Unit Three: Chemistry In Society

This unit explores how chemistry connects to the real world around you. You'll learn about industrial processes, environmental impact, and the calculations that help chemists make important decisions.

Understanding these concepts will help you see chemistry beyond the lab - from the fertilisers that grow our food to the materials in your smartphone.

Quick Tip: This unit combines practical maths skills with environmental awareness - perfect preparation for both exams and understanding current events!

Getting The Most from Reactants

Industrial processes aren't just about making products - they're about making money whilst protecting the environment. Companies must balance profit with sustainability, which creates fascinating challenges.

Key factors include the availability and cost of reactants, plus opportunities for recycling by-products. Smart factories use waste from one process as raw materials for another, reducing costs and environmental impact.

Energy requirements matter hugely since heating costs money. Exothermic reactions can actually provide energy for other processes, whilst endothermic ones need expensive heating.

Environmental considerations focus on minimising waste and avoiding toxic substances like carbon monoxide. Modern chemical design emphasises biodegradable products that won't harm ecosystems.

Real-World Connection: Think about how crisp packets are designed - they need to preserve food but also break down safely after disposal.

N5 Revision: Mass Calculations

Mole calculations are your gateway to understanding chemical quantities. The key formula connects mass, moles, and formula mass: n = m/g $where n = moles, m = mass, g = formula mass$.

Let's tackle a typical problem: calculating aluminium mass from aluminium oxide electrolysis. Using the equation 2Al₂O₃ → 4Al + 3O₂, you can see that 2 moles of oxide produce 4 moles of aluminium.

With 51g of aluminium oxide (0.5 mol), you'll get 1 mol of aluminium. Since aluminium's formula mass is 27g/mol, you'll produce 27g of aluminium metal.

Concentration calculations use C = n/V, linking moles to volume. Master these formulas and you'll breeze through stoichiometry problems.

Exam Strategy: Always write out the balanced equation first - it shows the mole ratios you need for calculations.

Molar Volume

Molar volume is brilliantly simple - one mole of any gas occupies roughly 24 litres at standard conditions. This makes gas calculations much easier than you'd expect.

For the nitrogen and hydrogen reaction $N₂ + 3H₂ → 2NH₃$, 400cm³ of nitrogen produces 800cm³ of ammonia. Just use the mole ratios from the balanced equation.

Gas volume calculations follow the same ratio patterns as mole calculations. In the butane combustion example $C₄H₁₀ + 6.5O₂ → 4CO₂ + 5H₂O$, 250cm³ of butane needs 1,625cm³ of oxygen.

The beauty of molar volume is its universality - whether you're dealing with hydrogen, oxygen, or carbon dioxide, one mole always occupies the same space.

Memory Trick: Remember "24 litres per mole" - it's your key to solving any gas volume problem quickly.

Complex Gas Calculations

Excess reactant problems require careful tracking of what's used and what remains. In the methane combustion $CH₄ + 2O₂ → CO₂ + 2H₂O$, 50cm³ methane uses 100cm³ oxygen, leaving 300cm³ oxygen unreacted.

The Haber process example shows nitrogen as the limiting reactant. With 10cm³ N₂ and 50cm³ H₂, you'll produce 20cm³ ammonia and have 20cm³ hydrogen left over.

Solid-gas reactions need extra attention to units. When magnesium reacts with hydrochloric acid, convert everything to consistent units before calculating the hydrogen gas volume.

Practice makes perfect with these multi-step problems. Break them down systematically and check your mole ratios carefully.

Problem-Solving Tip: Always identify which reactant runs out first - this determines how much product you can make.

Excess Calculations

Limiting reactants determine how much product forms, whilst excess reactants ensure complete conversion. It's like baking - if you're short of eggs, extra flour won't help you make more cakes.

For magnesium and hydrochloric acid $Mg + 2HCl → MgCl₂ + H₂$, calculate moles of each reactant separately. With 0.01 mol Mg but only 0.01 mol HCl, you need 0.02 mol HCl - so hydrochloric acid is limiting.

Economic considerations drive these calculations. Industries use cheap reactants in excess to ensure expensive ones react completely, maximising profit from costly materials.

The magnesium-nitric acid example shows 0.139 mol Mg with 0.15 mol HNO₃. Since you need 0.278 mol acid for complete reaction, magnesium is in excess this time.

Industrial Insight: Pharmaceutical companies often use expensive active ingredients as limiting reactants, with cheaper solvents in large excess.

Percentage Yield

Percentage yield measures reaction efficiency - how much product you actually get compared to theoretical maximum. It's calculated as: (actual yield ÷ theoretical yield) × 100%.

Esterification reactions never give 100% yield because they're reversible. The methanol-ethanoic acid reaction producing methyl ethanoate typically gives around 83% yield, which is pretty good for this type of reaction.

Real industrial processes rarely achieve perfect yields due to side reactions, incomplete conversion, and product losses during purification. Understanding these limitations helps chemists optimise conditions.

The ammonia synthesis example shows how scaling up affects calculations. With 790kg hydrogen producing 850kg ammonia instead of the theoretical 4,476kg, the yield drops to just 19%.

Industry Reality: Higher yields mean better profits and less waste - that's why chemical engineers spend years optimising reaction conditions.

Atom Economy

Atom economy measures how much of your starting materials end up in useful products. It's calculated as: (mass of desired product ÷ total mass of reactants) × 100%.

Making hydrogen from coal $C + 2H₂O → CO₂ + 2H₂$ has terrible atom economy - only 8.3%. Most of your reactants become carbon dioxide waste rather than useful hydrogen.

The Haber process $N₂ + 3H₂ ⇌ 2NH₃$ achieves perfect 100% atom economy because ammonia is the only product. This makes it environmentally and economically attractive.

Green chemistry prioritises high atom economy reactions. Modern chemical design focuses on processes where most reactant atoms end up in the final product rather than waste.

Environmental Impact: Low atom economy means more waste disposal costs and environmental damage - expensive problems for chemical companies.

Equilibrium Basics

Chemical equilibrium occurs in closed systems where forward and reverse reaction rates become equal. Concentrations stay constant, but reactions continue happening in both directions.

Le Chatelier's principle predicts equilibrium shifts: systems respond to changes by opposing them. Think of it like a seesaw - disturb the balance and it adjusts to restore stability.

Temperature changes favour endothermic reactions when heated. For N₂O₄ ⇌ 2NO₂ (endothermic forward), increasing temperature shifts right, producing more brown NO₂ gas.

The four main factors affecting equilibrium are temperature, concentration, pressure, and catalysts. Understanding these helps chemists control industrial processes like ammonia synthesis.

Visual Learning: Imagine equilibrium as a busy two-way street - equal traffic in both directions, but cars still moving constantly.

Equilibrium Factors

Concentration changes shift equilibrium to counteract the disturbance. Adding HCl to the bromine-water equilibrium $Br₂ + H₂O ⇌ 2H⁺ + Br⁻ + BrO⁻$ increases H⁺ concentration, so the equilibrium shifts left.

Adding NaOH removes H⁺ ions, forcing the equilibrium right to replace them. These acid-base effects are crucial in buffer systems and industrial processes.

Pressure changes affect gaseous equilibria. For 2SO₂ + O₂ ⇌ 2SO₃, increasing pressure favours the side with fewer gas molecules - the products (2 moles vs 3 moles).

Catalysts speed up both forward and reverse reactions equally, so they don't change equilibrium position. However, they help systems reach equilibrium faster, which matters enormously in industry.

Exam Success: Remember that catalysts affect rate, not position - a common source of exam mistakes!