Welcome to Gas Chemistry!
Welcome to your study notes for Gas Chemistry! Gases are all around us — from the air we breathe to the bubbles in our fizzy drinks. In this topic, we will explore the gases that make up our atmosphere, how to prepare and test different gases in the laboratory, their physical and chemical properties, and their everyday uses.
Don't worry if this seems like a lot of information at first! We will break each gas down step-by-step using clear patterns, memorable tests, and simple tips to help you master your CCEA GCSE Chemistry exam.
1. The Earth's Atmosphere
The Earth is surrounded by a layer of gases called the atmosphere. The air we breathe is not a single substance, but a mixture of different gases in relatively constant proportions.
Approximate Composition of Clean, Dry Air:
• Nitrogen (\( \text{N}_2 \)): approximately \( 78\% \)
• Oxygen (\( \text{O}_2 \)): approximately \( 21\% \)
• Argon (\( \text{Ar} \)): approximately \( 0.9\% \)
• Carbon dioxide (\( \text{CO}_2 \)): approximately \( 0.04\% \)
• Other noble gases & water vapour (\( \text{H}_2\text{O} \)): variable trace amounts
Memory Trick: Remember the order from most to least abundant using "N-O-A-C" (Nitrogen, Oxygen, Argon, Carbon dioxide) — "No One Argues Chemistry!"
Key Takeaway: Nitrogen makes up almost four-fifths (\( \approx 78\% \)) of the air, while oxygen makes up about one-fifth (\( \approx 21\% \)).
2. Nitrogen Gas (\( \text{N}_2 \))
Physical Properties of Nitrogen:
• Colourless and odourless gas
• Insoluble (or only very slightly soluble) in water
• Density is similar to air
Why is Nitrogen so Unreactive?
Nitrogen exists as a diatomic molecule (\( \text{N}_2 \)). The two nitrogen atoms are joined by a very strong covalent triple bond (\( \text{N} \equiv \text{N} \)). A huge amount of energy is required to break this triple bond before nitrogen can react. Therefore, nitrogen is very unreactive (inert) at room temperature.
Uses of Nitrogen:
• Food Packaging: Nitrogen is pumped into crisp packets and food containers to provide an inert atmosphere. This displaces oxygen and prevents the food from going stale or oxidising.
• Liquid Nitrogen (Coolant): Because of its extremely low boiling point (\( -196\text{ }^\circ\text{C} \)), liquid nitrogen is used to freeze biological samples, blood, and cells, and in cryosurgery to remove warts.
Quick Review: Nitrogen is unreactive because of its strong triple covalent bond. Its main uses are in food packaging (inert atmosphere) and as a coolant (liquid nitrogen).
3. Ammonia Gas (\( \text{NH}_3 \))
Ammonia is a compound made of nitrogen and hydrogen. It is one of the few common alkaline gases you will encounter.
Physical Properties of Ammonia:
• Colourless gas
• Characteristic sharp, choking, pungent smell
• Less dense than air
• Extremely soluble in water, forming an alkaline solution of ammonium hydroxide (\( \text{NH}_4\text{OH} \))
Chemical Tests for Ammonia:
Test 1: Damp Red Litmus Paper
Hold a piece of damp red litmus paper near the gas. Ammonia is alkaline, so it turns the damp red litmus paper blue.
Test 2: Reaction with Concentrated Hydrochloric Acid (\( \text{HCl} \))
Dip a glass rod into concentrated hydrochloric acid and hold it near the ammonia gas.
Observation: Dense white fumes (smoke) are formed.
Word Equation: Ammonia \( + \) Hydrogen chloride \( \rightarrow \) Ammonium chloride
Symbol Equation: \( \text{NH}_3\text{(g)} + \text{HCl(g)} \rightarrow \text{NH}_4\text{Cl(s)} \)
Uses of Ammonia:
• The primary use of ammonia is in the manufacture of fertilisers (such as ammonium nitrate) to provide essential nitrogen to crops for healthy plant growth.
Key Takeaway: Ammonia is a pungent, alkaline gas that turns damp red litmus paper blue and produces dense white fumes with concentrated \( \text{HCl} \).
4. Hydrogen Gas (\( \text{H}_2 \))
Laboratory Preparation of Hydrogen:
Hydrogen is prepared in the lab by reacting a metal (such as zinc or magnesium) with a dilute acid (such as dilute hydrochloric acid or dilute sulfuric acid).
Word Equation: Zinc \( + \) Hydrochloric acid \( \rightarrow \) Zinc chloride \( + \) Hydrogen
Symbol Equation: \( \text{Zn(s)} + 2\text{HCl(aq)} \rightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\text{(g)} \)
Physical Properties of Hydrogen:
• Colourless and odourless gas
• Least dense gas (much less dense than air)
• Insoluble in water
Test for Hydrogen Gas:
Apply a lit splint to the gas. Hydrogen burns rapidly with a characteristic "squeaky pop".
Equation: \( 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)} \)
Uses of Hydrogen:
• Meteorological Balloons (Weather Balloons): Hydrogen is used to lift weather balloons high into the atmosphere because it is less dense than air.
• Clean Fuel / Rocket Fuel: When hydrogen burns in oxygen, the only product is water (\( \text{H}_2\text{O} \)). It produces no polluting carbon emissions or greenhouse gases.
Did You Know? Although helium is often used for party balloons because it is non-flammable, hydrogen has the lowest density of any element in the universe!
5. Oxygen Gas (\( \text{O}_2 \))
Laboratory Preparation of Oxygen:
Oxygen is prepared by the catalytic decomposition of hydrogen peroxide (\( \text{H}_2\text{O}_2 \)).
Hydrogen peroxide naturally breaks down very slowly into water and oxygen, but adding a black powder called manganese(IV) oxide (\( \text{MnO}_2 \)) acts as a catalyst to speed up the reaction rapidly.
Word Equation: Hydrogen peroxide \( \xrightarrow{\text{MnO}_2} \) Water \( + \) Oxygen
Symbol Equation: \( 2\text{H}_2\text{O}_2\text{(aq)} \xrightarrow{\text{MnO}_2} 2\text{H}_2\text{O(l)} + \text{O}_2\text{(g)} \)
Physical Properties of Oxygen:
• Colourless and odourless gas
• Slightly heavier/denser than air
• Slightly soluble in water (crucial for aquatic animals to survive!)
Test for Oxygen Gas:
Place a glowing splint into a tube of the gas. Oxygen will relight the glowing splint.
Combustion Reactions with Oxygen:
Oxygen supports combustion. When elements burn in oxygen, they form oxides:
1. Carbon (Non-metal):
• Observation: Burns with an orange flame; produces a colourless gas.
• Equation: \( \text{C(s)} + \text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)} \)
• Nature of Oxide: Acidic oxide (dissolves in water to form carbonic acid, turning universal indicator red/orange).
2. Sulfur (Non-metal):
• Observation: Melts to a red liquid and burns with a brilliant blue flame, producing a colourless, pungent gas (sulfur dioxide).
• Equation: \( \text{S(s)} + \text{O}_2\text{(g)} \rightarrow \text{SO}_2\text{(g)} \)
• Nature of Oxide: Acidic oxide (dissolves in water to form sulfurous acid, turning universal indicator red).
3. Magnesium (Metal):
• Observation: Burns with a bright white light to form a white ash / powder (magnesium oxide).
• Equation: \( 2\text{Mg(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MgO(s)} \)
• Nature of Oxide: Basic oxide (reacts with acids; slightly soluble in water to form an alkaline solution, turning universal indicator blue/purple).
4. Iron (Metal):
• Observation: Glows red-hot and produces orange sparks, leaving a black solid (triiron tetroxide).
• Equation: \( 3\text{Fe(s)} + 2\text{O}_2\text{(g)} \rightarrow \text{Fe}_3\text{O}_4\text{(s)} \)
• Nature of Oxide: Basic oxide.
5. Copper (Metal):
• Observation: Does not burn with a flame; the shiny reddish-brown metal surface turns black as copper(II) oxide forms.
• Equation: \( 2\text{Cu(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{CuO(s)} \)
• Nature of Oxide: Basic oxide.
Rule of Thumb for Oxides:
• Metal oxides are typically basic (neutralise acids).
• Non-metal oxides are typically acidic (form acidic solutions in water).
Uses of Oxygen:
• Medicine: Oxygen masks and cylinders in hospitals for patients with respiratory difficulties.
• Welding: Oxy-acetylene torches produce an extremely hot flame used to cut and weld steel.
6. Carbon Dioxide Gas (\( \text{CO}_2 \))
Laboratory Preparation of Carbon Dioxide:
Carbon dioxide is prepared in the lab by reacting calcium carbonate (marble chips) with dilute hydrochloric acid.
Word Equation: Calcium carbonate \( + \) Hydrochloric acid \( \rightarrow \) Calcium chloride \( + \) Water \( + \) Carbon dioxide
Symbol Equation: \( \text{CaCO}_3\text{(s)} + 2\text{HCl(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)} \)
Physical Properties of Carbon Dioxide:
• Colourless and odourless gas
• Denser than air (can be poured like a liquid!)
• Slightly soluble in water
Chemical Tests for Carbon Dioxide:
The Limewater Test:
Bubble the gas through limewater (a colourless solution of calcium hydroxide, \( \text{Ca(OH)}_2 \)).
• Initial Observation: Limewater turns milky / cloudy white.
Reason: An insoluble white precipitate of calcium carbonate (\( \text{CaCO}_3 \)) forms.
Equation: \( \text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)} \)
• Prolonged Bubbling: If you continue bubbling \( \text{CO}_2 \) through the cloudy mixture, the precipitate dissolves and the solution turns colourless again.
Reason: Insoluble calcium carbonate reacts with water and excess \( \text{CO}_2 \) to form soluble calcium hydrogencarbonate.
Equation: \( \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)} \rightarrow \text{Ca(HCO}_3)_2\text{(aq)} \)
Reaction of Carbon Dioxide with Water:
Carbon dioxide dissolves in water to form a weakly acidic solution called carbonic acid (\( \text{H}_2\text{CO}_3 \)).
Equation: \( \text{CO}_2\text{(g)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{H}_2\text{CO}_3\text{(aq)} \)
This gives carbonated drinks their slightly tangy, acidic taste.
Uses of Carbon Dioxide:
• Fire Extinguishers: Carbon dioxide is denser than air and does not support combustion. It blankets the fire, cutting off the oxygen supply.
• Carbonated Drinks: High pressure dissolves \( \text{CO}_2 \) in beverages to create fizz.
Key Takeaway: Carbon dioxide is a dense, acidic gas that turns limewater cloudy, but turns it colourless again upon prolonged bubbling.
7. Summary Table of Gas Tests
Here is a quick-reference summary of the tests you must know for your exam:
• Hydrogen (\( \text{H}_2 \)): Apply a lit splint \( \rightarrow \) burns with a squeaky pop.
• Oxygen (\( \text{O}_2 \)): Insert a glowing splint \( \rightarrow \) relights the glowing splint.
• Carbon Dioxide (\( \text{CO}_2 \)): Bubble through limewater \( \rightarrow \) turns cloudy / milky (turns colourless on prolonged bubbling).
• Ammonia (\( \text{NH}_3 \)): Hold damp red litmus paper \( \rightarrow \) turns blue; hold glass rod dipped in conc. \( \text{HCl} \) \( \rightarrow \) gives dense white fumes of \( \text{NH}_4\text{Cl} \).
8. Common Exam Mistakes to Avoid!
• Confusing Splints: Make sure you specify a glowing splint for oxygen (which relights) and a lit splint for hydrogen (which pops). Do not just write "splint"!
• Limewater Changes: Remember both stages of the limewater test: it turns cloudy first, and with prolonged bubbling it becomes colourless.
• Nitrogen's Triple Bond: If asked why nitrogen is unreactive, always mention the high energy required to break the triple covalent bond.
• Litmus Paper: Remember that litmus paper must be damp so that the gas can dissolve and show its acidic or alkaline nature.