Flashcards · H432 · Topic 3

Periodic table and energy flashcards

44 cards on Topic 3, subtopics 3.1.1 to 3.2.3. Pick a subtopic or study them all.

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3.1.1 Periodicity

First ionisation energy: The energy required to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of gaseous 1+ ions. X(g) → X⁺(g) + e⁻
Why is there a small fall in first ionisation energy from N to O?: In oxygen the electron is removed from a 2p orbital that holds a pair, and repulsion between the paired electrons makes it easier to remove.
Why is there a small fall in first ionisation energy from Be to B (or Mg to Al)?: The electron is removed from a p sub-shell, which is higher in energy than the s sub-shell.
How did Mendeleev’s table gain acceptance?: He grouped elements by properties, left gaps and predicted missing elements; the discovery of germanium (1886), matching his forecast, supported the Periodic Law.
Why does graphene conduct and why is it strong?: Delocalised electrons move along the sheet; every atom is held by strong covalent bonds across the sheet.
Why is silicon’s melting point far above phosphorus’s?: Silicon is giant covalent, so many strong Si–Si bonds must break. Phosphorus is simple molecular P₄: only weak London forces between molecules are overcome.
Ionisation energy equation: Gaseous atom loses an electron. Not a solid, not a gained electron.

3.1.2 Group 2

Equation for the second ionisation energy of strontium: Sr⁺(g) → Sr²⁺(g) + e⁻
Why does reactivity increase down Group 2?: Atomic radius and shielding increase, so nuclear attraction on the outer electrons decreases (despite more protons). The first and second ionisation energies fall, so the two outer electrons are lost more easily.
Magnesium with steam: Mg(s) + H₂O(g) → MgO(s) + H₂(g). The product is the oxide (with cold water, Mg(OH)₂ forms very slowly).
Group 2 oxides with water: trend in pH, and is it redox?: MO + H₂O → M(OH)₂. The hydroxides get more soluble down the group, so the solutions become more alkaline (about pH 10 for MgO to 13 or above for BaO). No oxidation numbers change: not redox.
Mass of Mg(OH)₂ (Mr 58.3) that neutralises 25.0 cm³ of 0.400 mol dm⁻³ HCl: n(HCl) = 0.0100 mol; n(Mg(OH)₂) = 0.00500 mol; mass = 0.292 g.
Two uses of Group 2 compounds as bases: Ca(OH)₂ neutralises acidic soil. Mg(OH)₂ (and CaCO₃) are antacids for excess stomach acid.

3.1.3 The halogens

Why does the boiling point increase from Cl₂ to I₂?: More electrons per molecule (34, 70, 106), so stronger London forces between molecules, which need more energy to overcome. The X–X bond stays intact.
Colours of Cl₂, Br₂ and I₂ in water and in cyclohexane: Cl₂: pale green in both. Br₂: orange in both. I₂: brown in water, violet in cyclohexane.
Why does halogen reactivity decrease down Group 17?: Atomic radius and shielding increase, so there is weaker nuclear attraction for an incoming electron and a 1− ion forms less easily.
Disproportionation: Oxidation and reduction of the same element in one reaction.
Chlorine with cold, dilute NaOH: equation and oxidation numbers: Cl₂ + 2NaOH → NaCl + NaClO + H₂O. Chlorine goes from 0 to −1 (NaCl) and 0 to +1 (NaClO, sodium chlorate(I)).
Benefits and risks of chlorinating drinking water: Benefit: kills bacteria, preventing diseases such as cholera; a residual amount keeps water safe in the pipes. Risks: chlorine is toxic; chlorinated hydrocarbons can form, some suspected carcinogens.
Silver halides: colours and solubility in ammonia: AgCl: white, soluble in dilute NH₃. AgBr: cream, soluble in concentrated NH₃ only. AgI: yellow, insoluble in both.

3.1.4 Qualitative analysis

Order of anion tests, and why: Carbonate, then sulfate, then halide. BaCO₃ and Ag₂SO₄ are both insoluble, so carbonate must go before Ba²⁺ is added and sulfate before Ag⁺ is added.
Why acidify with nitric acid rather than HCl or H₂SO₄ before anion tests?: HCl adds Cl⁻ and H₂SO₄ adds SO₄²⁻, which would give false positives. Nitrate ions precipitate with none of the reagents.
Test for sulfate ions in a sequence of tests: Add Ba(NO₃)₂(aq) (not BaCl₂, which adds chloride): a white precipitate of BaSO₄.
Test for ammonium ions: Warm with NaOH(aq); hold damp red litmus at the mouth of the tube: it turns blue. NH₄⁺ + OH⁻ → NH₃ + H₂O
Why must the litmus be damp in the ammonium test?: NH₃ must dissolve in water to give OH⁻ ions before the litmus responds.

3.2.1 Enthalpy changes

Activation energy, Ea: The minimum energy required for a reaction to take place.
Standard state: The physical state of a substance under standard conditions, e.g. Br₂(l), Na(s), N₂(g).
Enthalpy change of neutralisation, ΔneutH: The enthalpy change when an acid and a base react to form 1 mol of water, under standard conditions, all in standard states.
Standard enthalpy change of combustion, ΔcH: The enthalpy change when 1 mol of a substance burns completely in oxygen, under standard conditions, all in standard states.
2.12 g LiCl (M = 42.4) dissolves in 60.0 g water; temperature rises 7.4 K (c = 4.18 J g⁻¹ K⁻¹). ΔH?: q = 60.0 × 4.18 × 7.4 = 1856 J = 1.856 kJ; n = 0.0500 mol; ΔH = −1.856 ÷ 0.0500 = −37.1 kJ mol⁻¹
Why are experimental ΔcH values usually less exothermic than data-book values?: Heat loss to the air, heat absorbed by the can, incomplete combustion, evaporation of fuel, and non-standard conditions.
Average vs actual bond enthalpy: The average comes from many compounds; the actual value in one molecule may be different. This is why bond enthalpy answers differ from ΔfH-based answers.

3.2.2 Reaction rates

How does increasing concentration or gas pressure increase the rate?: More particles per unit volume, so more frequent collisions, and more successful collisions per second. The energy of each collision is unchanged.
How does a catalyst increase the rate?: It provides a different route with a lower activation energy, so a greater proportion of molecules have energy ≥ Ea. It is not consumed overall and ΔH is unchanged.
Homogeneous vs heterogeneous catalyst, with examples: Homogeneous: same physical state as the reactants (H⁺ in esterification; Cl radicals in ozone breakdown). Heterogeneous: different physical state (Fe in ammonia production; Ni in hydrogenation).
How do catalysts improve sustainability?: Industrial processes reach a useful rate at a lower temperature, so less energy is needed, less fossil fuel is burned and less CO₂ is emitted. Against this, some catalysts are toxic.
Boltzmann distribution: effect of raising the temperature: The peak is lower and further right, with the same area; the Ea line stays put, and a greater proportion of molecules have energy ≥ Ea.
Converting a gas-volume rate of 1.2 cm³ s⁻¹ to mol s⁻¹ at RTP: 1.2 ÷ 24 000 = 5.0 × 10⁻⁵ mol s⁻¹

3.2.3 Chemical equilibrium

Dynamic equilibrium: In a closed system, the forward and reverse reactions run at the same rate, so no concentration in the mixture changes.
2CO(g) + O₂(g) ⇌ 2CO₂(g), forward exothermic. Effect of raising pressure and raising temperature: Pressure: shifts right (3 gas moles → 2), more CO₂. Temperature: shifts left (endothermic direction), less CO₂.
[Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O turns blue when warmed. What does that show?: The position shifts right on heating, so the forward reaction is endothermic. (Pink [Co(H₂O)₆]²⁺, blue [CoCl₄]²⁻.)
Haber process compromise conditions: About 450 °C (low temperature favours yield but is too slow), about 2 × 10⁴ kPa (high pressure favours yield but is costly and a safety risk), iron catalyst; unreacted gases are recycled.
2NOCl ⇌ 2NO + Cl₂; [NOCl] = 0.250, [NO] = 0.0500, [Cl₂] = 0.0250 mol dm⁻³. Kc?: Kc = [NO]²[Cl₂] ÷ [NOCl]² = (0.0500² × 0.0250) ÷ 0.250² = 1.00 × 10⁻³
What does the size of Kc tell you?: Kc ≫ 1: position far to the right (mostly products). Kc ≈ 1: similar amounts. Kc ≪ 1: far to the left (mostly reactants).