Periodicity Period 3 - elements Can you write one fact about each of the following elements? Na Mg Al Si P S Cl Ar Candidates should be able to: •describe qualitatively (and indicate the periodicity in) the variations in atomic radius, ionic radius, melting point and electrical conductivity of the elements (see the Data Booklet) •explain qualitatively the variation in atomic radius and ionic radius •interpret the variation in melting point and in electrical conductivity in terms of the presence of simple molecular, giant molecular or metallic bonding in the elements •explain the variation in first ionisation energy. Na [Ne] 3s1 Mg [Ne] 3s2 Al [Ne] 3s2 3px1 Si P S Cl Ar [Ne] 3s2 3px1 3py1 [Ne] 3s2 3px1 3py1 3pz1 [Ne] 3s2 3px2 3py1 3pz1 [Ne] 3s2 3px2 3py2 3pz1 [Ne] 3s2 3px2 3py2 3pz2 Atomic radius 0.25 Atomic radius (nm) 0.2 0.15 0.1 0.05 0 Na Mg Al Si P Element S Cl Ar Atomic radius – what about argon? www.chemguide.co.uk/atoms/properties/atradius.html What is atomic radius? www.rjclarkson.demon.co.uk/found/period3_atomrad.gif Explaining the trend An atomic radius is a measure of the distance from the nucleus to the bonding pair of electrons. From sodium to chlorine, the bonding electrons are all in the 3rd shell being screened by the electrons in the first and second levels, i.e. the screening remains fairly constant. The increasing nuclear charge as you go across the period pulls the bonding electrons more tightly towards it. You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases. Leaving the noble gases out, atoms get smaller as you go across a period. IONIC RADIUS Ions aren't the same size as the atoms they come from. Compare the sizes of sodium and chloride ions with the sizes of sodium and chlorine atoms. Positive ions Positive ions are smaller than the atoms they come from. Sodium is 2,8,1; Na+ is 2,8. You've lost a whole layer of electrons, and the remaining 10 electrons are being pulled in by the full force of 11 protons. Negative ions Negative ions are bigger than the atoms they come from. Chlorine is 2,8,7; Clis 2,8,8. Although the electrons are still all in the 3-level, the extra repulsion produced by the incoming electron causes the atom to expand. There are still only 17 protons, but they are now having to hold 18 electrons. Pauling electronegativity value Electronegativity 3.5 3 2.5 2 1.5 1 0.5 0 Na Mg Al Si P Element S Cl Ar First ionisation energy (kJ mol-1) First ionisation energy 1600 1400 1200 1000 800 600 400 200 0 Na Mg Al Si P Element S Cl Ar Electronegativity Bonding, Structure and Properties Structure Type of element Bonding Formula Type of force broken on melting/ boiling Does the element conduct electricity ? Na Mg Al Si P S Cl Ar Giant metallic Metal Giant metallic Metal Giant metallic Metal Giant Simple Simple Simple Monatomic covalent molecule molecule molecule Non-metal Non-metal Non-metal Non-metal Non-metal Metallic Metallic Metallic Covalent Metallic bond Metallic bond Metallic bond Covalent bond Yes Yes Yes No Covalent P4 vdW Covalent S8 vdW Covalent Cl2 vdW Ar vdW No No No No Structure and Properties Electrical conductivity Melting and boiling points 3000 2500 2000 1500 1000 500 0 Periodicity Period 3 - oxides Starter activity Complete these sketches to show how these properties change as you go along Period 3 from Na to Ar. Candidates should be able to: describe the reactions, if any, of the elements with oxygen to give Na2O, MgO, Al2O3, P4O10, SO2 and SO3. state and explain the variation in oxidation number of the oxides. describe the reactions of the oxides with water. describe and explain the acid/base behaviour of oxides and hydroxides, including, where relevant, amphoteric behaviour in reaction with NaOH and acids. Reactions with oxygen Group number Element in Period 3 Nuclear charge [Ne] electronic configuration Trend in Atomic radius Trend in 1st ionisation energy Trend in electronegativity Formula of oxide/s Oxidation state 1 2 Na Mg 11+ 3s1 12+ 3s2 3 Al 4 Si 5 P 6 S 13+ 14+ 15+ 16+ 3s2 3p1 3s2 3p2 3s2 3p3 3s2 3p4 decreases increases Na2O MgO Al2O3 SiO2 P4O10 +1 +2 +3 +4 +5 increases SO2/ SO3 +4/+6 Reactions with oxygen Sodium Sodium burns in oxygen with an orange flame to produce the white solid sodium oxide. Magnesium Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide. Reactions with oxygen Aluminium Aluminium will burn in oxygen if it is powdered, otherwise the strong oxide layer on the aluminium tends to inhibit the reaction. If you sprinkle aluminium powder into a Bunsen flame, you get white sparkles. White aluminium oxide is formed. Silicon Silicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced. Reactions with oxygen Phosphorus White phosphorus catches fire spontaneously in air, burning with a white flame and producing clouds of white smoke - a mixture of phosphorus(III) oxide and phosphorus(V) oxide. The proportions of these depend on the amount of oxygen available. In an excess of oxygen, the product will be almost entirely phosphorus(V) oxide: Reactions with oxygen Sulphur Sulphur burns in air or oxygen on gentle heating with a pale blue flame. It produces colourless sulphur dioxide gas. In an excess of pure oxygen, some SO3 is also formed. This utilises the highest oxidation state of sulphur. Reactions with oxygen Which is which? Melting points of oxides Na2O MgO Al2O3 SiO2 P4O10 SO2 Tm/K 1548 3125 2345 1883 573 200 Bonding Ionic Ionic Ionic Covalent Covalent Covalent Giant lattice Giant lattice Giant lattice Giant lattice Structure Simple Simple molecular molecular Giant ionic and covalent solids contain only strong bonds and have high melting points. Simple molecules have weak vdW forces between molecules and have low melting points. Acid/base properties of the oxides Reaction with water pH Na2O + H2O 2Na+ + 2OH- 14 MgO + H2O Mg2+ + 2OH- 9 Al2O3 insoluble – no reaction 7 SiO2 insoluble – no reaction 7 P4O10 + 6H2O 4H3PO4 0 SO2 + H2O H2SO3 3 SO3 + H2O H2SO4 0 Acid/base behaviour of aluminium oxide BASE: Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O ACID: Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4 Periodicity Period 3 - chlorides Candidates should be able to: •describe the reactions, if any, of the elements with chlorine to give NaCl, MgCl2, Al2Cl6, SiCl4, and PCl5. •describe and explain the reactions of the chlorides with water. Reaction with chlorine Element Na Mg Description of reaction with Very vigorous Vigorous chlorine Formula of NaCl MgCl2 chloride/s Oxidation state of +1 +2 period 3 element State of chloride at Solid Solid r.t.p. b.pt. of 1465 1418 chloride (oC) Structure of Giant lattice chloride Bonding in Ionic chloride Al Si P Vigorous Slow Slow Al2Cl6 SiCl4 PCl3 / PCl5 +3 +4 +3 / +5 Solid Liquid liquid / solid 423 57 74 / 164 Simple molecular Covalent Structure of Al2Cl6 Acid/base properties of the chlorides Reaction with water NaCl(s) Na+(aq) + Cl-(aq) MgCl2(s) Mg2+(aq) + 2Cl-(aq) pH 7 6/7 Al2Cl6(s) + 12H2O(l) 2[Al(H2O)6]3+ 6Cl-(aq) 3 SiCl4(l) + 2H2O(l) SiO2(s) + 4HCl(g) 0 PCl5(l) + 4H2O(l) H3PO4(aq) + 5HCl(g) 0 Periodicity Group 7 - Halogens Starter activity Can you complete task 1 – IGCSE revision on the halogens Candidates should be able to: •describe the trends in volatility and colour of chlorine, bromine and iodine. •interpret the volatility of the elements in terms of van der Waals’ forces. •describe the relative reactivity of the elements as oxidising agents. •describe and explain the reactions of the elements with hydrogen. •describe and explain the relative thermal stabilities of the hydrides. •interpret these relative stabilities in terms of bond energies. The Halogens Ionic Covalent The Halogens Trend in colour and physical state Like dissolves like! Periodicity Group 7 – the halides Starter activity Can you complete the table ‘Reducing power of the halides’? Trend in reducing ability NaX NaF NaCl NaBr NaI Observations steamy fumes steamy fumes steamy fumes colourless gas Products HF HCl HBr SO2 brown fumes steamy fumes colourless gas Br2 HI SO2 yellow solid S smell of bad eggs H2S Grey solid, purple I2 fumes Type of reaction acid-base (F- acting as a base) acid-base (Cl- acting as a base) acid-base (Br- acting as a base) redox (reduction product of H2SO4) redox (oxidation product of Br-) acid-base (I- acting as a base) redox (reduction product of H2SO4) redox (reduction product of H2SO4) redox (reduction product of H2SO4) redox (oxidation product of I-) Candidates should be able to: •describe and explain the reactions of halide ions with aqueous silver ions followed by aqueous ammonia concentrated sulphuric acid. •describe and interpret in terms of changes of oxidation number the reaction of chlorine with cold, and with hot, aqueous sodium hydroxide. •explain the use of chlorine in water purification. •recognise the industrial importance and environmental significance of the halogens and their compounds, (e.g. for bleaches; PVC; halogenated hydrocarbons as solvents, refrigerants and in aerosols). Summary of reducing power Fluoride Relative reducing power Chloride Bromide Iodide Testing for halide ions ion present observation F- No precipitate Cl- White precipitate Br- Cream precipitate I- Yellow precipitate Silver fluoride is soluble, and so you don't get a precipitate. Confirmatory tests original precipitate E.g. observation AgCl Soluble in dilute NH3(aq) AgBr Soluble in concentrated NH3(aq) AgI Insoluble in concentrated NH3(aq) AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Cl-(aq) Reactions of chlorine with NaOH Cl2(g) + 2NaOH(aq) → NaClO(aq) + NaCl(aq) + H2O(l) 0 +1 -1 3Cl2(g) + 6NaOH(aq) → NaClO3(aq) + 5NaCl(aq) + 3H2O(l) Uses of halogens and their compounds bleach refrigerants PVC solvents Aerosol propellants Periodicity Nitrogen and Sulphur Starter activity Can you use information in your textbooks to complete task 9? Candidates should be able to: •explain the lack of reactivity of nitrogen. •describe the oformation, and structure of, the ammonium ion othe displacement of ammonia from its salts. •understand the environmental consequences of the uncontrolled use of nitrate fertilisers. •understand and explain the occurrence, and catalytic removal, of oxides of nitrogen. •explain why atmospheric oxides of nitrogen are pollutants, including their catalytic role in the oxidation of atmospheric sulphur dioxide. •describe the formation of atmospheric sulphur dioxide from the combustion of sulphur contaminated carbonaceous fuels. •state the role of sulphur dioxide in the formation of acid-rain and describe the main environmental consequences of acid-rain. •understand the industrial importance of sulphuric acid. •describe the use of sulphur dioxide in food preservation. Gases in the air Structure of nitrogen Bond dissociation enthalpy = +946 kJ/mol Nitrogen Enzymes High energy The Haber Process The ammonium ion Ammonium NH4+ Add sodium hydroxide solution Ammonia gas is to a solution of the given off. substance and gently heat. Fertiliser n. Any of a large number of natural and synthetic materials, including manure and nitrogen, phosphorus, and potassium compounds, spread on or worked into soil to increase its capacity to support plant growth. Fertilisers Ammonium nitrate Ammonium phosphate Potassium chloride Nitric acid NH3(g) + 2O2(g) HNO3(l) + H2O(l) Pt/Rh catalyst 900oC The environment and fertilisers fertilisers applied to farm land if too much used, at the wrong time of year, during wet weather, excess washed into rivers and lakes causes Excessive growth of aquatic plants. The bacteria which live on dead plants thrive and use up the oxygen in the water. The lack of oxygen causes death of fish. This is called eutrophication. excess contaminates underground drinking water supplies causes Harm to infants - called ‘blue baby’ syndrome Natural sources of SO2 and NOx Sulphur dioxide - SO2 Used as a food preservative: Inhibits growth of moulds, yeasts and aerobic bacteria. Acts as a reducing agent and retards the oxidation of foodstuffs. Acid rain SO2(g) + NO2(g) SO3(g) + NO(g) SO3(g) + H2O(l) H2SO4(l) Uses of sulphuric acid Detergents Paints , pigments and dyestuffs Fertilisers AS Chemistry Periodicity Group 2 Candidates should be able to interpret, and make predictions from, the trends in physical and chemical properties of the elements and their compounds. Trends in Atomic Radius The radius of an atom is governed by: the distance of the outer electrons from the nucleus the nuclear charge, and the amount of shielding Trends in First Ionisation Energy Trends in First Ionisation Energy Ionisation energy is also governed by: the charge on the nucleus, the amount of shielding by the inner electrons, the distance between the outer electrons and the nucleus. Increased nuclear charge ‘off-set’ by increased shielding. Outermost electron increasingly distant from pull of nucleus, less energy needed to remove it. Trends in Electronegativity Trends in Electronegativity Trends in Electronegativity Trend in Electronegativity As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it. Increasing nuclear charge ‘off-set’ by more shielding. In other words, as you go down the Group, the elements become less electronegative. Trend in Melting Point Trend in Melting point Going down Group 2: •the number of delocalised electrons remains the same ... •the charge on each metal cation stays the same at 2+, but ... •the ionic radius increases ... •so the attraction between the delocalised electrons and the metal cations decreases. AS Chemistry Periodicity Group 2 – chemical properties Candidates should be able to: •describe the reactions of the elements with oxygen and water. •describe the behaviour of the oxides with water. •describe the thermal decomposition of the nitrates and carbonates. •interpret, and make predictions from, the trends in chemical properties of the elements and their compounds. •explain the use of magnesium oxide as a refractory lining material and calcium carbonate as a building material. •describe the use of lime in agriculture.