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Chemistry: Changes to materials

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Chemistry: Changes to materials
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In this chapter

  1. 1The Reactivity Series of Metals
  2. 2Understanding Displacement Reactions
  3. 3Conservation of Mass
  4. 4Preparing Soluble Salts
  5. 5Chemical Tests and Deductions

1. The Reactivity Series of Metals

The reactivity series is a list that ranks metals in order of their reactivity, from most reactive at the top to least reactive at the bottom. A metal's reactivity is its tendency to form positive ions by losing electrons. This order is determined by observing how vigorously metals react with substances like water and dilute acids. For example, potassium reacts explosively with water, while gold does not react at all. Hydrogen and carbon are often included in the series as important reference points, particularly for extracting metals from their ores.

More Reactive Metal + Cold Water → Metal Hydroxide + Hydrogen

Less Reactive Metal + Steam → Metal Oxide + Hydrogen

Metal (above Hydrogen) + Dilute Acid → Salt + Hydrogen

Key term

Reactivity Series: A list of metals arranged in order of their reactivity, from the most reactive to the least reactive.

Examiner insight

Examiners look for precise observations. Instead of 'a reaction happened', use specific terms like 'fizzing', 'colour change', or 'solid dissolved'.

Common pitfall

Forgetting that hydrogen is the benchmark for reactions with acids. Metals below hydrogen in the series (like copper, silver, gold) will not react with dilute acids like HCl or H₂SO₄.

Worked example 14 marks

A small piece of calcium is added to a beaker of water. Describe two observations you would expect to see and write a balanced symbol equation for the reaction.

  1. 1

    Step 1: Identify the reactants and their positions in the reactivity series. Calcium is a reactive metal, placed well above hydrogen.

  2. 2

    Step 2: Predict the products based on the general equation: Metal + Water → Metal Hydroxide + Hydrogen. The products will be calcium hydroxide and hydrogen gas.

  3. 3

    Step 3: Describe the expected observations. The production of hydrogen gas will cause fizzing or bubbles. The formation of calcium hydroxide, which is slightly soluble, may make the water cloudy or milky. The reaction is exothermic, so the beaker might get warm.

  4. 4

    Step 4: Write the balanced symbol equation. Calcium is Ca, water is H₂O, calcium hydroxide is Ca(OH)₂, and hydrogen is H₂. The balanced equation is: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g).

Recap

  • The reactivity series orders metals by how readily they react.
  • Potassium is at the top (most reactive) and gold is at the bottom (least reactive).
  • A metal's position determines if it will react with water or acids.
  • Metals above hydrogen in the series will displace hydrogen from dilute acids.
  • Hydrogen and carbon are non-metals included for comparison.

Quick check

  1. Place the following metals in order of decreasing reactivity: Iron, Gold, Magnesium, Copper.2 marks

2. Understanding Displacement Reactions

A displacement reaction is a chemical reaction where a more reactive element takes the place of a less reactive element in a compound. In the context of metals, a more reactive metal will displace a less reactive metal from its salt solution. For example, if you place a piece of zinc (more reactive) into a blue solution of copper(II) sulfate, the zinc will displace the copper. The blue colour fades as colourless zinc sulfate is formed, and a layer of reddish-brown copper metal is deposited.

More Reactive Metal + Salt of Less Reactive Metal → Salt of More Reactive Metal + Less Reactive Metal

Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Key term

Displacement Reaction: A reaction in which a more reactive element takes the place of a less reactive element in one of its compounds.

Examiner insight

High-scoring answers always include an explicit comparison of the metals' positions in the reactivity series to justify why a displacement reaction does or does not occur.

Fun fact

The spectacular Thermite reaction, where aluminium displaces iron from iron(III) oxide, produces so much heat that the iron becomes molten. It's used to weld thick steel, like railway tracks, in remote locations.

Worked example 14 marks

A strip of magnesium is placed into a beaker of iron(II) chloride solution. Predict whether a reaction will occur. If so, describe the changes you would see and write a balanced symbol equation.

  1. 1

    Step 1: Locate magnesium and iron in the reactivity series. Magnesium is more reactive than iron.

  2. 2

    Step 2: Since magnesium is more reactive, it will displace iron from the iron(II) chloride solution. A reaction will occur.

  3. 3

    Step 3: Describe the observations. The magnesium strip will dissolve or get smaller. A dark grey solid (iron) will form on the surface of the magnesium or in the solution. The pale green solution of iron(II) chloride will become colourless as magnesium chloride is formed.

  4. 4

    Step 4: Write the balanced symbol equation. Reactants are Mg and FeCl₂. Products are MgCl₂ and Fe. The equation is already balanced: Mg(s) + FeCl₂(aq) → MgCl₂(aq) + Fe(s).

Recap

  • A more reactive metal can displace a less reactive metal from its salt solution.
  • If the metal is less reactive than the one in the salt, no reaction happens.
  • The reactivity series is essential for predicting the outcome of displacement reactions.
  • Observations often include a colour change in the solution and the formation of a solid metal deposit.

Quick check

  1. Would you expect a reaction if a copper wire is placed in a solution of silver nitrate? Explain your answer.2 marks

3. Conservation of Mass

The Law of Conservation of Mass is a fundamental principle in chemistry. It states that in a closed system, the total mass of the substances before a chemical reaction (the reactants) is equal to the total mass of the substances after the reaction (the products). Atoms are not created or destroyed during a reaction; they are simply rearranged to form new substances. Sometimes, mass might appear to be 'lost' or 'gained'. This is usually because the reaction is in an open container and a gas has either escaped (e.g., CO₂ in a reaction with a carbonate) or been used up from the air (e.g., oxygen in combustion).

Total mass of reactants = Total mass of products

Key term

Conservation of Mass: The principle that mass cannot be created or destroyed in a chemical reaction, so the total mass of reactants equals the total mass of products.

Examiner insight

Examiners frequently test understanding of conservation of mass by asking students to explain why the mass measured on a balance goes down (gas escapes) or up (gas from the air reacts).

Common pitfall

Mistaking an apparent change in mass in an open system for a violation of the law of conservation of mass. Always consider if a gas is involved.

Worked example 12 marks

In an experiment, 5.6 g of iron filings react completely with 3.2 g of sulfur powder when heated. What is the mass of the iron sulfide produced?

  1. 1

    Step 1: Identify the reactants and products. Reactants are iron and sulfur. The product is iron sulfide.

  2. 2

    Step 2: Apply the Law of Conservation of Mass. Total mass of reactants = Total mass of products.

  3. 3

    Step 3: Calculate the total mass of reactants. Mass = Mass of iron + Mass of sulfur = 5.6 g + 3.2 g = 8.8 g.

  4. 4

    Step 4: State the mass of the product. The mass of iron sulfide produced is 8.8 g.

Worked example 23 marks

A student adds 10 g of calcium carbonate to an open beaker containing 100 g of hydrochloric acid. After the fizzing stops, the total mass of the beaker and its contents is 105.6 g. Explain this result with reference to the conservation of mass.

  1. 1

    Step 1: Calculate the initial total mass. Initial mass = 10 g + 100 g = 110 g.

  2. 2

    Step 2: Compare the initial and final masses. The final mass (105.6g) is less than the initial mass (110 g). The apparent mass loss is 110 g - 105.6 g = 4.4 g.

  3. 3

    Step 3: Explain the observation. The reaction between calcium carbonate and hydrochloric acid produces carbon dioxide gas. Since the beaker is open, this gas escapes into the surroundings.

  4. 4

    Step 4: Link to conservation of mass. The 'lost' mass is actually the mass of the CO₂ gas that escaped. Mass is still conserved; the mass of the products (solution + gas) would equal the mass of the reactants if the gas were collected.

Recap

  • Mass is always conserved in any chemical reaction.
  • Total mass of reactants equals the total mass of products.
  • Atoms are rearranged during a reaction, not lost or gained.
  • Apparent mass changes in open systems are due to gases escaping or being consumed from the air.
  • Balancing symbol equations ensures the conservation of atoms.

Quick check

  1. If 2.4 g of magnesium burns in air to produce 4.0 g of magnesium oxide, what mass of oxygen was used in the reaction?1 mark

4. Preparing Soluble Salts

A salt is a compound formed when the hydrogen in an acid is replaced by a metal. There are several ways to make a soluble salt in the lab. The method chosen depends on the reactants. Common methods include:

  1. Acid + Reactive Metal: Produces salt and hydrogen gas. This is only suitable for metals that are reactive enough to react with acid but not dangerously so (e.g., zinc, magnesium, iron).
  2. Acid + Insoluble Base (a metal oxide or hydroxide): Produces salt and water. This is a very common and safe method, especially for making salts of less reactive metals like copper.
  3. Acid + Carbonate (metal carbonate): Produces salt, water, and carbon dioxide gas.

A typical preparation involves adding an excess of the solid reactant (e.g., copper(II) oxide) to the acid, filtering off the unreacted solid, and then crystallising the salt from the resulting solution by heating it to evaporate some water and then allowing it to cool.

Acid + Metal → Salt + Hydrogen

Acid + Insoluble Base → Salt + Water

Acid + Carbonate → Salt + Water + Carbon Dioxide

Key term

Salt: An ionic compound formed in a neutralisation reaction between an acid and a base, or when the hydrogen in an acid is replaced by a metal.

Examiner insight

For salt preparation questions, marks are awarded for each logical step in the practical procedure. Mentioning specific apparatus (beaker, filter funnel, evaporating basin) and techniques (stirring, warming, filtering, cooling) is crucial.

Common pitfall

Confusing the gases produced. Remember: Metal + Acid → Hydrogen ('squeaky pop' test). Carbonate + Acid → Carbon Dioxide (turns limewater milky).

Worked example 16 marks

Describe a method to prepare a pure, dry sample of copper(II) sulfate crystals, starting from copper(II) oxide and sulfuric acid.

  1. 1

    Step 1: Gently warm a beaker of dilute sulfuric acid. Do not boil.

  2. 2

    Step 2: Add copper(II) oxide powder to the warm acid a little at a time, stirring continuously, until no more powder dissolves. This means the copper(II) oxide is in excess and all the acid has reacted.

  3. 3

    Step 3: Filter the mixture using filter paper and a funnel to remove the excess unreacted black copper(II) oxide. The blue copper(II) sulfate solution is collected as the filtrate.

  4. 4

    Step 4: Pour the filtrate into an evaporating basin and heat it gently to evaporate about half of the water. This concentrates the solution.

  5. 5

    Step 5: Leave the concentrated solution to cool slowly. Blue crystals of copper(II) sulfate will form.

  6. 6

    Step 6: Once crystallisation is complete, filter off the crystals, wash them with a small amount of cold distilled water, and then dry them between sheets of filter paper.

Recap

  • Soluble salts can be made by reacting an acid with a metal, a base, or a carbonate.
  • Using an excess of an insoluble reactant ensures all the acid is used up.
  • Filtration is used to separate an insoluble solid from a liquid.
  • Crystallisation is used to obtain a pure solid salt from its solution.
  • The final crystals should be washed and dried.

Quick check

  1. When making a salt from an acid and an insoluble base, why is it important to add the base in excess?1 mark

5. Chemical Tests and Deductions

In chemistry, you often act like a detective, using observations and simple tests to identify unknown substances. Key tests include identifying gases produced in reactions. Hydrogen gas, typically made from a reactive metal and acid, gives a 'squeaky pop' when a lit splint is introduced. Carbon dioxide, made from an acid and a carbonate, turns limewater (calcium hydroxide solution) from clear to cloudy or milky. By combining these tests with knowledge of reaction types (like displacement or neutralisation), you can deduce the identities of reactants and products.

Test for Hydrogen: 2H₂(g) + O₂(g) → 2H₂O(l) (causes a 'squeaky pop')

Test for Carbon Dioxide: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l) (forms a white precipitate)

Key term

Limewater: A saturated aqueous solution of calcium hydroxide, Ca(OH)₂, used as the standard chemical test for carbon dioxide.

Examiner insight

Examiners reward clear, logical answers that describe the test, the expected positive result, and the conclusion that can be drawn from it for each substance being identified.

Worked example 16 marks

A student has three unlabelled powders: magnesium, magnesium oxide, and magnesium carbonate. They have access to dilute hydrochloric acid and limewater. Describe a series of tests to identify each powder.

  1. 1

    Step 1: Plan the procedure. Add a small sample of each powder to three separate test tubes. Then add dilute hydrochloric acid to each.

  2. 2

    Step 2: Predict the reaction of magnesium carbonate. It will react with the acid to produce bubbles of gas. If this gas is bubbled through limewater, the limewater will turn milky. This identifies the magnesium carbonate.

  3. 3

    Step 3: Predict the reaction of magnesium. It will react with the acid to produce bubbles of gas. If this gas is collected and tested with a lit splint, it will give a 'squeaky pop'. This identifies the magnesium.

  4. 4

    Step 4: Predict the reaction of magnesium oxide. It is a base and will react with the acid, but it will not produce any gas. The solid will dissolve to form a colourless solution. This identifies the magnesium oxide.

  5. 5

    Step 5: Summarise the identification. The one that makes a gas that turns limewater milky is magnesium carbonate. The one that makes a gas that gives a squeaky pop is magnesium. The one that reacts without producing a gas is magnesium oxide.

Recap

  • Hydrogen gas is identified by the 'squeaky pop' test with a lit splint.
  • Carbon dioxide gas is identified by the 'milky limewater' test.
  • Oxygen gas (not covered here, but useful) relights a glowing splint.
  • By observing reactions and testing products, you can identify unknown chemicals.
  • Linking observations to chemical principles is a key skill.

Quick check

  1. A gas produced in a reaction is bubbled through limewater, which remains clear. What can you conclude about the gas?1 mark

End-of-chapter exercise

Test yourself on the whole chapter. Work through these before moving on.

  1. Define the term 'displacement reaction' and give an example using a word equation.2 marks
  2. A student adds a piece of zinc metal to a test tube containing silver nitrate solution. Explain what would be observed and why, with reference to the reactivity series.3 marks
  3. Write the balanced symbol equation for the reaction between magnesium and sulfuric acid.2 marks
  4. A 10.0 g piece of magnesium ribbon is heated in a crucible. The final mass of the white powder (magnesium oxide) formed is 16.6 g. Calculate the mass of oxygen that reacted with the magnesium.2 marks
  5. Describe, in a series of numbered steps, how you would prepare a pure, dry sample of zinc sulfate crystals starting from solid zinc carbonate and a suitable acid.6 marks
  6. When calcium carbonate reacts with hydrochloric acid in an open flask on a balance, the mass on the balance decreases. Explain why this happens, even though mass is conserved in a chemical reaction.4 marks
  7. A metal 'M' displaces copper from copper(II) sulfate solution but does not react with a solution of zinc sulfate. What can you deduce about the position of metal 'M' in the reactivity series? Explain your reasoning.3 marks
  8. An unknown powdered substance 'X' is added to dilute nitric acid. A gas is evolved which gives a 'squeaky pop' with a lit splint. The resulting solution is evaporated to produce a crystalline salt. Identify the type of substance 'X' is, the gas produced, and the general name for the crystalline salt.3 marks
  9. Iron is extracted from its ore, iron(III) oxide (Fe₂O₃), by heating it with carbon in a blast furnace. The reaction is: 2Fe₂O₃ + 3C → 4Fe + 3CO₂. Explain why this is considered a displacement reaction.2 marks
  10. A student reacts 1.3 g of zinc powder with 50 cm³ of 1.0 mol/dm³ copper(II) sulfate solution. The temperature increases from 20°C to 45°C. When the experiment is repeated with 1.3 g of iron powder, the temperature only increases to 28°C. Explain this difference in temperature change in terms of the reactivity of the metals.4 marks

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