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Physical Science Enhancement - Chemistry 3

Oxygen – the ‘acid maker’

The name "oxygen" is derived from the Greek, meaning "acid maker". It was originally thought that all acids contain oxygen, whereas there are many that do not, for example hydrochloric acid, HCl.

However, non-metal oxides, such as carbon dioxide, sulfur dioxide and nitrogen dioxide make acids when dissolved in water.

Content

Scheme of work

C3.1 — The properties of oxygen

Background

Oxygen makes up about one-fifth of the Earth’s atmosphere. It is colourless, odourless and supports combustion, but does not itself burn.

Combustion reactions require oxygen (usually from the air), heat and a fuel source.

If insufficient oxygen is present for the total fuel, then the combusion will be incomplete and the products will include less oxidised species. For example, when methane burns in plentiful air or oxygen the following reaction takes place:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)

Complete combustion of methane

However, if there is restricted supply of air (or oxygen) then the methane cannot fully react and incomplete combustion products, carbon monoxide, carbon microparticulates, can form.

CH4(g) + 1½ O2(g) → CO(g) + 2H2O(l)

Incomplete combustion of methane

In incomplete combustion the carbon monoxide produced is toxic and potentially fatal. Every year there are cases of accidental poisoning and fatalities caused by heating systems that do not burn the fuel completely.

Case study

Incomplete combustion creates carbon microparticulates, PM10, that show up as dark smoke or fumes. This leads to air pollution and is a major problem in cities and population centres. The pollution is a leading cause of lung disease and cancer.

Preparation of oxygen

Oxygen may be prepared in the laboratory in several ways. Joseph Priestly first isolated oxygen by heating a sample of mercuric oxide, HgO, using the focussed heat of the sun.

Joseph Priestly's experiment

Now we know that there are far simpler means of making oxygen gas in the laboratory.

Activity — Preparing oxygen in the lab

Apparatus and chemicals

  • Conical flask with side arm and bung
  • Delivery tube and trough of water
  • Hydrogen peroxide solution (10–15%)
  • Manganese(IV) oxide catalyst (a spatula)
  • Glowing splint

Procedure

  • Set up the apparatus with the delivery tube leading into an inverted test tube over water.
  • Add a small amount of MnO2 catalyst to the flask.
  • Pour in hydrogen peroxide and quickly replace the bung.
  • Collect the gas over water, discarding the first tube (mostly air).
  • Test a fresh tube with a glowing splint — it relights in oxygen.

Safety

  • Wear goggles and gloves: H2O2 is an oxidiser and irritant.
  • Handle MnO2 powder carefully, avoid dust.
  • Keep all flames small and controlled.

Teacher demonstrations

Residue tests

Summary

Oxygen can be prepared in the lab by catalysed decomposition of hydrogen peroxide. A glowing splint relights in oxygen. Metals and non-metals burn differently: iron produces sparks, carbon forms CO2, sulfur produces choking SO2. The acid-base nature of the resulting oxide (acidic, neutral, insoluble) can be tested with indicators.

Check your understanding

  • Why is MnO2 described as a catalyst in the decomposition of hydrogen peroxide?
  • Which combustion demo produced the most vigorous reaction? Explain using the fire triangle.
  • How can you test whether an oxide is acidic or neutral?

C3.2 — Metal and non-metal oxides

Background

This lesson builds on C3.1, where oxygen was prepared and tested. Here, you will compare how oxides behave in water and with acids. Non-metal oxides such as carbon dioxide typically form acidic solutions in water, whereas many metal oxides are basic and react with acids to form salts and water. You will prepare CO2, choose an appropriate collection method, and show that copper(II) oxide behaves as a basic oxide by reacting with dilute sulfuric acid.

Activity — Comparing oxides: CO₂ preparation and CuO reaction

Apparatus and chemicals

  • Boiling tube(s), test tubes, delivery tube, rubber bung
  • Conical flask for gas generation; trough or beaker of water
  • Marble chips (CaCO3) or sodium hydrogencarbonate (NaHCO3)
  • Hydrochloric acid (dilute)
  • Limewater (Ca(OH)2(aq)) in a test tube
  • Copper(II) oxide powder (CuO)
  • Dilute sulfuric acid (≈ 0.5–1.0 mol dm−3)
  • Heat-proof mat, test-tube holder, small spatula
  • Source of gentle heat (if needed to speed CO2 generation)

Procedure

  • Prepare CO2: Place a few marble chips (or a spatula of NaHCO3) in a conical flask. Add dilute HCl and quickly fit the bung with delivery tube.
  • Choose collection method: Either collect over water in an inverted test tube, or use downward delivery into an upright dry test tube (CO2 is denser than air).
  • Test for CO2: Bubble the gas through limewater. Record the observation (turns milky due to CaCO3 formation). Briefly try a small flame near a CO2-filled tube to observe extinguishing.
  • Show a basic oxide: Place a small amount of CuO in a test tube. Add warm dilute H2SO4 in portions, swirling until a blue solution forms (CuSO4). Decant/filtrate if solid remains.
  • Write balanced equations for each reaction in your notes.

Safety

  • Wear goggles and gloves. Dilute acids are corrosive to eyes and skin.
  • Handle CuO powder carefully; avoid inhalation. Clean spills promptly.
  • Do not inhale CO2. Work in a well-ventilated area; keep flames small and controlled.
  • Glassware can become hot; use holders and heat-proof mats.

Key observations

Worked examples

  • CO2 preparation (using marble chips): CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
  • Limewater test: Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)
  • Basic oxide reaction: CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)

Choosing a gas collection method

Summary

Non-metal oxides like CO2 generally form acidic solutions in water and can be identified with the limewater test. Many metal oxides behave as bases; CuO reacts with dilute sulfuric acid to give copper(II) sulfate and water. CO2 is denser than air, so downward delivery is an efficient collection method.

Check your understanding

  • Explain why downward delivery is usually better than collection over water for CO2.
  • Classify each oxide and justify briefly: CO2, CuO, Al2O3.
  • Write the ionic reason why limewater turns milky when exposed to CO2.

C3.3 — Oxygen in the air

Background

This lesson links to C3.2 (acidic vs basic oxides). When a small candle burns in a sealed space, oxygen is consumed and carbon dioxide and water are produced. If carbon dioxide is absorbed by dilute sodium hydroxide, the volume change is dominated by oxygen consumption. Measuring the trapped liquid rise after cooling lets us estimate the percentage of oxygen in air.

Activity — Estimate the percentage of oxygen in air

Apparatus and chemicals

  • Wide trough or shallow tray with water
  • Short candle or tealight on a small floating platform (e.g., foil on a bit of cork)
  • Large gas jar or wide-mouth glass beaker (inverted over the candle)
  • Ruler or adhesive scale for external height readings
  • Dropper or small beaker with dilute sodium hydroxide solution (≈ 0.5–1.0 mol dm−3)
  • Marker pen or wax pencil
  • Timer

Procedure

  • Place the floating candle in the centre of the water tray. Add a few millilitres of dilute NaOH to the water so that any CO2 formed is absorbed when it dissolves into the water under the jar.
  • Light the candle. Invert the gas jar over the candle so its rim is below the water surface. Avoid trapping bubbles along the rim.
  • Observe: the flame will burn for a short time then go out. Wait at least 2–3 minutes for cooling to occur and for CO2 to be absorbed.
  • Mark the final water level on the outside of the jar. Also mark the initial water level (before covering) if not already done.
  • Measure the change in water height inside the jar and estimate the fraction of the jar’s internal height that has filled with water.
  • Calculate %O2 ≈ (volume fraction replaced by water) × 100%.

Safety

  • Wear goggles and gloves. Dilute NaOH is caustic; rinse spills with plenty of water.
  • Keep flames small and supervised; tie back hair; no loose sleeves.
  • Handle hot glass carefully; wait for cooling before dismantling.

Key chemistry

Notes on measurement

Worked example

  • A jar 16 cm tall shows a 3.2 cm water rise after cooling. Estimated %O2 ≈ (3.2 ÷ 16) × 100% = 20% (close to 21%).
  • Major uncertainties: reading the water line (±0.2 cm), incomplete CO2 absorption, small leaks at the rim, and temperature effects.

Summary

Burning a small candle in a closed space consumes oxygen. If CO2 is absorbed by NaOH and the system is allowed to cool, the water rise indicates how much of the original air has lost oxygen. A simple height-fraction estimate gives an oxygen percentage close to 21% when errors are minimised.

Check your understanding

  • Why must you wait for the gas to cool before reading the water level? What error appears if you do not?
  • State one design change that would reduce each of these: leaks at the rim; incomplete CO2 absorption; reading uncertainty.
  • Your class mean is 18%. Give two likely reasons for the low value and one improvement to approach 21%.

C3.4 — Fire and fuels

Background

This lesson extends C3.3 (oxygen content of air) to fuels and combustion stoichiometry.

Hydrocarbons such as methane burn in oxygen to release energy. The ratio of fuel to air is critical: too little oxygen means incomplete combustion; too much air means the mixture will not ignite well. Explosive limits define the concentration range in which a gas–air mixture can burn. Demonstrations with methane and student inquiry with hydrogen–oxygen bottle rockets illustrate these concepts.

Activity — Methane can demonstration (teacher only)

Apparatus and chemicals

  • Clean metal can with small hole in the top lid
  • Gas inlet tube reaching to the bottom of the can
  • Methane supply (laboratory natural gas)
  • Safety screen and hearing protection
  • Ignition source (taper, long match or spark igniter)

Procedure

  • Insert the gas tube into the can and allow methane to displace air from the bottom upward for a few seconds.
  • Ignite the small hole at the top: initially a small flame appears.
  • Turn off the gas and remove the inlet tube. As air is drawn into the can, the flame grows, changes colour, then finally produces a loud “whoomph” when the mixture reaches the critical ratio.

Safety

  • Teacher demonstration only. Use a clear safety screen between apparatus and class.
  • Do not overfill with gas. Ventilate the area well.
  • Wear hearing protection if available. Ensure no flammables are nearby.

Activity — Hydrogen/oxygen bottle rockets (student inquiry)

Apparatus and chemicals

  • Small plastic bottles (≈250 mL, wide-neck, no screw caps)
  • Hydrogen generator (e.g., Zn + HCl) placed well away from ignition zone
  • Oxygen source (prepared by catalyst decomposition of H2O2)
  • Water trough for collecting gases by displacement
  • Launcher tube or push-fit connector
  • Safety screen, tape measure, and clear floor “range” area
  • Ignition system controlled by teacher (spark, long taper)

Procedure

  • Fill bottles with varying H2:O2 ratios (e.g. 1:1, 2:1, 3:1) by collecting gases over water.
  • Quickly invert and fit to the launcher under the safety screen.
  • Teacher ignites. Observe launch distance and record ratio used.
  • Compare results across mixtures and plot distance vs ratio.

Safety

  • Only the teacher ignites. Students remain behind the safety line.
  • Keep hydrogen generator far from ignition zone. Use small bottles only.
  • Do not screw-cap or seal bottles — risk of explosion. Always launch away from people.

Key chemistry

Worked example

  • For methane: 1 mol CH4 needs 2 mol O2. Since air is 21% O2, moles of air needed = 2 ÷ 0.21 ≈ 9.5 mol.
  • For hydrogen: 2H2 + O2 → 2H2O. Optimum ratio = 2:1 by volume.

Summary

Combustion depends critically on the fuel-to-air ratio. Methane shows a characteristic ignition pattern in the methane can demonstration, while hydrogen–oxygen mixtures in bottle rockets highlight the importance of stoichiometry. The most energetic launches occur close to the stoichiometric 2:1 H2:O2 ratio.

Check your understanding

  • Explain why the methane flame changes before the “whoomph” in the methane can demo.
  • Why is 2:1 the optimum volumetric ratio for H2:O2 mixtures?
  • Give one systematic factor (other than ratio) that may affect rocket distance.

C3.5 — Measuring the energy of a fuel

Background

This lesson follows PS3.O4 on fuels and mixtures. Here, you will measure the energy released by burning different alcohol fuels using simple calorimetry. The aim is to connect combustion enthalpy (ΔHc) with experimental measurements, while recognising heat losses and uncertainty. Spirit burner experiments highlight the difference between theoretical and measured energy values and allow discussion of experimental design.

Activity — Calorimetry with alcohol fuels

Apparatus and chemicals

  • Spirit burners with ethanol and methanol fuels (with caps)
  • 100 mL graduated cylinder, beaker or metal calorimeter can (≈200 mL)
  • Digital balance (±0.01 g)
  • Thermometer (±0.5 °C)
  • Clamp stand with draught shield (optional: simple foil wind shield)
  • Heat-proof mat, matches/taper

Procedure

  • Measure 100 g (100 mL) of water into the calorimeter/beaker. Record its initial temperature.
  • Weigh the filled spirit burner with its lid. Record mass.
  • Place the burner under the beaker at a fixed distance. Light the wick and stir the water gently while heating.
  • Extinguish after a ~20 °C rise in temperature. Reweigh the burner with lid.
  • Repeat with the other fuel. Record all data carefully.

Safety

  • Alcohols are highly flammable. Cap burners immediately when not in use.
  • Do not refill hot burners. Allow to cool first.
  • Wear goggles; keep flames small and supervised. Fire blanket should be nearby.

Key chemistry

Worked example

  • Water mass = 100 g; ΔT = 18 °C → q = 100 × 4.18 × 18 ≈ 7.5 kJ
  • Fuel mass lost = 0.25 g methanol; n = 0.25 ÷ 32 ≈ 0.0078 mol
  • ΔHc ≈ −7.5 ÷ 0.0078 ≈ −960 kJ mol−1 (literature ~−715 kJ mol−1 — difference due to heat loss, incomplete combustion, evaporation)

Design and uncertainty

Summary

Simple calorimetry shows how much energy fuels release when burned. Measured ΔHc values are less exothermic than literature values due to heat loss and incomplete combustion. Careful experimental design reduces these errors. Comparing ethanol and methanol illustrates how molecular structure affects combustion energy.

Check your understanding

  • Explain why measured ΔHc values are always less negative than literature values.
  • If your thermometer has ±0.5 °C readability and ΔT = 18 °C, what is the % uncertainty from this source?
  • Suggest one design improvement that would most reduce heat loss, and explain why.

C3.6 — Sulfur dioxide and nitrogen oxides

Background

This lesson builds on PS3.O5 (energy from fuels). Burning non-metals often produces acidic oxides that dissolve in water to form acids. Sulfur forms sulfur dioxide (SO2) which dissolves to make sulfurous acid, and nitrogen from the air can form nitrogen oxides (NO, NO2) in hot flames, contributing to acid rain. You will prepare and test SO2, and discuss environmental impacts of acidic oxides.

Activity — Preparing and testing sulfur dioxide

Apparatus and chemicals

  • Small spatula of powdered sulfur
  • Combustion spoon or deflagrating spoon with lid
  • Gas jar with cover plate
  • Universal indicator solution and water
  • Heat-proof mat, tongs
  • Bunsen burner

Procedure

  • Place a few drops of water with universal indicator in the bottom of a gas jar.
  • Put a small amount of sulfur into the combustion spoon. Ignite strongly in a Bunsen flame until it burns with a blue flame.
  • Quickly lower the spoon into the gas jar and cover with the plate to trap the gas.
  • Swirl gently to dissolve SO2 in the indicator solution. Record colour change.
  • Write a balanced chemical equation for the combustion of sulfur.

Safety

  • SO2 is toxic and irritant to lungs — carry out in a fume cupboard or well-ventilated area.
  • Burn only small amounts of sulfur. Do not inhale fumes.
  • Wear goggles and keep flames small and controlled.

Key chemistry

Worked examples

  • SO2 in water: forms H2SO3, which donates H+ ions → acidic solution.
  • NO2 in rainwater: 4NO2(g) + 2H2O(l) + O2(g) → 4HNO3(aq)
  • Environmental link: Acid rain damages limestone, corrodes metals, and harms aquatic life.

Discussion

Summary

Sulfur dioxide is an acidic oxide that dissolves in water to form sulfurous acid. Nitrogen oxides form in high-temperature flames and contribute to acid rain. Acidic oxides are responsible for many environmental issues including corrosion and ecosystem damage.

Check your understanding

  • Write the balanced equation for sulfur burning in oxygen. What type of oxide is formed?
  • Explain why nitrogen oxides are formed more in car engines than in ordinary laboratory flames.
  • Suggest one environmental problem caused by acid rain and one possible solution.

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