Flashcards · IB Diploma Programme Chemistry · Topic 4

Reactivity 1 – What drives chemical reactions? flashcards

27 cards on Topic 4, subtopics 1.1 to 1.4. Pick a subtopic or study them all.

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1.1 Measuring enthalpy changes

Heat vs temperature: Energy on the move (J) vs a measure of mean particle kinetic energy (K). Heat depends on the amount of substance; temperature does not.
The thermometer rises during a reaction in solution. Exothermic or endothermic, and the sign of ΔH?: The thermometer reads the surroundings. They gained energy, so the system lost it: exothermic, ΔH negative.
Standard enthalpy change, ΔH⦵: ΔH under standard conditions (100 kPa; values usually for 298 K) with every substance in its standard state, in kJ mol⁻¹.
0.0150 mol of limiting reactant raises 40.0 g of solution by 4.3 K. ΔH?: Q = mcΔT = 40.0 × 4.18 × 4.3 = 719 J = 0.719 kJ. ΔH = −Q ÷ n = −0.719 ÷ 0.0150 = −47.9 kJ mol⁻¹.
Why is an experimental ΔH usually less negative than the data value?: Energy is lost to the air and bench; energy warms the cup or can; combustion is incomplete; fuel evaporates; products are not in standard states.
ΔH vs Ea on an energy profile: Gap between reactant and product levels vs gap from reactants up to the peak.

1.2 Energy cycles in reactions

Bond enthalpy: The energy absorbed when one mole of a particular covalent bond is broken in the gas phase; always positive. Average bond enthalpies are means over a range of compounds.
CH₂=CHCH₃ + Cl₂ → CH₂ClCHClCH₃. C=C 614, Cl–Cl 242, C–C 346, C–Cl 324 kJ mol⁻¹. ΔH?: Broken: 614 + 242 = 856. Formed: 346 + 2 × 324 = 994. ΔH = Σ(broken) − Σ(formed) = 856 − 994 = −138 kJ mol⁻¹.
Why is a ΔH from bond enthalpies only an estimate?: The values are averages over many molecules, and they apply to gases, so intermolecular forces in liquids or solids are ignored.
Hess's law: ΔH for a reaction is the same whatever pathway links the initial and final states (conservation of energy). Reverse an equation: change the sign of ΔH; scale it: scale ΔH.
Standard enthalpy of formation, and its value for an element: ΔH for making one mole of a substance from its elements, every species in its standard state. It is zero for an element in its standard state. (HL only)
Calculating ΔH⦵ from formation data vs combustion data: Formation: Σ ΔHf⦵(products) − Σ ΔHf⦵(reactants). Combustion: Σ ΔHc⦵(reactants) − Σ ΔHc⦵(products). (HL only)
Born–Haber: Li(s) → Li(g) +159; IE₁ +520; ½I₂ → I(g) +107; EA −295; lattice enthalpy +759 kJ mol⁻¹. ΔHf⦵ of LiI?: Elements to gaseous ions: 159 + 520 + 107 − 295 = +491. ΔHf⦵ = 491 − 759 = −268 kJ mol⁻¹ (lattice enthalpy here is for separating the solid). (HL only)

1.3 Energy from fuels

Complete vs incomplete combustion of an organic fuel: Complete: excess oxygen; CO₂ and H₂O only; blue, clean flame. Incomplete: limited oxygen; carbon ends up as CO and/or C (soot); yellow, sooty flame; less energy per mole.
Balanced equation for the complete combustion of propan-2-ol, C₃H₇OH: 2C₃H₇OH + 9O₂ → 6CO₂ + 8H₂O (3CO₂ + 4H₂O need 10 O atoms; subtract the 1 in the alcohol: 4½O₂ per molecule).
Why is carbon monoxide dangerous?: It attaches to haemoglobin more tightly than oxygen, so the blood carries less oxygen; it has no colour or smell.
Coal, crude oil and natural gas: energy per gram and CO₂ per unit of energy: Natural gas (mostly methane): highest energy per gram, lowest CO₂ per unit energy. Coal (mostly carbon): lowest energy per gram, highest CO₂ per unit energy. Crude oil in between.
How do greenhouse gases warm the lower atmosphere?: The ground absorbs visible light and emits infrared. CO₂, H₂O and CH₄ absorb the infrared (their bonds vibrate) and re-emit it randomly, so part of it returns to the ground.
Why is a biofuel rarely fully carbon neutral?: Its carbon was fixed recently by photosynthesis (6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂), but farming, fertiliser, processing and transport usually burn fossil fuels.
Hydrogen fuel cell half-equations with an acidic electrolyte: Anode (negative, oxidation): H₂ → 2H⁺ + 2e⁻. Cathode (positive, reduction): O₂ + 4H⁺ + 4e⁻ → 2H₂O. Overall 2H₂ + O₂ → 2H₂O.

1.4 Entropy and spontaneity (Additional higher level)

Entropy, and the standard entropy of an element: A measure of how dispersed the matter and energy of a system are (J K⁻¹ mol⁻¹). S⦵ is positive for every pure substance above 0 K, elements included; a perfect crystal at 0 K has zero entropy. (HL only)
When is ΔS of the system positive?: When the moles of gas increase (strongest effect); when a solid melts, a liquid boils or a solid sublimes; when a solid dissolves or substances mix; when a substance is heated. (HL only)
Gibbs energy equation and its units: ΔG⦵ = ΔH⦵ − TΔS⦵, with ΔH and ΔG in kJ mol⁻¹, ΔS divided by 1000 to kJ K⁻¹ mol⁻¹, T in K. Spontaneous when ΔG is negative. (HL only)
Spontaneity from the signs of ΔH and ΔS: ΔH −, ΔS +: at all temperatures. ΔH +, ΔS −: never. ΔH −, ΔS −: below T = ΔH/ΔS. ΔH +, ΔS +: above T = ΔH/ΔS. (HL only)
ΔH = +65.0 kJ mol⁻¹, ΔS = +160 J K⁻¹ mol⁻¹. When is the change spontaneous?: T = 65.0 ÷ 0.160 = 406 K. Both positive, so spontaneous above 406 K. (HL only)
Link between ΔG⦵ and K: ΔG⦵ = −RT lnK (ΔG⦵ in J mol⁻¹, R = 8.31 J K⁻¹ mol⁻¹). Negative ΔG⦵ gives K > 1, products favoured; positive gives K < 1. At equilibrium ΔG = 0. (HL only)
ΔG for a mixture that is not at equilibrium: ΔG = ΔG⦵ + RT lnQ. If Q < K, ΔG is negative and the forward reaction proceeds, even if ΔG⦵ is positive. (HL only)