Flashcards · 4CH1 · Topic 2

Inorganic chemistry flashcards

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a Group 1 (alkali metals)

Observations when lithium, sodium or potassium is added to water: Floats (less dense than water), fizzes (hydrogen), moves around, gets smaller and disappears; universal indicator turns blue/purple (alkaline hydroxide).
General equation for an alkali metal with water: metal + water → metal hydroxide + hydrogen 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)
Lithium, sodium and potassium with water: how do they differ?: Lithium fizzes steadily and stays solid. Sodium fizzes rapidly and melts into a ball. Potassium reacts violently; the hydrogen ignites with a lilac flame.
Why are alkali metals stored under oil?: To stop them reacting with oxygen and water vapour in the air.
Group 1 trends down the group: reactivity, melting point, hardness: Reactivity increases. Melting point decreases. The metals get softer.
Formulae of caesium hydroxide and rubidium oxide: CsOH and Rb₂O — every Group 1 metal forms a 1+ ion.
Explain why reactivity increases down Group 1 (Paper 2 only): More occupied shells, so the outer electron is further from the nucleus, with more shielding. The attraction to the nucleus is weaker, so the outer electron is lost more easily.

b Group 7 (halogens)

Appearance of chlorine, bromine and iodine at room temperature: Chlorine: pale green (yellow-green) gas. Bromine: red-brown liquid (orange-brown vapour). Iodine: grey-black solid (purple vapour).
Group 7 trends down the group: Colour darkens; gas → liquid → solid; melting and boiling points increase; reactivity decreases.
Predict the state and colour of fluorine, and of astatine: Fluorine: a gas, paler than chlorine (pale yellow). Astatine: a solid, darker than iodine (black).
Halogen displacement rule: A more reactive halogen displaces a less reactive halogen from a solution of its halide.
Chlorine water is added to potassium iodide solution. What do you see? Give the ionic equation.: The solution turns brown (iodine forms). Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)
Bromine water is added to sodium chloride solution. What happens?: No reaction — it stays orange. Bromine is less reactive than chlorine, so it cannot displace it.
In Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, what is oxidised and what is reduced?: Bromide ions are oxidised (lose electrons). Chlorine is reduced (gains electrons) — it is the oxidising agent.
Explain why reactivity decreases down Group 7 (Paper 2 only): The outer shell is further from the nucleus, with more shielding, so the attraction for an incoming electron is weaker and an electron is gained less easily.

c Gases in the atmosphere

Composition of dry air: the four most abundant gases: Nitrogen 78%, oxygen 21%, argon 0.9%, carbon dioxide 0.04% (by volume).
Three ways to find the percentage of oxygen in air: Damp iron wool rusting in an inverted measuring cylinder over water; burning phosphorus over water; passing air over hot copper between two gas syringes.
200 cm³ of air over damp iron wool falls to a constant 158 cm³. Percentage of oxygen?: (200 − 158) ÷ 200 × 100 = 21.0%. Always divide by the starting volume.
Gas-syringe method: why let the apparatus cool before the final reading?: Warm gas has expanded, so the final volume reads too large and the percentage of oxygen comes out too low.
Magnesium, hydrogen and sulfur burning in oxygen: observation and product: Magnesium: very bright white flame → white magnesium oxide. Hydrogen: pale blue flame → water. Sulfur: blue flame → sulfur dioxide (SO₂, not SO₃).
Thermal decomposition of copper(II) carbonate: colour change and equation: Green → black. CuCO₃(s) → CuO(s) + CO₂(g)
Heating a carbonate with a delivery tube into limewater: why remove the tube before stopping heating?: To prevent suck-back of cold limewater into the hot tube.
How does carbon dioxide affect the Earth’s temperature?: The surface gives out infrared radiation; CO₂ absorbs some and re-emits some back. More CO₂ (burning fossil fuels, deforestation) keeps more heat in, which may contribute to climate change.

d Reactivity series

The 4CH1 reactivity series, most to least reactive: Potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold.
Which reactions are used to place metals in the reactivity series?: Reactions with water; with dilute hydrochloric or sulfuric acid; and displacement reactions (metal + metal oxide, metal + aqueous metal salt).
Will zinc react with copper(II) sulfate solution?: Yes. Zinc is more reactive than copper, so it displaces it: zinc sulfate and copper metal form. Zinc is oxidised (Zn → Zn²⁺); copper ions are reduced (Cu²⁺ → Cu).
Oxidation and reduction in terms of oxygen: Oxidation: gain of oxygen. Reduction: loss of oxygen.
Oxidation and reduction in terms of electrons: Oxidation: loss of electrons. Reduction: gain of electrons.
Redox, oxidising agent and reducing agent: Redox: oxidation and reduction happen at the same time. Oxidising agent: is reduced (gains electrons). Reducing agent: is oxidised (loses electrons).
Conditions needed for iron to rust: Both water and oxygen together; neither alone is enough.
Three ways to prevent rusting, and how each works: Barrier (paint, oil): keeps out water and oxygen. Galvanising: a zinc coating — a barrier, and sacrificial if scratched. Sacrificial protection: a more reactive metal (zinc, magnesium) corrodes instead of the iron.

e Extraction and uses of metals

Ore (Paper 2 only): A rock containing enough of a metal compound to make extracting the metal worthwhile.
Why is gold found as the uncombined element? (Paper 2 only): It is very unreactive, so it does not react with substances in the ground.
How does reactivity decide the extraction method? (Paper 2 only): Carbon sits between aluminium and zinc. Below carbon (Zn, Fe, Cu): heat the oxide with carbon. Above carbon: electrolysis of a molten compound, because carbon cannot remove the oxygen.
Reduction of iron(III) oxide: equation, and what is reduced and oxidised (Paper 2 only): Fe₂O₃ + 3CO → 2Fe + 3CO₂ Iron(III) oxide is reduced; carbon monoxide is oxidised and is the reducing agent.
Why is aluminium oxide dissolved in molten cryolite? (Paper 2 only): Al₂O₃ melts above 2000 °C; the solution in cryolite can be electrolysed below 1000 °C, saving energy. (Cryolite is not a catalyst.)
Electrode half-equations in aluminium extraction, and why the anodes are replaced (Paper 2 only): Cathode: Al³⁺ + 3e⁻ → Al Anode: 2O²⁻ → O₂ + 4e⁻ The oxygen reacts with the hot carbon anodes to form CO₂, so they burn away.
Uses of aluminium and copper, with the property behind each (Paper 2 only): Aluminium: aircraft and power cables (low density); cans and foil (resists corrosion). Copper: wiring (very good electrical conductor, ductile); water pipes (does not react with water).
Mild steel, high-carbon steel and stainless steel: property and use (Paper 2 only): Mild (low-carbon): strong, malleable, rusts — car bodies. High-carbon: very hard but brittle — cutting tools. Stainless (chromium, often nickel): resists corrosion — cutlery.
Alloy, and why an alloy is harder than the pure metal (Paper 2 only): A mixture of a metal with one or more other elements. Atoms of a different size distort the regular layers, so the layers slide over each other less easily.
Maximum mass of iron from 400 tonnes of Fe₂O₃ (Ar: Fe = 56, O = 16) (Paper 2 only): Mr Fe₂O₃ = 160, containing 112 of iron. 400 × 112 ÷ 160 = 280 tonnes.

f Acids, alkalis and titrations

Colours of litmus, phenolphthalein and methyl orange in acid and in alkali: Litmus: red in acid, blue in alkali. Phenolphthalein: colourless in acid, pink in alkali. Methyl orange: red in acid, yellow in alkali.
pH bands: 0–3 strongly acidic; 4–6 weakly acidic; 7 neutral; 8–10 weakly alkaline; 11–14 strongly alkaline.
Universal indicator colours at pH 1, 7 and 14: pH 1: red. pH 7: green. pH 14: purple. It gives only an approximate pH.
Which ions do acids and alkalis provide in aqueous solution?: Acids: H⁺ ions. Alkalis: OH⁻ ions.
Ionic equation for neutralisation of an acid by an alkali: H⁺(aq) + OH⁻(aq) → H₂O(l)
Acid–alkali titration: outline method (Paper 2 only): Pipette 25.0 cm³ of alkali into a conical flask; add phenolphthalein or methyl orange; add acid from a burette with swirling, dropwise near the end; stop at the first permanent colour change. Repeat until titres are concordant.
Titres: 19.10 (rough), 18.45, 18.80, 18.50, 18.40 cm³. Mean titre? (Paper 2 only): Use only the concordant titres (within 0.10 cm³): (18.45 + 18.50 + 18.40) ÷ 3 = 18.45 cm³.
Why is universal indicator not used in a titration? (Paper 2 only): It changes colour gradually, so there is no sharp end point.

g Acids, bases and salt preparations

Which salts are always soluble?: All sodium, potassium and ammonium compounds, and all nitrates.
Insoluble chlorides and insoluble sulfates: Chlorides: silver and lead(II). Sulfates: barium, calcium and lead(II).
Which carbonates and hydroxides are soluble?: Carbonates: sodium, potassium, ammonium. Hydroxides: sodium, potassium, and calcium (only slightly).
Acid and base in terms of protons: Acid: a proton (H⁺) donor. Base: a proton acceptor.
Base and alkali: Bases: metal oxides, metal hydroxides and ammonia. An alkali is a base that is soluble in water — so magnesium hydroxide is a base but not an alkali.
acid + metal → acid + base → acid + carbonate →: salt + hydrogen (not with nitric acid) salt + water salt + water + carbon dioxide
Making hydrated copper(II) sulfate from copper(II) oxide: Warm dilute sulfuric acid; add CuO until in excess; filter off the excess; evaporate some water until saturated; cool to crystallise; filter and dry. Do not evaporate to dryness.
Making a soluble sodium, potassium or ammonium salt (Paper 2 only): Titrate to find the volumes needed, then mix the same volumes again without indicator, and crystallise. Both reactants are soluble, so there is no excess to filter off.
Making an insoluble salt such as lead(II) sulfate (Paper 2 only): Mix two solutions (e.g. lead(II) nitrate and a soluble sulfate), filter off the precipitate, wash with distilled water, dry. Pb²⁺(aq) + SO₄²⁻(aq) → PbSO₄(s)

h Chemical tests

Tests for hydrogen and oxygen: Hydrogen: lighted splint at the mouth of the tube — squeaky pop. Oxygen: glowing splint — relights.
Tests for carbon dioxide, ammonia and chlorine: CO₂: bubble through limewater — turns milky. NH₃: damp red litmus paper — turns blue. Cl₂: damp litmus paper — bleached white.
Flame test: method: Clean a nichrome or platinum wire in concentrated HCl and heat until there is no colour; dip it in the acid then the solid; hold it in a blue (non-luminous) Bunsen flame; record the colour.
Flame colours of Li⁺, Na⁺, K⁺, Ca²⁺ and Cu²⁺: Li⁺: red. Na⁺: yellow. K⁺: lilac. Ca²⁺: orange-red. Cu²⁺: blue-green.
Sodium hydroxide solution: precipitate colours with Cu²⁺, Fe²⁺ and Fe³⁺: Cu²⁺: blue. Fe²⁺: green. Fe³⁺: brown.
Test for the ammonium ion, NH₄⁺: Add sodium hydroxide solution and warm. No precipitate; the gas given off (ammonia) turns damp red litmus blue.
Test for halide ions, and the precipitate colours: Add dilute nitric acid, then silver nitrate solution. Cl⁻: white. Br⁻: cream. I⁻: yellow.
Test for sulfate ions (4CH1 reagents): Add dilute hydrochloric acid, then barium chloride solution: white precipitate of BaSO₄.
Why use nitric acid (not hydrochloric) before silver nitrate?: Hydrochloric acid adds Cl⁻ ions, which always give a white precipitate. The acid is there to remove carbonate ions, which would also give a precipitate.
Chemical test for water, and physical test for pure water: Anhydrous copper(II) sulfate turns from white to blue (water present). Pure water boils at exactly 100 °C and freezes at exactly 0 °C.