OCR GCSE Chemistry: Metals and Reactivity – Extraction and Electrolysis

GCSE Chemistry: Metals and Reactivity - Extraction and Electrolysis

Introduction

This tutorial explores the reactivity of metals, their extraction methods, and the concept of electrolysis. These topics are crucial for understanding how we obtain and utilize metals in our daily lives.

Reactivity Series

Metals differ in their reactivity, which determines how easily they react with other substances. We can arrange metals in a reactivity series based on their tendency to lose electrons and form positive ions.

• Highly reactive metals (e.g., potassium, sodium, lithium) readily lose electrons and react vigorously with water and acids.
• Less reactive metals (e.g., copper, silver, gold) are less likely to lose electrons and require stronger conditions for reactions.

Here is a simplified reactivity series (most reactive at the top):

Potassium Sodium Lithium Calcium Magnesium Aluminum Zinc Iron Tin Lead Copper Silver Gold

Key Observations:

• Metals higher in the series displace metals lower in the series from their compounds.
• Metals above hydrogen react with acids to produce hydrogen gas.
• Metals below hydrogen do not react with acids.

Metal Extraction

Metals are rarely found in their pure form in nature and need to be extracted from their ores. The extraction method depends on the metal's reactivity:

Extraction of Less Reactive Metals:

Carbon Reduction:

• Used for metals below carbon in the reactivity series (e.g., copper, iron).
• The ore is heated strongly with carbon (usually coke).
• Carbon acts as a reducing agent, removing oxygen from the metal oxide.
• Example: Extraction of Iron:
  • Iron ore (iron(III) oxide) is heated with coke in a blast furnace.
  • The chemical reaction: Fe?O? + 3C ? 2Fe + 3CO

Extraction of Highly Reactive Metals:

Electrolysis:

• Used for metals above carbon in the reactivity series (e.g., sodium, aluminium).
• Electrolysis involves using an electric current to break down a compound into its elements.
• Electrolyte: The molten or dissolved compound that conducts electricity.
• Electrodes: Conductive materials that allow the current to flow into and out of the electrolyte.
• Example: Electrolysis of Aluminium:
  • Aluminium oxide (Al?O?) is melted and electrolyzed.
  • At the cathode: Al³? + 3e? ? Al
  • At the anode: 2O²? ? O? + 4e?

Formation of Ionic Compounds

Metals react with non-metals to form ionic compounds. This occurs due to the transfer of electrons from metal atoms to non-metal atoms.

• Metal atoms lose electrons to form positive ions (cations).
• Non-metal atoms gain electrons to form negative ions (anions).
• The oppositely charged ions are attracted to each other, forming an ionic compound.

Example:

• Sodium (Na) reacts with chlorine (Cl?) to form sodium chloride (NaCl).
• Sodium atom loses one electron to form Na? ion.
• Chlorine atom gains one electron to form Cl? ion.
• The oppositely charged ions attract each other to form the ionic compound NaCl.

Practical Applications

• Metal displacement reactions: Demonstrate the reactivity series by observing which metals displace others from their solutions.
• Ionic equations: Use ionic equations to represent the reactions involved in metal extraction and displacement.
• Resource use and sustainability: Discuss the environmental impact of metal extraction and the need for sustainable practices.

Summary

This tutorial provides a foundation for understanding the reactivity of metals, their extraction methods, and the formation of ionic compounds. By applying these concepts, you can analyze the chemical processes involved in obtaining and utilizing metals in our society while considering the environmental implications.