Mission 1: When a cold pack goes cold, has the chemical inside lost energy or gained it?

System and surroundings

• The system is the reacting chemicals: reactants and products.
• The surroundings are everything else: water, container, thermometer, air.
• Energy moves between them, but the total is conserved.
• A thermometer in a solution reads the surroundings, not the chemicals.

Heat is not temperature

• Heat: energy transferred because of a temperature difference, in joules.
• Temperature: in kelvin, proportional to the average kinetic energy of the particles.
• A bath and a cup of water at 40 °C have the same temperature.
• The bath can transfer far more heat as it cools: it has many more particles.

Exothermic and endothermic

• Exothermic: energy passes from system to surroundings. Surroundings warm. ΔH is negative.
• Endothermic: energy passes from surroundings to system. Surroundings cool. ΔH is positive.
• ΔH describes the system. Always write its sign: −184 or +178 kJ mol⁻¹.
• Exothermic: combustion, neutralization. Endothermic: thermal decomposition of CaCO₃, photosynthesis.

Energy profiles

• Axes: potential energy (y) against reaction coordinate (x). Not time.
• Activation energy, Ea: from the reactant level up to the top of the barrier.
• ΔH: from the reactant level to the product level.
• Products lower → exothermic. Products higher → endothermic.
• Both kinds of reaction have an activation-energy barrier.

Stability decides the direction

• Lower potential energy means more stable.
• Exothermic: the products are more stable than the reactants.
• Endothermic: the products are less stable; the system must take in energy.

Mission 2: Where does half the energy go when you burn a fuel under a can of water?

The standard enthalpy change, ΔH⦵

• ΔH: the heat exchanged with the surroundings when the pressure is kept constant.
• ΔH⦵: measured at standard pressure, 100 kPa, with every substance in its standard state. Usually 298 K.
• Units: kJ mol⁻¹, for the amounts in the equation.
• States matter: forming H₂O(l) releases more energy than forming H₂O(g).

Q = mcΔT

• Q = mcΔT (data booklet). Q is in J.
• m: mass of the water or solution whose temperature changes, in g. Not the reactant's mass.
• c: specific heat capacity of water, 4.18 J g⁻¹ K⁻¹ (data booklet).
• ΔT: the temperature change. Same in K and °C; never add 273.
• For dilute solutions, take density 1.00 g cm⁻³ and c of water.

ΔH = −Q/n

• ΔH = −Q/n. n is the amount of the limiting reactant, in mol.
• Divide Q by 1000 to get kJ.
• The minus sign: the surroundings gain Q, so the system loses it.
• A temperature rise gives a negative ΔH; a fall gives a positive ΔH.

Finding the true ΔT

• Heat is exchanged with the air while the reaction is still going.
• So the highest (or lowest) reading falls short of the change with no heat exchange.
• Plot temperature against time. Extrapolate the line after the highest (or lowest) reading back to the time of mixing.
• ΔT = extrapolated value − starting temperature.

Worked example

• Excess zinc powder is added to 50.0 cm³ of 0.200 mol dm⁻³ CuSO₄(aq): Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s).
• From the graph: ΔT = 31.00 − 21.40 = 9.60 K.
• Q = 50.0 × 4.18 × 9.60 = 2006 J = 2.006 kJ.
• n(CuSO₄) = 0.0500 dm³ × 0.200 mol dm⁻³ = 0.0100 mol. Zinc is in excess, so CuSO₄ is limiting.
• ΔH = −2.006 ÷ 0.0100 = −201 kJ mol⁻¹.

Why measured values fall short

• Heat is lost to the air and apparatus, so measured ΔH is less exothermic than the accepted value.
• Burning a fuel under a can loses most: much of the heat warms the air and the can.
• Fuels may also burn incompletely, and conditions are not standard.
• Reduce losses with a lid, insulation, and the burner close to the can.

Mission 3: If breaking bonds takes in energy, how can burning a fuel give energy out?

Breaking takes in, forming gives out

• Breaking a bond is endothermic: energy is absorbed.
• Forming the same bond releases exactly the same energy: exothermic.
• A reaction is exothermic when the bonds formed are stronger in total than those broken.
• Bonds never release energy as they break.

Bond enthalpy

• Bond enthalpy: energy needed to break 1 mol of a bond, with all species gaseous.
• Always positive. Average values are in the data booklet.
• ΔH = Σ(bonds broken) − Σ(bonds formed).
• Draw the structures first, and count every bond broken and formed.

Worked example

• H₂(g) + Cl₂(g) → 2HCl(g).
• Data booklet: H–H 436, Cl–Cl 242, H–Cl 431 kJ mol⁻¹.
• Broken: 436 + 242 = 678 kJ. Formed: 2 × 431 = 862 kJ.
• ΔH = 678 − 862 = −184 kJ mol⁻¹.

Why it is only an estimate

• Bond enthalpies are averages over many compounds; a real bond depends on its molecule.
• They apply only to gases: energy for changes of state is left out.
• So a value from bond enthalpies can differ from an experimental one.

Mission 4: Carbon always burns partly to CO₂. So how do we know ΔH for making carbon monoxide?

Hess's law

• Hess's law: ΔH is the same whatever route links the same initial and final states.
• It follows from conservation of energy.
• So ΔH for a reaction that is hard to measure = the sum of ΔH along another route.

Three moves

• Reverse an equation: change the sign of ΔH.
• Multiply an equation: multiply ΔH by the same number.
• Add equations: add their ΔH values. Species on both sides cancel.
• Arrange the given equations so they add up to the target equation.

Worked example

• Target: C(s) + ½O₂(g) → CO(g).
• (1) C(s) + O₂(g) → CO₂(g), ΔH = −393.5 kJ mol⁻¹.
• (2) CO(g) + ½O₂(g) → CO₂(g), ΔH = −283.0 kJ mol⁻¹.
• Keep (1). Reverse (2): CO₂(g) → CO(g) + ½O₂(g), ΔH = +283.0.
• Add: CO₂ and ½O₂ cancel. ΔH = −393.5 + 283.0 = −110.5 kJ mol⁻¹.

Mission 5: Why can a heater burning in a closed tent kill, even with fuel to spare?

Combustion: reaction with oxygen

• Combustion: an exothermic reaction with oxygen when a substance is heated in it.
• Reactive metals form ionic oxides: 2Mg(s) + O₂(g) → 2MgO(s).
• Non-metals form covalent oxides: S(s) + O₂(g) → SO₂(g); C(s) + O₂(g) → CO₂(g).
• Write oxygen as O₂. Work out ionic oxide formulas from the ion charges: Al₂O₃.

Burning hydrocarbons and alcohols completely

• Complete combustion, in plenty of O₂: all C → CO₂, all H → H₂O.
• Balance C, then H, then O last.
• Propane: C₃H₈ + 5O₂ → 3CO₂ + 4H₂O.
• Alcohols already contain an O atom: subtract it before finding the O₂.
• Butan-1-ol: C₄H₉OH + 6O₂ → 4CO₂ + 5H₂O (13 O needed, 1 from the alcohol).

Incomplete combustion

• In limited O₂, some carbon forms carbon monoxide, CO, or carbon (soot), C.
• Hydrogen still forms H₂O.
• To CO: 2C₃H₈ + 7O₂ → 6CO + 8H₂O. To C: C₃H₈ + 2O₂ → 3C + 4H₂O.
• Less energy is released than in complete combustion.
• "Incomplete" means different products, not fuel left over.

Why CO and soot are harmful

• Carbon monoxide is colourless, odourless and toxic.
• It binds to haemoglobin more strongly than oxygen does, so blood carries less O₂.
• Soot particles damage the lungs and blacken surfaces; they give a smoky yellow flame.

Mission 6: Coal, oil, gas: for the same heat, which puts out least CO₂?

Fossil fuels

• Fossil fuels: coal, crude oil and natural gas. They formed from the remains of organisms over millions of years.
• Non-renewable: we use them far faster than they form.
• Burning them releases carbon locked away long ago, adding CO₂ to the atmosphere.

Energy per gram, CO₂ per kilojoule

• Specific energy: energy released per gram = energy per mol ÷ molar mass.
• Propane releases 2219 kJ mol⁻¹ (data booklet); M = 44.11 g mol⁻¹ → 50.3 kJ g⁻¹.
• C₃H₈ + 5O₂ → 3CO₂ + 4H₂O: 3 × 44.01 = 132.03 g CO₂ per 2219 kJ.
• That is 59.5 g of CO₂ per 1000 kJ.
• More hydrogen per carbon: more energy comes from forming water, so less CO₂ per kJ.

The greenhouse effect

• Sunlight, mostly visible, passes through the atmosphere and warms the Earth's surface.
• The warm surface emits infrared radiation.
• CO₂, H₂O and CH₄ molecules absorb infrared, then re-emit it in all directions, partly back down.
• More CO₂ → more infrared absorbed → the enhanced greenhouse effect: a warmer surface.

Coal, oil and natural gas compared

Biofuels

• Renewable: replenished naturally at least as fast as it is used.
• Biofuels come from recent biomass; plants fixed the carbon by photosynthesis within months or years.
• Photosynthesis: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g), using light energy.
• Examples: bioethanol from fermented sugars, biodiesel from plant oils, biogas from waste.

Biofuels: for and against

• For: renewable, and the CO₂ released was recently taken from the air.
• For: can be made locally, some from waste.
• Against: not carbon neutral. Farming, fertilizer, transport and processing burn fossil fuels.
• Against: crops compete with food for land and water; clearing land releases stored carbon.
• Against: often a lower energy per gram than the fossil fuel replaced.

Mission 7: Can a car run on hydrogen without burning it?

What a fuel cell does

• A fuel cell turns a fuel's chemical energy directly into electrical energy.
• Fuel and oxygen are fed in continuously, to separate electrodes. There is no flame.
• Anode (negative): the fuel is oxidized, losing electrons.
• Cathode (positive): oxygen is reduced, gaining electrons.
• Electrons flow through the external circuit from anode to cathode.

The hydrogen fuel cell, acidic electrolyte

• Anode: H₂ → 2H⁺ + 2e⁻.
• Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O.
• H⁺ ions move through the electrolyte to the cathode.
• Double the anode equation and add: 2H₂ + O₂ → 2H₂O. The only product is water.

Half-equations for carbon fuels

• The methanol fuel cell: overall CH₃OH + 1½O₂ → CO₂ + 2H₂O. It releases CO₂.
• Its cathode reaction is the same as in the hydrogen cell.
• At the anode, the fuel → CO₂. Balance: C atoms; O with H₂O; H with H⁺; charge with e⁻.
• Worked, ethanol: C₂H₅OH + 3H₂O → 2CO₂ + 12H⁺ + 12e⁻.
• Check: 4 O, 12 H and zero charge on each side.

The big ideas

• Energy moves between system and surroundings; the total is conserved. Heat is not temperature.
• Exothermic: surroundings warm, ΔH negative. Endothermic: surroundings cool, ΔH positive.
• Energy profile: potential energy against reaction coordinate. ΔH = products − reactants; Ea is the barrier.
• Q = mcΔT, with m the mass of water or solution; ΔH = −Q/n, with n for the limiting reactant.
• ΔH = Σ(broken) − Σ(formed) gives an estimate. Hess's law: reverse → change the sign; multiply; add.
• Complete combustion gives CO₂ + H₂O; limited O₂ gives CO or C. Compare fuels by kJ per gram and CO₂ per kJ.
• Biofuels are renewable but not carbon neutral. In a fuel cell, the fuel is oxidized at the anode.