Properties Of Period 3 Elements & Their Oxides

## Compare sodium and magnesium with cold water - Sodium reacts vigorously with cold water, forming sodium hydroxide solution and hydrogen. Heat can melt the metal; sodium has a low melting point for a metal. - Magnesium reacts much more slowly with cold water. Its oxide/hydroxide surface layer limits contact, and sparingly soluble Mg(OH)₂ is mainly a solid product. - The dissolved hydroxide raises pH. A measured pH depends on amount, water volume, temperature and equilibrium; it is not a fixed property of the metal name. ## Distinguish steam from cold-water products - Heated magnesium reacts with steam to give MgO(s) and H₂(g). This is faster than the usual cold-water reaction. - Use Mg+H₂O(g)→MgO+H₂ for the steam model, rather than copying the hydroxide product from cold water. - One mole of magnesium forms one mole of hydrogen in either balanced equation; the water coefficient and solid product differ. ## Track electron transfer in metal–water reactions - Sodium changes from 0 to+1, while magnesium changes from 0 to+2. Both metals lose electrons and act as reducing agents. - Hydrogen changes from +1 in water to 0 in H₂, so water is reduced in the bookkeeping. Oxygen remains−2. - Balance atoms and total electron transfer. Two Na atoms supply the two electrons required to make one H₂ molecule. ## Write the specified sodium and magnesium oxides - The AQA simple-oxide models are4Na+O₂→2Na₂O and 2Mg+O₂→2MgO. Their products are white ionic solids. - The specified sodium oxide is not a claim that every sodium–oxygen experiment gives only Na₂O. Peroxide formation can compete, especially in excess oxygen. - State the product model before calculating amounts. For Na₂O the Na:oxide ratio is2:1; for MgO the Mg:oxide ratio is1:1. ## Explain aluminium protection and silicon oxidation - Aluminium forms a thin adherent oxide layer that restricts further attack. Heated powder can react much more readily because surface area and coating conditions differ. - 4Al+3O₂→2Al₂O₃ and Si+O₂→SiO₂ conserve both element counts. The products have different bonding and structure. - A lack of rapid reaction at room temperature can reflect kinetics and surface protection; it does not mean oxide formation is thermodynamically impossible. ## Form phosphorus(V) oxide with the oxygen condition stated - Excess oxygen favours phosphorus(V) oxide, written P₄O₁₀ in this molecular model. Its empirical formula is P₂O₅. - P₄+5O₂→P₄O₁₀ is a balanced molecular equation. Atomic P notation can represent other phosphorus allotropes without claiming that all contain P₄ molecules. - Limited oxygen can give phosphorus(III) oxide. Do not claim that every combustion must produce the highest oxidation state. ## Separate sulfur combustion from further oxidation - Burning sulfur mainly gives SO₂, with sulfur in oxidation state +4. SO₃ contains sulfur at+6. - Sulfur trioxide is formed by further oxidation:2SO₂+O₂⇌2SO₃, using the catalytic Contact-process conditions. - A catalyst speeds approach to equilibrium; it does not change the equilibrium constant. Avoid replacing this route with an unqualified claim of direct sulfur combustion to SO₃. ## Use structure rather than a metal/non-metal slogan - Na₂O and MgO form extended ionic lattices. Al₂O₃ is commonly treated as ionic at this level, with significant covalent character from polarization. - Silicon dioxide forms an extended covalent network, while SO₂ gas and a molecular form of P₄O₁₀ consist of discrete molecules. - Across these oxides the increasing covalent contribution is useful, but composition alone does not specify every solid form or predict every property. ## Explain high ionic melting temperatures with limits - Melting an ionic oxide requires disruption of strong electrostatic attractions throughout its lattice. These are not attractions between separate Na₂O molecules. - Mg²⁺ is smaller and more highly charged than Na⁺; its stronger attraction to oxide helps explain MgO’s higher melting temperature. - Do not extrapolate a monotonic charge-only sequence through Al₂O₃. Crystal structure, covalent character and entropy also influence melting. ## Connect silica connectivity to its formula and melting - The reviewed network illustration shows silicon linked through oxygen bridges. Each Si has four O neighbours; each bridging O links two Si atoms. - Sharing oxygen between neighbouring Si centres gives the overall SiO₂ ratio, rather than a set of isolated O=Si=O molecules in ordinary silica. - Melting requires extensive reorganisation of a strong covalent network. It does not require every bond in the specimen to break simultaneously. ## Qualify the molecular-oxide melting pattern - The broad highest-oxide pattern is high melting for Na₂O, MgO, Al₂O₃ and SiO₂, then much lower phase-change temperatures for molecular forms of P₄O₁₀ and SO₃. MgO is highest in the usual comparison; the sequence is not monotonic. - SO₂ is a gas at 25°C. The specified γ-SO₃ form is liquid between its roughly16.8°C melting and 44.5°C boiling temperatures at ordinary pressure. - SO₃ has multiple condensed forms; P₄O₁₀ also has different structures and can sublime. Avoid an exact universal melting curve or the claim that both sulfur oxides are always room-temperature gases. ## Form hydroxide when sodium oxide reacts with water - Na₂O(s)+H₂O(l)→2NaOH(aq) forms a strongly alkaline solution when enough oxide is present. The pH still depends on concentration. - Oxide ions in the lattice accept protons from water. The formal bookkeeping O²⁻+H₂O→2OH⁻ does not mean free oxide ions persist in aqueous solution. - For the supplied ideal calculation, 0.00500 mol Na₂O supplies0.0100 mol OH⁻ in 0.100 dm³, so[OH⁻]=0.100 mol dm⁻³ and pH 13.00 at 25°C. ## Distinguish hydroxide solubility from base strength - MgO can hydrate to Mg(OH)₂. The hydroxide is sparingly soluble, so a suspension can contain much more solid than dissolved material. - A moderately alkaline saturated solution reflects limited dissolved hydroxide; it is not evidence that the dissolved hydroxide ion is a weak base. - Use supplied concentrations or solubility data for pH. Do not assign a universal pH 10 to every mass and dilution of MgO or Mg(OH)₂. ## Avoid equating water insolubility with neutral behaviour - Al₂O₃ and SiO₂ show no appreciable reaction with cold water under ordinary conditions. Strong lattices/networks and kinetic barriers matter. - This observation does not make either oxide neutral. Aluminium oxide is amphoteric; silicon dioxide is acidic in its reactions with suitable bases. - Acid/base classification uses appropriate acid and base reactions, not simply whether a solid visibly dissolves in water. ## Form phosphoric acid without calling it strong - Phosphorus(V) oxide reacts with excess water to form phosphoric acid:H₃PO₄. One mole P₄O₁₀ yields four moles of acid. - Phosphoric acid is a weak, triprotic acid. A low measured pH does not by itself make an acid strong, and its three ionisation steps are not equally extensive. - The reaction can be very exothermic. The equation teaches composition and stoichiometry, rather than serving as an instruction to perform an uncontrolled hydration. ## Describe sulfur dioxide in water as an equilibrium system - Sulfur dioxide dissolves in water and gives an acidic solution. H₂SO₃ is conventional sulfurous-acid notation for the aqueous sulfur(IV) system. - Do not describe the solution as pure molecular H₂SO₃. Dissolved SO₂, hydrated species and acid-base equilibria contribute. - Hydrogensulfite HSO₃⁻ and sulfite SO₃²⁻ arise on successive proton loss in the formal acid model; sulfur remains at oxidation state +4. ## Connect sulfur trioxide to sulfuric acid - SO₃ reacts with water to form sulfuric acid, retaining sulfur at oxidation state +6. This hydration is very exothermic. - The conceptual equation is SO₃+H₂O→H₂SO₄. It does not prescribe direct mixing as a classroom or industrial procedure. - In dilute aqueous solution, the first proton loss of H₂SO₄ is treated as complete. The second is an equilibrium and depends on concentration. ## Read the phosphoric-acid displayed structure - The original displayed formula shows a central P with four oxygen neighbours: one conventional P=O and three P–O–H groups. - The oxygen arrangement around P is approximately tetrahedral; a flat displayed projection does not imply a square planar molecule. - The formula specifies three acidic O–H hydrogens. It does not contain three P–H bonds. Single/double-bond placement is a useful formal Lewis description. ## Preserve the oxygen skeleton through phosphate proton loss - Successive proton loss gives H₂PO₄⁻, HPO₄²⁻ and PO₄³⁻. Each step removes H⁺ and makes the remaining species one charge more negative. - The four-oxygen skeleton remains. In a conventional Lewis contributor, an O–H becomes O⁻ when its proton is lost. - Actual anions have delocalised bonding; the drawn single and double bonds are not a claim of permanently distinct measured P–O bond orders. ## Read sulfurous-acid and sulfite-family structures - The formal H₂SO₃ contributor has one S=O and two S–O–H groups. Sulfur has three oxygen neighbours and a lone pair. - One or two proton losses give HSO₃⁻ and SO₃²⁻, keeping the three-oxygen framework. The S–O connections are not S–H bonds. - This formal model supports composition and proton accounting; aqueous speciation and resonance need more than one fixed drawing. ## Read sulfuric-acid and sulfate-family structures - The displayed conventional H₂SO₄ contributor has two S=O bonds and two S–O–H groups. Sulfur has four oxygen neighbours. - First proton loss gives HSO₄⁻; second gives SO₄²⁻. Hydrogens attach to oxygen, and each proton loss changes one O–H group to formal O⁻. - The approximately tetrahedral oxygen arrangement is projected onto the page. Sulfate has delocalised S–O bonding, so the chosen Lewis contributor is not a set of fixed experimental bond types. ## Separate acid dissociation from full neutralisation - In the dilute aqueous model, H₂SO₄’s first dissociation is effectively complete; the second must be represented by an equilibrium. - Do not automatically assume[H⁺]=2c in every unneutralised sulfuric-acid solution. The extent of the second step depends on equilibrium and concentration. - Complete neutralisation by sufficient strong base still requires two moles OH⁻ per mole H₂SO₄; stoichiometric capacity is not the same as initial free[H⁺]. ## Neutralise acid with a basic oxide - Na₂O and MgO act as bases by accepting protons from acids. MgO+2HCl→MgCl₂+H₂O is a molecular example. - In the ionic equations, oxide is consumed and water forms. Magnesium remains +2 and sodium remains +1, so this is not a metal oxidation reaction. - For the supplied MgO amount, 0.0200 mol acid equivalents are consumed, leaving0.00500 mol when the stated acid is in excess. ## Show the basic role of amphoteric aluminium oxide - Aluminium oxide reacts with acid to form aluminium salt and water. Its basic role is shown by proton consumption. - Al₂O₃+6H⁺→2Al³⁺+3H₂O conserves atoms and charge. The aluminium oxidation state remains +3. - This acid reaction alone shows basic behaviour. To demonstrate amphoterism, also show a suitable reaction with base. ## Show aluminium oxide reacting with base - Hot concentrated aqueous alkali can convert Al₂O₃ to tetrahydroxoaluminate(III), [Al(OH)₄]⁻. - The ionic model is Al₂O₃+2OH⁻+3H₂O→2[Al(OH)₄]⁻. Combined with its acid reaction, this demonstrates amphoterism. - Aluminium remains +3 in the complex. Proton transfer and complex formation do not automatically mean a change in oxidation state. ## Classify silica by its reaction with strong base - Silicon dioxide is an acidic oxide even though it does not react appreciably with cold water. Hot concentrated alkali can attack its network. - The simplified metasilicate equation is SiO₂+2OH⁻→SiO₃²⁻+H₂O, or SiO₂+2NaOH→Na₂SiO₃+H₂O. - Actual silicate speciation depends on conditions. Use the stated metasilicate model for this stoichiometric calculation. ## State the endpoint in phosphate neutralisation - With enough strong base for full neutralisation, P₄O₁₀+12OH⁻→4PO₄³⁻+6H₂O accounts for all twelve acidic hydrogen equivalents after hydration. - Partial neutralisation can give dihydrogenphosphate or hydrogenphosphate salts. Do not assume PO₄³⁻ is the only possible product at every pH. - A product statement and endpoint are essential before choosing a mole ratio for a polyprotic acid system. ## Distinguish sulfite, sulfate and partial neutralisation - Sulfur dioxide forms hydrogensulfite with one OH⁻ equivalent or sulfite with two, depending on conditions and endpoint. - Sulfur trioxide forms sulfate with sufficient OH⁻. Keep SO₃²⁻, the sulfite ion, separate from neutral SO₃, the sulfur(VI) oxide. - No oxidation-state change is required in these neutralisations: sulfur remains +4 in the sulfur(IV) set and+6 in the sulfur(VI) set. ## Interpret evidence with reagent conditions recorded - A water observation alone cannot distinguish Al₂O₃ from SiO₂. Compare suitable acid and alkali reactions as well. - Use separate small samples, comparable conditions and a blank when reviewing teacher-supplied evidence. Record solid, solution and pH observations distinctly. - The supplied unknown gives a silicate in hot alkali, supporting SiO₂. A neutral water suspension does not establish a neutral oxide. ## Combine reaction ratios with specified measurement data - For complete reaction, 1.20 g Mg atAr 24.0 corresponds to 0.0500 mol. The steam equation gives 0.0500 mol H₂. - Using the supplied molar volume24.0 dm³mol⁻¹ gives 1.20dm³ hydrogen. The molar volume is an input for stated conditions, not a universal gas constant. - Check limiting reagents, gas losses, units and significant figures before comparing a calculated amount with an experimental observation.