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Topic 3 On A Level Edxcel Inorganic Topic

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Inorganic chemistry A levels contains these conceptual topics to understand the patterns . all topics have been explained in the Inorganic chemistry

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  1. 4. Inorganic Chemistry and the Periodic Table 4A: Group 2 Atomic radius Atomic radius increases down the Group. As one goes down the group, the atoms have more shells Of electrons making the atom bigger. 1st ionisation energy Melting points Down the group the melting points decrease. The metallic bonding weakens as the atomic size increases. The distance between the positive ions and delocalized electrons increases. Therefore the electrostatic attractive forces between the positive ions and the delocalized electrons weaken. The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells. In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons Group 2 reactions Reactivity of group 2 metals increases down the group The reactivity increases down the group as the atomic radii increase there is more shielding. The nuclear attraction decreases and it is easier to remove (outer) electrons and so cations form more easily Reactions with oxygen. The group 2 metals will burn in oxygen. Mg burns with a bright white flame. 2Mg + 02 2Mgo MgO is a white solid with a high melting point due to its ionic bonding. Reactions with chlorine Mg will also react slowly with oxygen without a flame. Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen. 2Mg + 02 2Mgo This needs to be cleaned off by emery paper before doing reactions with Mg ribbon. If testing for reaction rates with Mg and acid, an un-cleaned Mg ribbon would give a false result because both the Mg and MgO would react but at different rates. Mg + 2HCl MgC12 + '-42 Mgo + 2HCl MgC12 + H20 The group 2 metals will react with chlorine Mg + MgC12 Reactions with water. Magnesium reacts in Steam to produce magnesium oxide and hydrogen. The Mg would burn with a bright white flame. Mg (s) + H20 (g) Mgo (s) + (g) Mg will also react with warm water, giving a different magnesium hydroxide product. Mg +2 H20 + This is a much slower reaction than the reaction with steam and there is no flame. The other group 2 metals will react with cold water with increasing vigour down the group to form hydroxides. Ca + 2 H20 (aq) + H2 (g) Sr + 2 H20 (aq) + H2 (g) Ba + 2 H20 (aq) + H2 (g) The hydroxides produced make the water alkaline One would observe: •fizzing, (more vigorous down group) •the metal dissolving, (faster down group) •the solution heating up (more down group) •and with calcium a white precipitate appearing (less precipitate forms down group) N Goalby chemrevise.org 1
  2. Reactions of the Oxides of Group 2 elements with water Group 2 ionic Oxides react with water to form hydroxides The ionic oxides are basic as the oxide ions accept protons to become hydroxide ions in this reaction (acting as a bronsted lowry base) Mgo (s) + H20 (l) (s) pH 9 Mg(OH)2 is only slightly soluble in water so fewer free OH' ions are produced and so lower pH CaO (s) + (l) (aq) PH 12 Reactions of the Oxides of Group 2 elements with Acids Mgo (s) + 2 HCI (aq) MgC12 (aq) + H20 (l) SrO (s) + 2 HCI (aq) srC12 (aq) + H20 (l) CaO (s) + H2S04 (aq) CaS04 (aq) + H20 (l) Reactions of the hydroxides of Group 2 elements with Acids 2HN03 (aq) + (aq) (aq)+ 2H20 (l) 2HCl (aq) + (aq) -5 MgC12 (aq)+ 2H20 (l) Solubility Of hydroxides Group Il hydroxides become more soluble down the group. All Group Il hydroxides when not soluble appear as white precipitates. Magnesium hydroxide is classed as insoluble in water. Simplest Ionic Equation for formation of Mg(OH)2 (s) Mg2+ (aq) + 20H (aq) A suspension of magnesium hydroxide in water will appear slightly alkaline (pH 9) so some hydroxide ions must therefore have been produced by a very slight dissolving. Magnesium hydroxide is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation. + 2HCl MgC12 + 2+420 It is safe to use because it so weakly alkaline. It is preferable to using calcium carbonate as it will not produce carbon dioxide gas. Solubility Of Sulfates Group Il sulfates become less soluble down the group. BaS04 is the least soluble. Calcium hydroxide is reasonably soluble in water. It is used in agriculture to neutralise acidic soils. An aqueous solution of calcium hydroxide is called lime water and can be used a test for carbon dioxide. The limewater turns cloudy as white calcium carbonate is produced. + C02 (g, CaC03(G) + Barium hydroxide would easily dissolve in water. The hydroxide ions present would make the solution strongly alkaline. + aq Ba2* (aq) + 20H (aq) If Barium metal is reacted with sufuric acid it will only react slowly as the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack. Ba + H2S04 BaS04 + H2 The same effect will happen to a lesser extent with metals going up the group as the solubility increases. The same effect does not happen with other acids like hydrochloric or nitric as they form soluble group 2 salts. N Goalby chemrevise.org 2
  3. Thermal decomposition of group 2 carbonates Group 2 carbonates decompose on heating to produce group 2 oxides and carbon dioxide gas. MgCOa(s) MgO(s) + C02(g) Cacog(s) CaO(s) + C02(g) Thermal decomposition is defined as the use of heat to break down a reactant into more than one product The ease of thermal decomposition decreases down the group Group 2 carbonates become more thermally stable going down the group. As the cations get bigger they have less of a polarising effect and distort the carbonate ion less. The C-O bond is weakened less so it less easily breaks down Group 1 carbonates do not decompose with the exception of lithium. As they only have +1 charges they don't have a big enough charge density to polarise the carbonate ion. Lithium is the exception because its ion is small enough to have a polarising effect Li2C03(s) Li20(s) + C02(g) There are a number of experiments that can be done to investigate the ease of decomposition. One is to heat a known mass of carbonate in a side arm boiling tube and pass the gas produced through lime water. Time for the first permanent cloudiness to appear in the limewater. Repeat for different carbonates using the same moles of carbonate/same volume of limewater/same Bunsen flame and height of tube above flame. Heat Thermal decomposition of group 2 nitrates (V) Group 2 nitrates decompose on heating to produce group 2 oxides, oxygen and nitrogen dioxide gas. lime water 2Mgo + 4N02 + 02 You would observe brown gas evolving (N02) and the The ease of thermal decomposition White nitrate solid is seen to melt to a colourless solution decreases down the group and then re-solidify. The explanation for change in thermal stability is the same as for carbonates Magnesium nitrate decomposes the easiest because the Mg2* ion is smallest and has the greater charge density. It causes more polarisation of the nitrate anion and weakens the N—O bond Group 1 nitrates, with the exception of lithium nitrate, do not decompose in the same way as group 2 nitrates. They decompose to give a nitrate (Ill) salt and oxygen. 2NaN03 2NaN02 + 02 Sodium Sodium nitrate(V) nitrate(lll) Lithium nitrate decomposes in the same way as group 2 nitrates 4 LiN03 2Li20 + 4N02 + 02 N Goalby chemrevise.org 3
  4. Flame tests Method Use a nichrome wire ( nichrome is an unreactive metal and will not give out any flame colour) Clean the wire by dipping in concentrated hydrochloric acid and then heating in Bunsen flame If the sample is not powdered then grind it up. Dip wire in solid and put in Bunsen flame and observe flame Explanation for occurrence Of flame In a flame test the heat causes the electron to move to a higher energy level. The electron is unstable at the higher energy level and so drops back down. As it drops back down from the higher to a lower energy level, energy is emitted in the form of visible light energy with the wavelength of the observed light Lithium : Scarlet red Sodium : Yellow Potassium : lilac Rubidium : red Caesium: blue Magnesium: no flame colour (energy emitted of a wavelength outside visible spectrum) Calcium: brick red Strontium: red Barium: apple green N Goalby chemrevise.org 4
  5. 4B Halogens Fluorine (F2): very pale yellow gas. It is highly reactive Chlorine : (C12) greenish, reactive gas, poisonous in high concentrations Bromine (Br2) : red liquid, that gives off dense brown/orange poisonous fumes Iodine (12) : shiny grey solid sublimes to purple gas. Trend in melting point and boiling point Increase down the group As the molecules become larger they have more electrons and so have larger London forces between the molecules. As the intermolecular forces get larger more energy has to be put into break the forces. This increases the melting and boiling points Trend in electronegativity Electronegativity is the relative tendency of an atom in a molecule to attract electrons in a covalent bond to itself. As one goes down the group the electronegativity of the elements decreases. As one goes down the group the atomic radii increases due to the increasing number of shells. The nucleus is therefore less able to attract the bonding pair of electrons The reactivity of the halogens decreases down the group as the atoms get bigger with more shielding so they less easily attract and accept electrons. They therefore form -1 ions less easily down the group 1. The Oxidation reactions Of halide ions by halogens. A halogen that is a strong oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds The oxidising strength decreases down the group. Oxidising agents are electron acceptors. know these Chlorine will displace both bromide and iodide ions; bromine will displace iodide ions potassium chloride (aq) potassium bromide (aq) potassium iod'de (aq) Chlorine (aq) Very pale green solution, no reaction Yellow solution, C has disp aced Br Brown solution, CI has disp aced I Brom ne (aq) Yellow solution, no reaction Yellow solution, no reaction Brown Solution, Br has displaced Observations if an Organic solvent is added potassium chloride (aq) potassium bromide (aq) potassium iodide (aq) Chlorine (aq) colourless, no reaction yellow, Cl has displaced Br purple, Cl has displaced I Bromine (aq) yellow, no reaction yellow, no reaction purple, Br has displaced I (aq) + Bra(aq) (aq) + 12(aq) (aq) + 12(aq) Iodine (aq) Brown solution, no reaction Brown solution, no reaction Brown Solution, no reaction Iodine (aq) purple, no reaction purple, no reaction purple, no reaction observations ! The colour of the solution in the test tube shows which free halogen is present in solution. Chlorine —very pale green solution (often colourless), Bromine = yellow solution Iodine = brown solution (sometimes black solid present) The colour of the organic solvent layer in the test tube shows which free halogen is present in solution. Chlorine = colourless Bromine = yellow Iodine = purple C12(aq) + 2Br- (aq) C12(aq) + 21 (aq) Br2(aq) + 21 (aq) 2Cl 2Cl 2Br N Goalby chemrevise.org 5
  6. The Oxidation reactions Of metals and metal ion by halogens. In all reactions where halogens 2Na 2Na• + 2e Br2(l) + 2Na (s) 2NaBr (s) + 2e•» 2Br are reacting with metals, the metals are being oxidised 3C12(g) +2 Fe (s) 2 FeC13 (s) Br2(l) + Mg (s) -Y MgBr2 (s) C12(g) + 2Fe2• (aq) -5 2 Cl• (aq) + 2Fe3+ (aq) 21• (aq) + 2Fe3• (aq) 12 (aq) + 2Fe2• (aq) Chlorine and Bromine can oxidise Fe2+ to Fe3*- Iodine is not strong enough oxidising agent to do this reaction. The reaction is reversed for Iodine The disproportionation reactions of chlorine. Disproportionation is the name for a reaction where an element simultaneously oxidises and reduces. Chlorine with water: C12(g) + HClO(aq) + NCI (aq) If some universal indicator is added to the solution it will first turn red due to the acidity of both reaction products. It will then turn colourless as the HCIO bleaches the colour. Chlorine is both simultaneously reducing and oxidising changing its oxidation number from O to -1 and Oto +1 The pale greenish colour of these solutions is due to the C12 Chlorine is used in water treatment to kill bacteria. It has been used to treat drinking water and the water in swimming pools. The benefits to health of water treatment by chlorine outweigh its toxic effects. Reaction Of halogens with cold dilute NaOH solution: C12, Br2, and 12 in aqueous solutions will react with cold sodium hydroxide. The colour of the halogen solution will fade to colourless C12(aq) + 2NaOH(aq) -5 NaCl (aq) + NaClO (aq) + H20(l) The mixture of NaCl and NaClO is used as Bleach and to disinfect/ kill bacteria Reaction Of halogens with hot dilute NaOH solution: With hot alkali disproportionation also occurs but the halogen that is oxidised goes to a higher oxidation state. 3C12 (aq) + 6 NaOH(aq) 5 NaC1 (aq) + NaC103 (aq) + 3H20 (l) 312 (aq) + 6NaOH (aq) -5 5 Nal (aq) + Na103 (aq) + 3H20 (l) 312 (aq) + 60H (aq) 5 (aq) + 10/ (aq) + 3H20 (1) In IIJPAC convention the various forms of sulfur and chlorine compounds where oxygen is combined are all called sulfates and chlorates with relevant oxidation number given in roman numerals. If asked to name these compounds remember to add the oxidation number. NaClO: sodium chlorate(l) NaClOa: sodium chlorate(V) K2S04 potassium sulfate(Vl) K2S03 potassium sulfate(lV) N Goalby chemrevise.org 6
  7. The reaction Of halide salts with concentrated sulfuric acid. The halides show increasing power as reducing agents as one goes down the group. This can be clearly demonstrated in the various reactions of the solid halides with concentrated sulfuric acid. Know the equations and Observations of these reactions very well. Fluoride and Chloride Explanation of differing reducing power Of halides A reducing agent donates electrons. The reducing power of the halides increases down group 7 They have a greater tendency to donate electrons. This is because as the ions get bigger it is easier for the outer electrons to be given away as the pull from the nucleus on them becomes smaller. The H2S04 is not strong enough an oxidising reagent to oxidise the chloride and fluoride ions. No redox reactions occur. Only acid-base reactions occur. NaF(s) + H2S04(l) + HF(g) Observations: White steamy fumes of HF are evolved. NaCl(s) + H2S04(l) NaHS04(s) + HC1(g) Observations: White steamy fumes of HCI are evolved. Bromide These are acid —base reactions and not redox reactions. H2S04 plays the role of an acid (proton donor). Br- ions are stronger reducing agents than Cl- and F- and after the initial acid- base reaction reduce the sulfur in H2S04 from +6 to + 4 in S02 Acid- base step: NaBr(s) + H2S04(l) NaHS04(s) + HBr(g) Redox step: 2HBr + H2S04 Br2(g) + S02(g) + 2H20(l) Ox 1/2 equation 2Br - -5 Br2 + 2e• Re 'h equation H2S04 +2 +2e- S02 +2 H20 Observations: White steamy fumes of HBr are evolved. Red fumes of Bromine are also evolved and a colourless, acidic gas Reduction product = sulfur dioxide Note the H2S04 plays the role of acid in the first step producing HBr and then acts as an oxidising agent in the second redox step. Iodide l- ions are the strongest halide reducing agents. They can reduce the sulfur from +6 in H2S04 to + 4 in S02, to O in S and -2 in H2S. Nal(s) + H2S04(l) NaHS04(s) + Hi(g) 2Hl + 12(s) + S02(g) + 2H20(l) 6Hl + H2S04 (s) + 4 H20 (l) 8Hl + H2S04 + H2S(g) + 4H20(l) Ox 1/2 equation 21 • 12 + 2e Re h equation H2S04 + +2e S02 +2 H20 Re h equation H2S04 + 6 +6e- S +4 H20 Re h equation H2S04 + 8 +8e- H2S + 4 H20 Observations: White steamy fumes of HI are evolved. Black solid and purple fumes of Iodine are also evolved A colourless, acidic gas S02 A yellow solid of Sulphur H2S (Hydrogen Sulphide), a gas with a bad egg smell, Reduction products = sulfur dioxide, sulfur and hydrogen sulfide Note the H2S04 plays the role of acid in the first step producing HI and then acts as an oxidising agent in the three redox steps Often in exam questions these redox reactions are worked out after first making the half-equations N Goalby chemrevise.org 7
  8. The reactions of halide ions with silver nitrate. This reaction is used as a test to identify which halide ion is present. The test solution is made acidic with nitric acid, and then silver nitrate solution is added dropwise. The role Of nitric ac'd is to react with any carbonates present to prevent formation Of the precipitate Ag2C03. This would mask the desired Observations 2 HN03 + Nazcos 2 NaN03 + H20 + C02 Effect Of Light on Silver Halides Fluorides produce no precipitate Chlorides produce a white precipitate + Cl- (aq) -5 AgCl(s) Bromides produce a cream precipitate + Br (aq) AgBr(s) Iodides produce a pale yellow precipitate + I- Agl(s) Effect Of ammonia on silver halides The silver halide precipitates can be treated with ammonia solution to help differentiate between them if the colours look similar: Producing hydrogen halides The precipitates ( except Agl) darken in sunlight forming silver. This reaction is used in photography to form the dark bits on photographic film Silver chloride dissolves in dilute ammonia to form a complex ion AgCl(s) + 2NH3(aq) (aq) + Cl- (aq) Colourless solution Silver bromide dissolves in concentrated ammonia to form a complex ion AgBr(s) + 2NH3(aq) (aq) + Br- (aq) Colourless solution Silver iodide does not react with ammonia — it is too insoluble. Hydrogen halides are made by the reaction of solid sodium halide salts with phosphoric acid NaCl(s) + H3P04(l) NaH2P04(s) + HCl(g) Observations: White steamy fumes of the hydrogen halides are evolved. The Steamy fumes of HCI are produced when the HCI meets the air because it dissolves in the moisture in the This is the apparatus used to make the hydrogen halide using phosphoric acid. Notice the downward delivery which is used because the hydrogen halides are more dense than air Phosphoric acid is not an oxidising agent and so does not oxidise HBr and HI. Phosphoric acid is more suitable for producing hydrogen halides than using concentrated sulfuric acid to make HCI, HBr, and HI. This is because there are no extra redox reactions taking place and no other products formed. Solubility in water : The hydrogen halides are all soluble in water. They dissolve to form acidic solutions. HCI H20(l) Hydrogen Halide The water quickl rises up the tube All the hydrogen halides react readily with ammonia to give the white smoke of the ammonium halide HCl(g) + NH2 (g) NH4Cl (s) HBr(g) + NHS (g) NH4Br (s) Hi(g) + NH: (g) NH41 (s) This can be used as a test for the presence of hydrogen halides N Goalby chemrevise.org 8
  9. 4C Analysis of Inorganic Compounds Testing for Negative ions (anions) Testing for presence of a carbonate C032- and hydrogencarbonates HC03- Add any dilute acid and observe effervescence. Bubble gas through limewater to test for C02 — will turn limewater cloudy 2HCl + Na2C03 2NaCl + H20 + C02 HCI + NaHCOg NaCl + + C02 Testing for presence of a sulfate 2H+ + COF- + C02 + HCOa- 1-420 + C02 Acidified BaC12 solution is used as a reagent to test for sulfate ions If Barium Chloride is added to a solution that contains sulfate ions a Fizzing due to C02 would be observed if a carbonate or a hydrogencarbonate was present Ba2+ (aq) + S042-(aq) BaS04 (s). Other anions should give a negative result which is no precipitate forming white precipitate forms The acid is needed to react with carbonate impurities that are often found in salts which would form a white Barium carbonate precipitate and so give a false result Testing for halide ions with silver nitrate. This reaction is used as a test to identify which halide ion is present. The test solution is made acidic with nitric acid, and then silver nitrate solution is added dropwise. Fluorides produce no precipitate Chlorides produce a white precipitate Ag+(aq) + Cl- (aq) AgCl(s) Bromides produce a Cream precipitate Ag+(aq) + Br (aq) -5 AgBr(s) Iodides produce a pale yellow precipitate Ag+(aq) + l- (aq)» Agl(s) Sulfuric acid cannot be used to acidify the mixture because it contains sulfate ions which would form a precipitate The role of nitric acid is to react with any carbonates present to prevent formation of the precipitate Ag2COg. This would mask the desired observations 2 HN03 + Na2COa 2 NaN03 + H20 + C02 Hydrochloric acid cannot be used to acidify the mixture because it contains chloride ions which would form a precipitate The silver halide precipitates can be treated with ammonia solution to help differentiate between them if the colours look similar: Silver chloride dissolves in dilute ammonia to form a complex ion AgCl(s) + 2NH3(aq) (aq) + Cl- (aq) Colourless solution Silver bromide dissolves in concentrated ammonia to form a complex ion Agar(s) + 2NHg(aq) (aq) + Br- (aq) Colourless solution Silver iodide does not react with ammonia — it is too insoluble. Testing for positive ions (cations) Test for ammonium ion NH', by reaction with warm NaOH(aq) forming NH, NH' +OH' NH3 + Ammonia gas can be identified by its pungent smell or by turning damp red litmus paper blue See flame tests on page 4 for more cation tests N Goalby chemrevise.org 9

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