Metals and the Reactivity Series

Welcome to this comprehensive lesson on Metals and the Reactivity Series, designed specifically for CXC Chemistry students. This lesson covers all aspects required by the 2024-2025 syllabus.

Learning Objectives

By the end of this lesson, you should be able to:

1. Introduction to Metals

Metals make up approximately 75% of all known elements and play crucial roles in our daily lives, from construction to electronics.

1.1 Physical Properties of Metals

Metals share several distinctive physical properties that make them valuable for various applications:

1.2 Chemical Properties of Metals

Chemically, metals exhibit these key properties:

2. The Reactivity Series

The reactivity series is an arrangement of metals in order of their reactivity, from the most reactive to the least reactive. It helps us predict how metals will behave in various chemical reactions.

Element Reactivity Potassium (K) Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Carbon (C)* Zinc (Zn) Iron (Fe) Tin (Sn) Lead (Pb) Hydrogen (H)* Copper (Cu) Silver (Ag) Gold (Au) Platinum (Pt) * Carbon and hydrogen are non-metals included for reference Decreasing reactivity

2.1 The Complete Reactivity Series

From most reactive to least reactive:

Important Note: Carbon and hydrogen are not metals but are included in the reactivity series for reference purposes as they help determine how metals react with certain compounds.

2.2 Categories within the Reactivity Series

For study purposes, we can divide the reactivity series into categories:

3. Reactions of Metals with Water

The reactivity of a metal determines how it reacts with water. This is one way to determine a metal's position in the reactivity series.

3.1 Patterns of Reactivity with Water

Metal Category Reaction with Water Products Examples
Very reactive metals React violently with cold water Metal hydroxide + Hydrogen gas Potassium, Sodium, Calcium
Reactive metals React with steam but not cold water Metal oxide + Hydrogen gas Magnesium, Aluminium, Zinc
Moderately reactive metals React slowly with steam Metal oxide + Hydrogen gas Iron, Tin, Lead
Less reactive metals Do not react with water or steam No reaction Copper, Silver, Gold, Platinum

3.2 Chemical Equations

Here are the balanced equations for some common reactions with water:

Experiment: Reaction of Metals with Water

Aim: To observe and compare the reactions of different metals with water.

Materials needed:

Procedure:

  1. Place a small piece of sodium in cold water (CAUTION: Stand back!)
  2. Place a small piece of calcium in cold water
  3. Place magnesium, zinc, and copper in separate test tubes with cold water
  4. For metals that don't react with cold water, try passing steam over heated samples
  5. Add a few drops of phenolphthalein to the water after reaction

Observations:

4. Reactions of Metals with Acids

The reaction of metals with acids is another important indicator of their reactivity. Most metals that are more reactive than hydrogen will displace hydrogen from acids.

4.1 General Pattern

When a metal reacts with an acid, it usually produces a salt and hydrogen gas:

Metal + Acid → Metal Salt + Hydrogen

4.2 Reactivity Pattern with Acids

Metal Category Reaction with Dilute Acids Products Examples
Very reactive metals React vigorously or explosively Metal salt + Hydrogen gas Potassium, Sodium, Calcium
Reactive metals React readily Metal salt + Hydrogen gas Magnesium, Aluminium, Zinc
Moderately reactive metals React more slowly Metal salt + Hydrogen gas Iron, Tin, Lead
Less reactive metals Do not react with most acids* No reaction Copper, Silver, Gold, Platinum

*Exception: Copper will react with nitric acid (HNO3) but this is an oxidizing acid and the reaction is not a simple displacement of hydrogen.

4.3 Chemical Equations

Here are balanced equations for some common reactions with hydrochloric acid:

Experiment: Reactivity of Metals with Acids

Aim: To compare the rates of reaction of different metals with dilute hydrochloric acid.

Materials needed:

Procedure:

  1. Place equal-sized pieces of each metal in separate test tubes
  2. Add equal volumes of dilute hydrochloric acid to each test tube
  3. Start timing and observe the rate of hydrogen gas production
  4. Record the time taken for the reaction to complete (or after 10 minutes)

Expected Results:

5. Reactions of Metals with Oxygen

When metals react with oxygen (usually from air), they form metal oxides. The ease and speed of this reaction also correlate with the metal's position in the reactivity series.

5.1 General Pattern

The general equation for metal oxidation is:

Metal + Oxygen → Metal Oxide

5.2 Reactivity Pattern with Oxygen

Metal Category Reaction with Oxygen Products Examples
Very reactive metals React vigorously, may burn with characteristic flames Metal oxide Potassium, Sodium, Calcium
Reactive metals React readily when heated Metal oxide Magnesium, Aluminium, Zinc
Moderately reactive metals React slowly when heated Metal oxide Iron, Tin, Lead
Less reactive metals React only at high temperatures or not at all Metal oxide (if reaction occurs) Copper, Silver, Gold, Platinum

5.3 Chemical Equations

Here are balanced equations for some common reactions with oxygen:

Metal Oxidation Comparison Sodium Stored in oil Magnesium Burns with bright flame Iron Forms rust (slowly) Copper Forms green patina (very slow) After Exposure to Oxygen Na₂O (white) MgO (white) Fe₂O₃ (rust) CuCO₃ (patina)

6. Displacement Reactions

Displacement reactions occur when a more reactive metal displaces a less reactive metal from its compound in solution. These reactions provide direct evidence for the relative positions of metals in the reactivity series.

6.1 Single Displacement Reactions

The general equation for a single displacement reaction is:

Metal A + Metal B Compound → Metal A Compound + Metal B

This reaction will only occur if Metal A is more reactive than Metal B.

6.2 Examples of Displacement Reactions (continued)

Important: The following reactions will NOT occur because the first metal is less reactive than the second:

Displacement Reaction Example Copper ion solution (blue) + Zinc strip added → Zinc ion solution (colorless) Copper metal deposits Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Experiment: Metal Displacement Reactions

Aim: To investigate displacement reactions and verify the reactivity series.

Materials needed:

Procedure:

  1. Label five test tubes for each metal salt solution
  2. Place each metal piece in all five solutions (25 combinations total)
  3. Record any color changes or evidence of reactions after 10-15 minutes

Expected Results:

Conclusion: The results can be used to establish a reactivity series: Mg > Zn > Fe > Pb > Cu

7. Applications of the Reactivity Series

Understanding the reactivity series has many practical applications in chemistry and industry.

7.1 Extraction of Metals

The method used to extract a metal from its ore depends on its position in the reactivity series:

Metal Reactivity Extraction Method Examples
Very reactive metals Electrolysis Potassium, Sodium, Calcium
Reactive metals Electrolysis or heat with carbon Magnesium, Aluminium
Moderately reactive metals Reduction with carbon Zinc, Iron, Tin, Lead
Less reactive metals Found native or simple heat Copper, Silver, Gold, Platinum

7.2 Electrolysis

Electrolysis is used to extract the most reactive metals like sodium and aluminium:

7.3 Carbon Reduction

Moderately reactive metals like zinc and iron are extracted by reduction with carbon:

7.4 Heat Treatment

Less reactive metals like mercury can be extracted by simple heating:

7.5 Finding Native

The least reactive metals (gold, platinum) are found naturally in their elemental form (native).

Metal Extraction Methods Most reactive Least reactive Electrolysis Carbon Reduction Heat/Native K, Na, Ca, Mg, Al Zn, Fe, Sn, Pb Cu, Hg, Ag, Au, Pt Cathode (-) Anode (+) Electrolysis Carbon Reduction Native/Heat

8. Corrosion and Protection

Corrosion is the deterioration of metals through reaction with their environment, particularly oxygen. The reactivity series can help predict how susceptible a metal is to corrosion.

8.1 Corrosion of Metals

More reactive metals tend to corrode more easily. Common examples include:

8.2 Protection Methods

Several methods are used to protect metals from corrosion:

8.2.1 Barrier Methods

8.2.2 Sacrificial Protection

8.2.3 Electroplating

Important concept: In sacrificial protection, the more reactive metal (higher in the reactivity series) corrodes preferentially, protecting the less reactive metal.

9. Alloys and Their Importance

Alloys are mixtures of a metal with other elements (usually other metals) to improve properties.

9.1 Common Alloys

Alloy Composition Properties/Uses
Brass Copper + Zinc Harder than copper, used for decorative items, musical instruments
Bronze Copper + Tin Harder than copper, used for statues, medals
Steel Iron + Carbon (small amount) Stronger and less brittle than iron, used in construction
Stainless Steel Iron + Chromium + Nickel Resistant to corrosion, used for cutlery, surgical instruments
Duralumin Aluminium + Copper + Magnesium Lightweight but strong, used in aircraft construction
Solder Lead + Tin Low melting point, used for joining electrical components

9.2 Advantages of Alloys

10. Glossary of Key Terms

Alloy
A mixture of a metal with other elements (usually metals) that has improved properties.
Anion
A negatively charged ion that is attracted to the anode during electrolysis.
Cathode
The negative electrode in electrolysis where reduction occurs and metals are deposited.
Cation
A positively charged ion that is attracted to the cathode during electrolysis.
Corrosion
The gradual destruction of metals by chemical reactions with substances in their environment.
Displacement reaction
A reaction where a more reactive element displaces a less reactive element from its compound.
Ductility
The ability of a material to be drawn into wire without breaking.
Electrolysis
The process of using electricity to drive a non-spontaneous chemical reaction.
Galvanizing
The process of coating iron or steel with a layer of zinc to prevent rusting.
Luster
The way light reflects off a surface, with metals typically having a shiny appearance.
Malleability
The ability of a material to be hammered or pressed into shape without breaking.
Native metal
A metal found in nature in its elemental form, rather than as a compound.
Ore
A naturally occurring rock or sediment containing minerals with economically important elements.
Oxidation
The loss of electrons by a substance during a chemical reaction.
Patina
A green or brown film that forms on the surface of bronze or similar metals due to oxidation.
Reactivity series
An arrangement of metals in order of their reactivity, from most reactive to least reactive.
Reduction
The gain of electrons by a substance during a chemical reaction.
Rusting
The corrosion of iron in the presence of oxygen and moisture to form hydrated iron(III) oxide.
Sacrificial protection
A method of protecting a metal by connecting it to a more reactive metal that corrodes preferentially.

11. Self-Assessment Questions

Multiple Choice Questions

  1. Which of the following metals is most reactive?
    1. Copper
    2. Zinc
    3. Magnesium
    4. Iron
  2. Metal X displaces metal Y from its salt solution, but does not displace metal Z. This suggests that:
    1. X is more reactive than Y but less reactive than Z
    2. X is more reactive than both Y and Z
    3. X is less reactive than both Y and Z
    4. X is less reactive than Y but more reactive than Z
  3. Which metal would react most vigorously with cold water?
    1. Lead
    2. Zinc
    3. Calcium
    4. Copper
  4. Which extraction method is most suitable for aluminium?
    1. Heating the oxide alone
    2. Reduction with carbon
    3. Electrolysis
    4. Found native in the earth
  5. When iron nails are placed in copper(II) sulfate solution:
    1. No reaction occurs
    2. The solution turns colorless and copper deposits form
    3. Hydrogen gas is produced
    4. The iron dissolves without any visible change

Short Answer Questions

  1. Explain why gold and platinum are found native in the earth while potassium and sodium are not.
  2. Describe the reaction between magnesium and steam, writing a balanced chemical equation.
  3. Why does zinc protect iron from rusting in galvanized iron? Explain using the reactivity series.
  4. A student places strips of copper, zinc, and magnesium into separate solutions of iron(II) sulfate. Predict and explain what would happen in each case.
  5. Explain why different methods are used to extract different metals from their ores.

Application Questions

  1. Design an experiment to determine the relative positions of four unknown metals in the reactivity series.
  2. A ship's hull is made of steel (iron) and has a brass propeller. Explain why the steel hull might corrode faster than expected.
  3. Aluminium is very reactive, yet aluminium cans do not corrode rapidly. Explain this apparent contradiction.
  4. In the extraction of iron, carbon monoxide can reduce iron(III) oxide to iron. Write balanced equations for this process and explain why it works in terms of reactivity.
  5. Suggest three different methods for protecting an iron bridge from corrosion, explaining the chemistry behind each method.

Multiple Choice Answers

  1. c) Magnesium - It is higher in the reactivity series than the other metals listed.
  2. a) X is more reactive than Y but less reactive than Z - This is the only scenario that fits the observed displacement reactions.
  3. c) Calcium - It is the most reactive of the metals listed and will react vigorously with cold water.
  4. c) Electrolysis - Aluminium is too reactive to be extracted by carbon reduction or heating alone.
  5. b) The solution turns colorless and copper deposits form - Iron displaces copper from the solution because iron is more reactive.

Short Answer Answers

  1. Gold and platinum are very unreactive metals (low in the reactivity series) and do not readily form compounds with elements in their environment. Potassium and sodium are extremely reactive (high in the reactivity series) and readily react with oxygen, water, and other elements in the environment, so they are always found as compounds in nature.
  2. Magnesium reacts with steam to produce magnesium oxide and hydrogen gas. The balanced equation is: Mg(s) + H2O(g) → MgO(s) + H2(g). This reaction occurs because magnesium is more reactive than hydrogen and can displace it from water when the reaction is activated by heat.
  3. Zinc protects iron by acting as a sacrificial metal. Since zinc is more reactive than iron (higher in the reactivity series), it oxidizes preferentially, forming zinc ions and releasing electrons. These electrons flow to the iron, preventing the iron from oxidizing (rusting). Even if the zinc coating is scratched, the zinc will continue to corrode instead of the iron.
  4. Copper: No reaction would occur because copper is less reactive than iron. Zinc: The zinc would displace iron from the solution, forming zinc sulfate and iron metal deposits. Zn(s) + FeSO4(aq) → ZnSO4(aq) + Fe(s) Magnesium: The magnesium would displace iron from the solution, forming magnesium sulfate and iron metal deposits. Mg(s) + FeSO4(aq) → MgSO4(aq) + Fe(s). The reaction with magnesium would be more vigorous than with zinc.
  5. Different extraction methods are used because metals have different reactivities. Very reactive metals like potassium and sodium require electrolysis because they bond too strongly with non-metals to be reduced by carbon. Moderately reactive metals like iron and zinc can be extracted using carbon reduction because carbon is more reactive with oxygen than these metals. Less reactive metals like copper and silver can be extracted by simply heating their oxides or may be found native because they don't readily form compounds in the first place.

Application Answers

  1. Design: Label the unknown metals W, X, Y, and Z. Step 1: Test each metal with water to identify those that react. Step 2: For metals that don't react with water, test with dilute acids. Step 3: For metals that don't react with acids, test with metal salt solutions through displacement reactions. Results: Metal W reacts vigorously with water, producing hydrogen gas and an alkaline solution, indicating it's likely potassium or sodium. Metal X reacts slowly with water, suggesting it may be calcium. Metal Y reacts with dilute acid but not water, indicating it could be zinc or iron. Metal Z doesn't react with water or acids, suggesting it's copper or silver. Conclusion: Based on these tests, we can place the metals in order of reactivity (from highest to lowest): W > X > Y > Z.
  2. To determine the best metal to use, I need to consider the reactivity series. The metal must be more reactive than iron to protect it through sacrificial protection, but not so reactive that it corrodes too quickly. From the options: Copper is less reactive than iron, so it would not provide protection. Zinc is more reactive than iron but not excessively reactive - this makes it ideal for galvanizing. Aluminum, while more reactive than zinc, forms a protective oxide layer that would prevent it from sacrificing itself. Magnesium is significantly more reactive than iron and would corrode too quickly in an outdoor environment. Silver is less reactive than iron and would not provide protection. The best choice is zinc because it sits at the right position in the reactivity series to provide effective sacrificial protection without corroding too rapidly.
  3. This phenomenon is due to the different positions of iron and copper in the reactivity series. Iron (Fe) is more reactive than copper (Cu), so when iron nails are placed in copper sulfate solution, a displacement reaction occurs. The iron atoms donate electrons to copper ions, causing solid copper to form on the nail surface while iron dissolves into the solution. The balanced equation is: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s). This demonstrates a key principle of the reactivity series: a more reactive metal will displace a less reactive metal from its compounds. This principle is used in metal extraction and in designing batteries where electron transfer between metals of different reactivity generates electric current.
  4. Calcium is very reactive, so it would need to be extracted by electrolysis of its molten chloride, requiring substantial electricity. Zinc is moderately reactive and can be extracted by carbon reduction of its oxide, which requires heating but less energy than electrolysis. Copper is low in the reactivity series and may be found native or can be extracted by simply heating its oxide. Therefore, copper has the lowest environmental impact in terms of extraction energy requirements, while calcium has the highest. To minimize environmental impact: (1) Use recycled metals instead of extracting new ones, (2) Develop more energy-efficient extraction technologies, (3) Use renewable energy sources for extraction processes, (4) Recover and reuse byproducts from extraction, and (5) Select appropriate metals for specific applications based on their properties and extraction requirements.
  5. When magnesium ribbon burns in air, it undergoes an oxidation reaction with oxygen to form magnesium oxide: 2Mg(s) + O2(g) → 2MgO(s). This reaction demonstrates the high reactivity of magnesium with oxygen. The intense white light produced shows the large amount of energy released during this exothermic reaction. Magnesium's high position in the reactivity series explains why it reacts so vigorously with oxygen. Modern applications of this reaction include flash photography, fireworks, and emergency flares. The reaction also demonstrates why magnesium structures require special protective coatings to prevent oxidation in air. In magnesium fire extinguishers, sand or dry powder must be used instead of water, as magnesium is reactive enough to split water into hydrogen and oxygen, potentially worsening the fire.