Sciences · Topic 2.2
Equilibrium and Thermochemistry
This topic covers two of the most important concepts in chemistry: dynamic equilibrium and energy changes in reactions. Understanding how to predict how a system responds to change (Le Chatelier's Principle) and how to classify and calculate energy changes (thermochemistry) is essential for the eAssessment. Environmental applications — including the greenhouse effect and climate change — are also assessed.
What You'll Learn
- Define chemical equilibrium and explain what makes it "dynamic"
- Apply Le Chatelier's Principle to predict shifts in equilibrium
- Classify reactions as exothermic or endothermic using ΔH values
- Draw and interpret energy profile diagrams
- Explain the role of polymers and greenhouse gases in society and the environment
- Evaluate environmental impacts of industrial chemistry from multiple perspectives
eAssessment Focus
Criterion A: Explain equilibrium, Le Chatelier's Principle, and ΔH with correct terminology.
Criterion B: Predict equilibrium shifts from given conditions; evaluate hypotheses about reaction rates.
Criterion C: Interpret energy diagrams; process ΔH data to draw conclusions.
Criterion D: Evaluate the social and environmental implications of industrial processes (e.g., Haber process, plastic pollution, climate change).
Key Vocabulary
| Term | Definition |
|---|---|
| Chemical equilibrium | State where forward and reverse reaction rates are equal; concentrations are constant |
| Dynamic equilibrium | Equilibrium where both reactions continue — they are equal, not stopped |
| Le Chatelier's Principle | If a system at equilibrium is disturbed, it shifts to counteract the disturbance |
| Enthalpy (ΔH) | Heat energy change in a reaction at constant pressure; ΔH < 0 = exothermic; ΔH > 0 = endothermic |
| Exothermic | Reaction that releases heat energy to surroundings; products have less energy than reactants |
| Endothermic | Reaction that absorbs heat energy from surroundings; products have more energy than reactants |
| Activation energy (Ea) | The minimum energy required for a reaction to occur |
| Catalyst | A substance that lowers activation energy without being consumed; increases reaction rate |
| Polymer | A large molecule made of many repeated monomer units joined by covalent bonds |
| Greenhouse gas | A gas that absorbs infrared radiation, trapping heat in the atmosphere (CO₂, CH₄, N₂O, H₂O) |
Sciences · Topic 2.2
Chemical Equilibrium
When a reversible reaction reaches equilibrium, both forward and reverse reactions continue at equal rates. This is called dynamic equilibrium — concentrations are constant, but the system is not static.
What Is Dynamic Equilibrium?
Key Idea
At equilibrium: Rate of forward reaction = Rate of reverse reactionConcentrations of reactants and products remain constant (not necessarily equal)
Example — Haber Process: N₂(g) + 3H₂(g) ⇋ 2NH (g)
Both nitrogen and hydrogen continue to react to form ammonia (forward). Ammonia continues to decompose back to N₂ and H₂ (reverse). At equilibrium, these rates are equal and concentrations stop changing.
Both nitrogen and hydrogen continue to react to form ammonia (forward). Ammonia continues to decompose back to N₂ and H₂ (reverse). At equilibrium, these rates are equal and concentrations stop changing.
Factors That Affect Equilibrium Position
| Factor | Effect on equilibrium position | Effect on rate |
|---|---|---|
| Concentration change | Shifts toward side that uses the added substance | Changes both rates; new equilibrium established |
| Temperature change | Shifts toward endothermic side if heated; exothermic side if cooled | Both rates change; ratio changes |
| Pressure change (gases) | Shifts toward side with fewer moles of gas | Both rates change |
| Catalyst added | No shift in position | Both rates increase equally; equilibrium reached faster |
Critical Point: A catalyst does NOT change the equilibrium position or yield. It only allows equilibrium to be reached faster. This is a very common misconception in exams.
Sciences · Topic 2.2
Le Chatelier's Principle
Le Chatelier's Principle is one of the most powerful tools in chemistry for predicting how a system responds to change. Master the three types of disturbance: concentration, temperature, and pressure.
The Principle
Le Chatelier's Principle (1884)
If a system at equilibrium is subjected to a change in condition (concentration, temperature, or pressure), the system will shift in the direction that counteracts (partially offsets) the change.
Application to the Haber Process: N₂ + 3H₂ ⇋ 2NH (ΔH = −92 kJ/mol)
| Disturbance | System Response | Effect on NH yield |
|---|---|---|
| Increase [N₂] or [H₂] | Shift right (→) to use up reactant | Yield increases |
| Remove NH (product) | Shift right (→) to replace product | Yield increases (continuous removal used in industry) |
| Increase temperature | Shift left (←) — endothermic direction absorbs heat | Yield decreases |
| Decrease temperature | Shift right (→) — exothermic direction releases heat to compensate | Yield increases (but rate is slow) |
| Increase pressure | Shift right (→) — 4 mol gas → 2 mol gas (fewer molecules) | Yield increases |
| Add catalyst (Fe) | No shift; equilibrium reached faster | No change in yield |
Industrial Compromise: Haber Process Conditions
Theoretically, low temperature and high pressure maximise yield. In practice:
- Temperature ~450°C (compromise: low enough for reasonable yield; high enough for acceptable rate)
- Pressure ~200 atm (compromise: high enough to shift equilibrium; not so high that equipment costs are excessive)
- Iron catalyst to increase rate without affecting yield
Sciences · Topic 2.2
Thermochemistry
Thermochemistry studies heat energy changes in chemical reactions. Every reaction either releases or absorbs energy, characterised by its enthalpy change (ΔH).
Exothermic vs Endothermic
Exothermic (ΔH < 0)
Energy released to surroundings. Products have LESS energy than reactants. Surroundings get warmer. Examples: combustion, neutralisation, respiration.
Endothermic (ΔH > 0)
Energy absorbed from surroundings. Products have MORE energy than reactants. Surroundings get colder. Examples: photosynthesis, thermal decomposition, dissolving ammonium nitrate.
Energy Profile Diagrams
Exothermic profile: Reactants are higher on the energy axis than products. The curve rises to the activation energy (Ea) peak, then falls below the reactant level.
ΔH = Energy of products − Energy of reactants (negative for exothermic)
ΔH = Energy of products − Energy of reactants (negative for exothermic)
Endothermic profile: Reactants are lower than products. The curve rises to the Ea peak, then falls to a level above the reactants.
ΔH is positive for endothermic.
ΔH is positive for endothermic.
Effect of a catalyst: Lowers the activation energy (Ea) — the peak is lower. ΔH is unchanged (same reactant and product energy levels).
Bond Energies and ΔH Calculation
Formula
ΔH = ∑(Bond energies broken) − ∑(Bond energies formed)Breaking bonds = endothermic (requires energy)
Forming bonds = exothermic (releases energy)
Example: H₂ + Cl₂ → 2HCl
Bonds broken: 1 H–H (436 kJ) + 1 Cl–Cl (243 kJ) = 679 kJ absorbed
Bonds formed: 2 H–Cl (2 × 432 = 864 kJ) released
ΔH = 679 − 864 = −185 kJ/mol (exothermic)
Bonds broken: 1 H–H (436 kJ) + 1 Cl–Cl (243 kJ) = 679 kJ absorbed
Bonds formed: 2 H–Cl (2 × 432 = 864 kJ) released
ΔH = 679 − 864 = −185 kJ/mol (exothermic)
Greenhouse Effect and Climate Chemistry
| Gas | Source | Global Warming Potential (relative to CO₂) |
|---|---|---|
| CO₂ | Combustion, deforestation, respiration | 1 |
| CH₄ (methane) | Livestock, landfills, natural gas leaks | ~25 |
| N₂O | Agriculture (fertilisers), vehicle emissions | ~298 |
| HFCs | Refrigerants, aerosols | Hundreds to thousands |
eAssessment Connection: The chemistry of greenhouse gases connects Criterion A (knowing the science) to Criterion D (evaluating climate change from scientific, economic, social, and ethical perspectives). Be prepared to discuss both the chemistry and the broader implications.
Sciences · Topic 2.2
Organic Chemistry & Polymers
Organic chemistry is the chemistry of carbon compounds. Polymers — long-chain molecules made from monomers — have transformed modern life but pose serious environmental challenges.
Hydrocarbons and Functional Groups
| Group | Functional group | Example | Use |
|---|---|---|---|
| Alkanes | C–C single bonds | CH₄ (methane) | Fuel |
| Alkenes | C=C double bond | C₂H₄ (ethene) | Making polymers |
| Alcohols | –OH group | C₂H₅OH (ethanol) | Fuel, solvent, drinks |
| Carboxylic acids | –COOH group | CHCOOH (ethanoic acid) | Food preservatives, vinegar |
Addition Polymerisation
Reaction
Many alkene monomers join together (opening the C=C double bond) to form a long chain polymer.n(CH₂=CH₂) → —(CH₂–CH₂)n— [polyethene from ethene]
Polymers: Benefits and Environmental Costs
| Benefits | Environmental problems |
|---|---|
| Lightweight, durable, cheap | Non-biodegradable — persist for hundreds of years |
| Medical uses (sterile packaging, implants) | Microplastics in oceans and food chains |
| Food preservation (packaging) | Releases toxic gases when incinerated |
| Electrical insulation, waterproofing | Derived from non-renewable fossil fuels |
Criterion D — Evaluating Plastic Use: A full response evaluates the trade-off between the utility of plastics (hygiene, food safety, medical benefits) and their environmental cost (pollution, wildlife harm, climate impact from fossil fuel extraction), then proposes evidence-based solutions (bioplastics, circular economy, extended producer responsibility).
Sciences · Topic 2.2
Worked Examples
These examples model the depth and precision expected in the eAssessment. Practise writing full explanations, not just one-word answers.
EXPLAINFor N₂ + 3H₂ ⇋ 2NH (ΔH = −92 kJ/mol), explain the effect of increasing temperature on the equilibrium yield of ammonia.
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Full Solution
According to Le Chatelier's Principle, when a system at equilibrium is disturbed, it shifts to oppose the change.
Increasing temperature adds heat energy to the system. The system shifts in the direction that absorbs this extra heat — the endothermic direction. For this reaction, the forward reaction is exothermic (ΔH = −92 kJ/mol), so the reverse reaction (decomposition of NH) is endothermic.
Therefore, increasing temperature shifts equilibrium to the left, decreasing the yield of ammonia. This is why the industrial Haber process uses a compromise temperature (~450°C) — high enough for a useful reaction rate, but not so high that equilibrium yield is severely reduced.
Increasing temperature adds heat energy to the system. The system shifts in the direction that absorbs this extra heat — the endothermic direction. For this reaction, the forward reaction is exothermic (ΔH = −92 kJ/mol), so the reverse reaction (decomposition of NH) is endothermic.
Therefore, increasing temperature shifts equilibrium to the left, decreasing the yield of ammonia. This is why the industrial Haber process uses a compromise temperature (~450°C) — high enough for a useful reaction rate, but not so high that equilibrium yield is severely reduced.
CALCULATECalculate ΔH for the reaction CH₄ + 2O₂ → CO₂ + 2H₂O using bond energies: C–H = 413, O=O = 498, C=O = 799, O–H = 463 kJ/mol.
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Full Solution
CH₄ + 2O₂ → CO₂ + 2H₂O
Bonds broken (endothermic):
4 × C–H = 4 × 413 = 1652 kJ
2 × O=O = 2 × 498 = 996 kJ
Total energy in = 1652 + 996 = 2648 kJ
Bonds formed (exothermic):
2 × C=O (in CO₂) = 2 × 799 = 1598 kJ
4 × O–H (in 2H₂O) = 4 × 463 = 1852 kJ
Total energy out = 1598 + 1852 = 3450 kJ
ΔH = Energy broken − Energy formed = 2648 − 3450 = −802 kJ/mol
ΔH < 0 → exothermic (combustion of methane releases energy)
Bonds broken (endothermic):
4 × C–H = 4 × 413 = 1652 kJ
2 × O=O = 2 × 498 = 996 kJ
Total energy in = 1652 + 996 = 2648 kJ
Bonds formed (exothermic):
2 × C=O (in CO₂) = 2 × 799 = 1598 kJ
4 × O–H (in 2H₂O) = 4 × 463 = 1852 kJ
Total energy out = 1598 + 1852 = 3450 kJ
ΔH = Energy broken − Energy formed = 2648 − 3450 = −802 kJ/mol
ΔH < 0 → exothermic (combustion of methane releases energy)
PREDICTFor the equilibrium: 2SO₂(g) + O₂(g) ⇋ 2SO(g) (ΔH = −196 kJ/mol). Predict the effect on yield of SO if pressure is doubled.
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Full Solution
By Le Chatelier's Principle: increasing pressure causes the equilibrium to shift toward the side with fewer moles of gas.
Count moles of gas: Left side = 2 + 1 = 3 moles. Right side = 2 moles.
Increasing pressure shifts equilibrium to the right (toward 2 moles of gas), which means more SO is produced. The yield of SO increases.
Count moles of gas: Left side = 2 + 1 = 3 moles. Right side = 2 moles.
Increasing pressure shifts equilibrium to the right (toward 2 moles of gas), which means more SO is produced. The yield of SO increases.
EVALUATEA student says a catalyst increases the yield of a reaction at equilibrium. Is this correct? Explain.
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Full Solution
The student is incorrect. A catalyst lowers the activation energy for both the forward and reverse reactions equally. This means both reactions speed up by the same factor, so the equilibrium position (and therefore yield) is unchanged.
What the catalyst does is allow the system to reach equilibrium faster. In industrial processes (like the Haber process), this is economically valuable — you reach the same yield in less time — but the equilibrium yield itself remains the same.
Only changes in temperature, concentration, or pressure (for gases) can shift the equilibrium position and change yield.
What the catalyst does is allow the system to reach equilibrium faster. In industrial processes (like the Haber process), this is economically valuable — you reach the same yield in less time — but the equilibrium yield itself remains the same.
Only changes in temperature, concentration, or pressure (for gases) can shift the equilibrium position and change yield.
ENVIRONMENTALExplain how CO₂ causes global warming at a molecular level.
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Full Solution
The Sun emits mainly shortwave radiation (visible light and UV), which passes through the atmosphere and warms Earth's surface.
Earth then re-emits energy as infrared (longwave) radiation. CO₂ molecules absorb this infrared radiation because the C=O bond vibrations match the frequency of infrared waves. The molecule then re-emits the energy in all directions, including back toward Earth's surface.
This trapping of heat is the enhanced greenhouse effect. Higher CO₂ concentrations (from burning fossil fuels) absorb more infrared radiation, reducing the amount escaping to space and warming the atmosphere. This is the mechanism of anthropogenic climate change.
Earth then re-emits energy as infrared (longwave) radiation. CO₂ molecules absorb this infrared radiation because the C=O bond vibrations match the frequency of infrared waves. The molecule then re-emits the energy in all directions, including back toward Earth's surface.
This trapping of heat is the enhanced greenhouse effect. Higher CO₂ concentrations (from burning fossil fuels) absorb more infrared radiation, reducing the amount escaping to space and warming the atmosphere. This is the mechanism of anthropogenic climate change.
EVALUATEEvaluate the environmental impact of plastic packaging from multiple perspectives.
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Full Solution
Benefits of plastic packaging: Extends food shelf life, reducing food waste (which has a higher carbon footprint than plastic production in some cases). Essential for medical supply hygiene. Lightweight → lower transport fuel use.
Environmental costs: Most plastics are non-biodegradable and persist for 400–1000 years. Marine plastic pollution kills seabirds, turtles, and fish. Microplastics have been found in drinking water, human blood, and placentas — health impacts are still under investigation. Plastic production from fossil fuels contributes to greenhouse gas emissions.
Social/economic: Plastic reduction policies can increase food prices and costs for small businesses. Recycling infrastructure is unevenly distributed globally — wealthy nations export waste to less-regulated countries.
Conclusion: A systemic transition to a circular economy (reuse, compostable materials, producer responsibility) is needed, but must be designed equitably to avoid placing the burden on low-income communities or developing nations.
Environmental costs: Most plastics are non-biodegradable and persist for 400–1000 years. Marine plastic pollution kills seabirds, turtles, and fish. Microplastics have been found in drinking water, human blood, and placentas — health impacts are still under investigation. Plastic production from fossil fuels contributes to greenhouse gas emissions.
Social/economic: Plastic reduction policies can increase food prices and costs for small businesses. Recycling infrastructure is unevenly distributed globally — wealthy nations export waste to less-regulated countries.
Conclusion: A systemic transition to a circular economy (reuse, compostable materials, producer responsibility) is needed, but must be designed equitably to avoid placing the burden on low-income communities or developing nations.
INTERPRETAn energy profile diagram shows Ea = 150 kJ/mol with a catalyst and 250 kJ/mol without. Describe what this tells you about the catalyst's effect on the reaction rate and yield.
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Full Solution
The catalyst lowers the activation energy from 250 kJ/mol to 150 kJ/mol — a reduction of 100 kJ/mol. This means a much larger proportion of reactant molecules have sufficient energy to overcome the energy barrier, so the reaction rate increases significantly.
However, the energy levels of reactants and products are unchanged, so ΔH remains the same. The equilibrium position is therefore unchanged — the catalyst has no effect on the final equilibrium yield, only on how quickly that equilibrium is reached.
However, the energy levels of reactants and products are unchanged, so ΔH remains the same. The equilibrium position is therefore unchanged — the catalyst has no effect on the final equilibrium yield, only on how quickly that equilibrium is reached.
Sciences · Topic 2.2
Practice Q&A
Attempt each question before revealing the model answer. For Le Chatelier's questions, always state the principle before applying it.
DEFINEWhat makes equilibrium "dynamic"?
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Model Answer
Equilibrium is "dynamic" because both the forward and reverse reactions continue to occur simultaneously at equal rates. It is not a static system where reactions have stopped — reactants continue forming products and products continue reforming reactants. The concentrations remain constant, but the reactions do not stop.
PREDICTFor A(g) + 2B(g) ⇋ 3C(g) (ΔH = +60 kJ/mol), what happens to yield of C if temperature is increased?
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Model Answer
By Le Chatelier's Principle: the system shifts to absorb the added heat energy → in the endothermic direction. ΔH = +60 kJ/mol means the forward reaction is endothermic. Therefore the equilibrium shifts right, increasing the yield of C.
CLASSIFYIdentify as exothermic or endothermic: (a) ΔH = −285 kJ/mol; (b) ΔH = +178 kJ/mol; (c) combustion of wood; (d) melting ice.
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Model Answer
(a) ΔH = −285 kJ/mol → exothermic (negative ΔH)
(b) ΔH = +178 kJ/mol → endothermic (positive ΔH)
(c) Combustion of wood → exothermic (heat and light released)
(d) Melting ice → endothermic (heat energy absorbed to break bonds in ice crystal)
(b) ΔH = +178 kJ/mol → endothermic (positive ΔH)
(c) Combustion of wood → exothermic (heat and light released)
(d) Melting ice → endothermic (heat energy absorbed to break bonds in ice crystal)
SUGGESTAn industrial chemist wants to maximise yield of SO from 2SO₂ + O₂ ⇋ 2SO (ΔH = −196 kJ/mol). What conditions should they use?
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Model Answer
To maximise yield:
1. Low temperature (exothermic reaction: lower T shifts equilibrium right) — but in practice a moderate temperature (~450°C) is used for acceptable rate
2. High pressure (right side has fewer gas moles: 2 vs 3) — shifts right
3. Remove SO as it forms (Le Chatelier: shifts right to replace product)
4. Excess O₂ (cheap and shifts equilibrium right)
5. Vanadium pentoxide catalyst (V₂O₅) to reach equilibrium faster without changing yield
1. Low temperature (exothermic reaction: lower T shifts equilibrium right) — but in practice a moderate temperature (~450°C) is used for acceptable rate
2. High pressure (right side has fewer gas moles: 2 vs 3) — shifts right
3. Remove SO as it forms (Le Chatelier: shifts right to replace product)
4. Excess O₂ (cheap and shifts equilibrium right)
5. Vanadium pentoxide catalyst (V₂O₅) to reach equilibrium faster without changing yield
CALCULATECalculate ΔH for H₂ + Br₂ → 2HBr. Bond energies: H–H = 436, Br–Br = 193, H–Br = 366 kJ/mol.
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Model Answer
Bonds broken: H–H (436) + Br–Br (193) = 629 kJ
Bonds formed: 2 × H–Br = 2 × 366 = 732 kJ
ΔH = 629 − 732 = −103 kJ/mol (exothermic)
Bonds formed: 2 × H–Br = 2 × 366 = 732 kJ
ΔH = 629 − 732 = −103 kJ/mol (exothermic)
EXPLAINWhy does a catalyst not change the equilibrium position of a reversible reaction?
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Model Answer
A catalyst works by providing an alternative reaction pathway with a lower activation energy. Crucially, it lowers the activation energy for both the forward and reverse reactions by the same amount. This means both reactions speed up equally, so the ratio of forward to reverse rate (and therefore the equilibrium constant) remains unchanged. The equilibrium position — and therefore the yield — does not change.
REAL-WORLDMethane is a greenhouse gas with a global warming potential 25 times that of CO₂. Explain why reducing methane emissions is an effective strategy against climate change.
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Model Answer
Methane's high global warming potential (GWP = 25) means each molecule traps significantly more heat than CO₂ over 100 years. Methane is also shorter-lived in the atmosphere (~12 years vs CO₂'s centuries), so reducing methane emissions produces faster climate benefits — concentrations fall more quickly once emissions stop.
Major sources include livestock (enteric fermentation), rice paddies, landfills, and natural gas leaks — many of which can be reduced through improved agricultural practices, landfill capture systems, and better pipeline maintenance. Given the urgency of climate action, targeting high-GWP gases like methane can deliver near-term temperature stabilisation while longer-term CO₂ reduction strategies are implemented.
Major sources include livestock (enteric fermentation), rice paddies, landfills, and natural gas leaks — many of which can be reduced through improved agricultural practices, landfill capture systems, and better pipeline maintenance. Given the urgency of climate action, targeting high-GWP gases like methane can deliver near-term temperature stabilisation while longer-term CO₂ reduction strategies are implemented.
DRAWDescribe what an energy profile diagram for an exothermic reaction with a catalyst would look like compared to without.
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Model Answer
Both diagrams show:
• Reactants at a higher energy level than products (exothermic: products lower)
• A "hump" (activation energy peak) between reactants and products
• The same reactant and product energy levels (same ΔH)
The difference: with a catalyst, the hump (Ea) is lower. The reactant and product levels are identical — only the peak height changes. This shows the catalyst lowers activation energy, increasing reaction rate, without affecting ΔH or equilibrium yield.
• Reactants at a higher energy level than products (exothermic: products lower)
• A "hump" (activation energy peak) between reactants and products
• The same reactant and product energy levels (same ΔH)
The difference: with a catalyst, the hump (Ea) is lower. The reactant and product levels are identical — only the peak height changes. This shows the catalyst lowers activation energy, increasing reaction rate, without affecting ΔH or equilibrium yield.
Sciences · Review
Flashcard Review
Tap each card to reveal the answer. Try to answer from memory first.
State Le Chatelier's Principle.
If a system at equilibrium is subjected to a change, it will shift in the direction that counteracts (partially offsets) the change.
Tap to reveal
What does ΔH < 0 mean?
The reaction is exothermic — energy is released to the surroundings. Products have less energy than reactants.
Tap to reveal
What does ΔH > 0 mean?
The reaction is endothermic — energy is absorbed from the surroundings. Products have more energy than reactants.
Tap to reveal
Does a catalyst change the equilibrium yield?
No. A catalyst lowers activation energy equally for both forward and reverse reactions, so equilibrium position is unchanged. It only speeds up reaching equilibrium.
Tap to reveal
For N₂ + 3H₂ ⇋ 2NH (ΔH = −92 kJ/mol): what happens to yield if temperature increases?
Yield decreases. The reaction is exothermic (forward) so increasing temperature shifts equilibrium left (endothermic direction) to absorb heat.
Tap to reveal
What is activation energy?
The minimum energy required for a reaction to occur — the energy barrier that must be overcome by reactant molecules.
Tap to reveal
What is dynamic equilibrium?
A state where the rate of the forward reaction equals the rate of the reverse reaction, so concentrations remain constant — but both reactions are still occurring.
Tap to reveal
ΔH formula using bond energies?
ΔH = ∑(Bond energies broken) − ∑(Bond energies formed). Breaking bonds requires energy; forming bonds releases energy.
Tap to reveal
What are greenhouse gases and why do they warm Earth?
Gases (CO₂, CH₄, N₂O) that absorb infrared radiation emitted by Earth's surface and re-emit it in all directions, trapping heat in the atmosphere.
Tap to reveal
What effect does increasing pressure have on N₂ + 3H₂ ⇋ 2NH?
Equilibrium shifts right (toward fewer gas moles: 4 mol gas → 2 mol gas), increasing NH yield.
Tap to reveal
What is a polymer?
A large molecule made of many repeating monomer units joined by covalent bonds. Example: polyethene from ethene monomers by addition polymerisation.
Tap to reveal
Why are most plastics an environmental problem?
They are non-biodegradable (persist for hundreds of years), create microplastic pollution in ecosystems, and are derived from non-renewable fossil fuels.
Tap to reveal
State the industrial conditions for the Haber process and the reason for each.
Temperature ~450°C (compromise: yield vs rate); pressure ~200 atm (high pressure favours products); iron catalyst (speeds up equilibrium attainment).
Tap to reveal
What happens to equilibrium if a product is removed?
The system shifts in the forward direction (right) to replace the removed product, as Le Chatelier's Principle predicts.
Tap to reveal
What is the global warming potential of methane compared to CO₂?
Methane has a GWP of approximately 25 over 100 years — meaning each molecule traps 25 times more heat than CO₂.
Tap to reveal
Sciences · Practice Test
Practice Test — 20 Questions
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