cell (Zn/Cu):
   • a) The anode was Zinc (Zn).
   • b) The cathode was Copper (Cu).
   • c) Electrons flowed from the zinc strip (anode), through the crocodile clip wire, to the copper strip (cathode).
3. The electrolyte's function is to complete the electrical circuit by allowing ions to flow between the two electrodes, balancing the charge as electrons flow through the external wire.
4. A current is generated by a difference in electrical potential between two different metals. This potential difference drives the flow of electrons. If you use two identical metals, there is no potential difference, so no electrons will flow, and no current is produced.

Part 2: Unwanted Electricity - The Science of Corrosion

1. • a) The process is oxidation (loss of electrons).
   • b) The chemical name for rust is hydrated iron(III) oxide.
2. • a) Blue color: Indicated the presence of iron(II) ions (Fe²⁺), meaning this was the site of iron oxidation (the anode).
   • b) Pink color: Indicated an increase in hydroxide ions (OH^-) and a more alkaline/basic environment. This was the site where oxygen was being reduced (the cathode).
3. Magnesium is more reactive (more active) than iron. When they are in contact in an electrolyte, magnesium becomes the anode and is preferentially oxidized (loses electrons). The iron becomes the cathode. Electrons flow from the magnesium to the iron, protecting the iron from being oxidized itself.
4. Rusting is the oxidation of iron (Fe → Fe²⁺ + 2e^-). By connecting the iron to the negative terminal, you are forcing a continuous supply of electrons onto it. This excess of electrons prevents the iron atoms from being able to lose their own electrons, thus inhibiting the oxidation process.

Part 3: Critical Thinking and Application

1. Zinc is more reactive (more active) than the iron in the steel hull. When in contact with the steel in saltwater (an electrolyte), the zinc acts as the anode and the steel hull acts as the cathode. The zinc block corrodes (is oxidized) instead of the ship's hull. This is a real-world example of sacrificial protection.
2. The iron would corrode much faster. Reasoning: According to the activity series, iron is more active than copper. Therefore, in an Fe-Cu galvanic cell, the iron will act as the anode and will be oxidized (rust), while the copper will act as the cathode. The presence of the less reactive copper actually accelerates the corrosion of the iron it is in contact with.