Summary sheets for periodicity , group 2,group 7 and more

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across a period. lonisation energy decreases going down a group. s-block 1s 2A -2s- +35 38 -4s -5s- -65 +75+ emperature/ 3000 2500 2000 1500 1000 500 48 58 Na 11 d-block 68 7888- 18 28 Mg 12 3d 4d 5d 6d Al 13 f-block Si lonization energy 14 15 atomic number Electron affinity P 3A 4A 5A Atomic p-block Nonmetallic character CHA Metallic character boiling point melting point 2p 3p 4p 5p 6p 16 CI 17 BA 1s+ Ar 18 Unit 2: Group II 1. Physical properties Trends in group II, Beryllium is not typical of the group and it is not considered here. Symbol Z Atomic radius (nm) 1st ionisation energy (kJ/mol) 0.160 0.197 0.215 0.218 Mg Ca Sr Ba 12 20 38 56 0 +1 0 M(s) + 2H₂O(l) M(OH)₂(aq) + H₂(g) 2. Reactivity with water Reactivity with water INCREASES down the group. Magnesium reacts slowly with liquid water, but rapidly when heated in the presence of steam. +2 3. Hydroxides and sulphates - solubility Hydroxides M(OH)₂ Varying solubility in water. Solubility INCREASES as you descend the group. pH of the hydroxide in water varies. pH increases as you descend the group. Sulphates - MSO4 Melting point (K) Chemical Mg(OH)₂ Ca(OH)2 BaSO4 4. Application of group II compounds Common name Milk of magnesia Slaked lime Barium meal 650 842 777 727 Colourless solids Solubility DECREASES as you descend the group. Thermal Decompose to form MO(s) and CO₂(g). Thermal stability increases as you descend the group. 738 590 550 503 Applications Treat indigestion, heartburns and wind. Neutralise fields and polluted lakes. Contrast medium for gut X-ray. Density P 1.74 1.54 2.60 3.52 ▬▬▬▬ Unit 3a: Group VII 1. Keywords Mean bond Enthalpy: Displacement: Electronegativity: bond enthalpy (kJ mol-1) lone pairs 300 200 100 0- The average enthalpy change when one mole of a specific bond is broken in a range of different gaseous compounds. A displacement reaction is a type of reaction in which part of one reactant is replaced by another reactant. The power of an atom to attract the electrons in a covalent bond. F-F CI-CI Br-Bri 1-1 As the atoms get smaller, lone pairs on the two atoms gef dose enough together to experience serious repulsion. $4 bonding pair 2. Physical properties Trends in group VII. A number of properties of Fluorine are untypical, this mainly stem from the fact that the mean bond enthalpy of the F-F bond is unexpectedly low. This is due to electron repulsion. Symbol Z F CI Br I F CI Br I 9 17 35 53 Electronegativity F₂ Cl₂ Br₂ 1₂ 3. Physical states The physical state of the halogens are summarised below. Symbol In pure form Pale yellow gas Pale green gas Dark red liquid Grey solid F. no 4.0 3.0 2.8 2.5 no Atomic (covalent) radius (nm) no 0.071 0.099 0.114 0.133 In non-polar solvents (Reacts with solvents) Pale green solution Orange solution Purple solution 4. Oxidising abilities - Displacement reactions The oxidising ability of the halogens decreases down the group. You cannot investigate the oxidising ability of Fluorine in aqueous solution because it reacts with water. CI yes - Melting point (K) 53 8172 266 387 no no Br yes yes Boiling point (K) 85 238 332 457 no In water (Reacts with water) Pale green solution Orange solution Insoluble |. yes yes yes Unit 3b: Reactivity of halide ions and chlorine. 1. Keywords Disproportionation: Precipitate (ppt): Symbol 2. Reducing strength The reducing ability of the halide ions increases down the group. F. Br CI I Br |· Symbol CI a reaction in which a substance is simultaneously oxidized and reduced, giving two different products.. deposited solid formed in a solution. 4. Reaction with silver ions All metal halides (but fluoride) react with silver ion to form an insoluble precipitate. Dilute nitric acid is added before the reaction to get rid of any carbonate or hydroxide impurities. Ag+ (aq) + X (aq) → AgX (s) Observation Atomic radius (nm) White ppt Cream ppt Pale yellow ppt 0.133 0.180 0.195 0.215 Halide salt solubility Dilute NH3 Concentrated NH3 Insoluble in NH3 3. Reaction with concentrated sulphuric acid Solid halides react with concentrated sulphuric acid giving different products based on their reducing powers. Chemical reactions. Products CI- Br |· Reaction A Reaction B Reaction C Reaction D Reaction A Reaction B MeHSO4 ✓ ✓ 5. Reactivity of chlorine Reactivity with water: Reactivity with water in sunlight: SO₂ X ✓ NaCIO(s) + H₂O(l) ✓ Reactivity with alkali: Summary table Reaction C 2Cl₂(g) + 2H₂O(l) → 4HCI(g) + O₂ (g) Alternative chlorination of swimming pools: Cl₂(g) + H₂O(l) HCl(aq) + HCIO(aq) "chlorine water" S X MeX(s) + H₂SO4(1)→ MeHSO4 (s) + HX(g) 2H+ (aq) + 2X* (aq) + H₂SO4 (I) → SO₂ (g) + 2H₂O (1) +X₂ (1) 6H+ (aq) + 6X (aq) + H₂SO4 (I) → S(s) + 4H₂O (1) + 3X₂ (s) 8H+ (aq) + 8X (aq) + H₂SO4 (1)→ H₂S(g) + 4H₂O (1) + 4X₂ (s) X Na+(aq) + OH (aq) + HCIO(aq) Reaction D Cl₂(g) + 2NaOH(aq) → Cl(aq) + NaCIO(aq) + H₂O(1) H₂S X X . . Observations Steamy fumes (HCI) Ppt (MeHSO4) Steamy fumes (HBr) Brown fumes (Br₂) Pungent gas (SO₂) Steamy fumes (HI) Black ppt (1₂) Rotten egg smell (H₂S) Yellow ppt (S) Types of reaction Acid-base Acid-base and redox Acid-base and redox Disproportionation reaction Goes from pale green to colourless Water is kept slightly acidic Disproportionation reaction Unit 4: Period 3 1. Keywords Amphoteric: 4. Reactivity of Period 3 oxides with acids/bases Na₂O MgO Al₂O3 melting points/C SiO₂ P4010 SO₂ 5. Reactivity of Period 3 oxides with water Na₂O(s) + H₂O(l) → 2NaOH(aq) MgO(s) + H₂O(l) → Mg(OH)₂(aq) Insoluble no reaction Na₂O MgO Al₂O3 SiO₂ P4010 SO₂ SO3 3500 3000 2500 2000 able to react both as a base and as an acid. Na₂O(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l) Al₂O3(s) + 6HCl(aq) → 2AlCl3(aq) + 3H₂O(l) Al₂O3(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)4(aq) SiO₂ (s) + 2NaOH(aq) → 2Na₂SiO3(aq) + H₂O(l) P4010(s) + 12NaOH(aq) → 4Na3PO4(aq) + 6H₂O(l) SO₂(g) + 2NaOH(aq) → Na₂SO3(aq) + H₂O(l) 1500 E 1000 Na₂O 500 0 Mgo P4010(s) + 6H₂O (1)→ 4H3PO4(aq) SO₂(g) + H₂O(l) → H₂SO3(aq) SO3(g) + H₂O(l) → H₂SO4(aq) Al₂O3 Insoluble SiO₂ oxides of Na to S P4010 SO₂ pH = 14 pH = 9 pH = 7 pH = 7 pH = 1-2 pH = 2-3 pH = 0-1 Alkali Alkali Amphoteric Weak acid Acid Acid Al₂O3/2345K SiO₂/1883K P4010 /573K SO3/290K SO₂/200K 2. Reactivity HO Reactivity with water (only Na, Mg, Cl₂): 3. Period 3 oxides melting points Na₂0/1548 K Mg0/ 3125K 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) Mg(s) + 2H₂O(l) →Mg(OH)₂(aq) + H₂(g) Mg(s) + H₂O(g) →MgO(s) + H₂(g) Cl₂(g) + H₂O(l) HCl(aq) + HCIO(aq) Reactivity with oxygen: 4Na(s) + O₂(g) → 2Na₂O(s) 2Mg(s) + O₂(g) → 2MgO(s) 4Al(s) + 30₂(g) → 2Al₂O3(s) Si(s) + O₂(g) → SiO₂ (s) 4P(s) + 50₂(g) → P4010(S) S(s) + O₂(g) → SO₂(g) lonic structure = high melting points. OH Sulfuric (IV) acid - H₂SO3 Bright yellow flame forming white sodium oxide Bright white flame forming white magnesium oxide. Bright white flame in a gas jar forming white Al₂O3. Reacts only when heated. Red P reacts when heated. White P reacts spontaneously + white smoke. Reacts when heated and lowered in a gas jar forming a colourless gas. lonic structure = high melting points. Smaller and more highly charged than Na giving a stronger attraction to the oxygen ions O=S=O Faster with steam lonic structure with covalent character. Very high charge density on the small Al Giant covalent structure. Strong covalent bonds Si-O. Simple molecular compounds = low melting points. Simple molecular compounds/weak intermolecular forces = low melting points Simple molecular compounds/weak intermolecular forces = low melting points HO-S-OH Sulfuric(VI) acid - H₂SO4 OH O=P-OH OH Phosphoric (V) acid - H3PO4 Unit 5a: Transition metals 1. Keywords Colorimetry: Complex ions: Coordination number: scientific technique that is used to determine the concentration of coloured compounds using its absorbance . metal ion surrounded by dative covalently bonded ligands. 2. Classification A transition metal is an element that forms at least one stable ion with a partially full d-shell. . is the number of dative covalent bonds that are formed with central transition metal ion. Scandium and Zinc are not considered to be transition metals. Scandium forms only a 3+ ion [Ar] 4sº 3dº Zinc forms only a 2+ ion [Ar] 4sº 3d¹⁰ Copper is a transition metal because its +2 ion has an incomplete d orbital. [Ar] 4sº 3d⁹ Chemical properties Physical properties Transition metals have incomplete d sub- shells. When the fourth period metals form ions the 4s electrons are the first to be lost then 3d electrons. Form coloured ions Form complex ions Act as good catalysts Variable oxidation states. . Shiny Strong Hard Good conductors 3. Coloured ion formation lons with d¹0 or dº are colourless, those partly filled tend to be coloured. Energy is absorbed when an electron is promoted to a higher energy level. The frequency of light is proportional to the energy difference. Electrons in lower 3d sublevels absorb energy from visible light which promotes them to a higher 3d sublevel. The rest of the frequencies are transmitted hence the complementary colour is observed. Factors affecting colour Size & type of ligand Oxidation state Complex shape . . ● Coordination number Strength of metal-ligand bond The size of the energy gap can be calculated if we know the frequency (v) or wavelength (A) of the light absorbed: AE = hv=hc λ AE = energy absorbed (J) c = speed of light (3.00 x 108 m s-¹) h = Planck's constant (6.63 x 10-34 J s) wavelength of light absorbed (m) λ= v = frequency of light absorbed (Hz) Colorimetry: we can then determine the concentration by comparing our sample against known concentrations of the same metal ion / ligand. 1. Make up solutions of known concentration. 2. Measure absorption of each known solution 3. Plot a graph of absorption vs concentration 4. Measure absorption of unknown and compare on the graph to give a concentration reading Compounds of Vanadium Aqueous ion VO₂ Oxidation +5 state Colour Aqueous ion Oxidation state Colour Compounds of Chromium 1.6 1.4 1.2 1.0 0.8 0.6 0.4 0.2 yellow 0.8.0 Ni²+(g) ● Cr₂O²( +6 Orange 1 VO² 0.5 +4 blue cro +6 Yellow V. +3 1.0 Concentration green Cr³ +3 Green V². TJT Ni²+(ag) +2 violet Cr²- +2 the d orbitals split into two groups with different energies in an aqueous ion Blue absorption of light at a particular frequency can promote an electron to a higher level 2.0 Unit 5b: Transition metals 1. Keywords Catalyst: Heterogeneous catalyst: Homogeneous catalyst: Ligand: Redox potential (electrode potential): Energy X, Y concentration of reactant Autocatalysis is a substance that increases the rate of a chemical reaction by providing an alternative route with a lower activation energy. The catalyst unchanged by the reaction. is in a different phase from the reactants. is in the same phase as the reactants. is an atom, ion or molecule that donates a pair of electrons to a central transition metal. is a measure of the intensity of its oxidising power, > affinity for electrons = more positive value. Reaction Progress The reaction starts slowly time Ea (no catalyst) is Ea (with catalyst) AG and then speeds up before slowing down again and eventually stopping. 2. Variable oxidation state The table shows the oxidation states of the transition metals. (red indicates the most common oxidation states found for that element). Not all of them are stable, only the lower oxidation states exist as stable ions (i.e. Mn 7+ exists only when in a covalent bond). The redox potential for a transition metal ion influenced by pH and by the ligand. In general, oxidation is favoured by alkaline conditions (since there is a higher tendency to form negative ions) and reduction is favoured by acidic conditions. +3 . Sc Heterogeneous catalyst: The reaction occurs at active sites on the surface. Homogeneous catalyst: The reaction proceeds through an intermediate species. Autocatalysis: where the catalyst is a product of the reaction. As a result the rate of reaction increases over time as more product, and therefore catalyst, is produced. +4 Heterogeneous catalysts can become poisoned by impurities blocking the active sites, reducing efficiency. Catalyst can also be lost from the support. This has a cost implication. +5 +4 +3 +3 +3 Ti V +6 +6 +5 +2 +2 +2 +4 Reactants Adsorbed reactants +7 +5 +4 +3 +2 +6 +5 +5 +4 +1 +1 +1 +1 +1 +1 +4 +4 +3 +3 Reactants +2 +2 5. Catalytic activity Transition metals are good as catalysts either due to a change in oxidation state or they adsorb other substances onto the surface. A partially filled d orbital can be used to form bonds with adsorbed reactants. Types of catalysts: Homogeneous catalysis Intermediate Cr Mn Fe Co Ni Cu Zn Heterogeneous catalysis +3 +3 Intermediate +2 Catalyst surface +2 +2 +1 +1 Product -Catalyst Product Adsorbed product Unit 6: Complex ions 1. Keywords Chelate: Coordinate (dative) bond : Coordination number: Ligand: Optical isomerism: Optical isomers These two ligands are in the vertical plane, but twisted at 90° to each other. compound containing a ligand bonded to a central metal atom at two or more points. H3N... H3N a covalent bond where both electrons in the shared pair come from the same atom Geometrical isomers CN NH3 is the number of dative covalent bonds that are formed with central transition metal ion. is an atom, ion or molecule that donates a pair of electrons to a central transition metal. when two or more forms of a compound with the same structure are mirror images of each other and typically differ in optical activity. Ionisation isomers NH3 CN NH3 SO4 Br N N N This ligand is in the horizontal plane. mirror H3N H3N 'N N NH3 NH3 CN CN NH3 Br OSO3 2. Complex ion formation All transition metals for co-ordinate bonds by accepting electron pairs from other ions called ligands. Different ligands form different strength bonds. Monodentateligands (single co- ordinate bond): H₂O (Aqua) NH3 (Ammino] CN . OH (Hydroxi CI- . small Coordination number Shape 2 3. Complex ion shapes The coordination number dictates the shape of the complex ions. large The charge of the complex ion depends on th charge of the transition metal and on the charge and number of the ligand. Linear Bidentate (two co-ordinate bonds): H₂NCH₂CH₂NH₂ ethane 1,2 diamine 4 . Tetrahedral or square planar (en) C₂O4²- ethanedioate (oxalate) Benzene 1,2 diol 2molecules :00CCH₂ :00CCH₂ 6 Octahedral NCH₂CH₂N: CN Types of isomerism can occur in complex ions: Geometrical isomerism - when two molecules have ligands in different position in space. • Optical isomerism - when two or more bidentate ligands are in a complex. Ni CH₂COO: the EDTA4-lon CH₂COO: CN _7 molecules CN 2- molecules generally Multidentate (can form many co- ordinate bonds): Haem EDTA 4 which can form 6 co- ordinate bonds Chi Cu lonisation isomerism - when a ligand has exchanged places with an anion or neutral molecule that was originally outside the coordination complex. the [Cu(EDTA)?- lon H₂N H₂N 4. Chelate effect When bidentate or multidentate ligands take the place of monodentate, there are more products than reactants hence entropy of system increases, so products are thermodynamically more stable. The enthalpy change of these reaction is very small but the increase in entropy makes the reaction possible. [Cu(H₂O)]²+ (aq) + EDTA4 (aq) ) [CuEDTA]2 (aq) + 6H₂O(l) NH3 Co" NH3 NH₂ NH3 3+ (CI) 3 Gibbs free-energy equation ΔG = ΔΗ - ΤΔ Entropy increases = favourable reaction Unit 7a: Reactivity of complex ions 1. Keywords Amphoteric Hydrolisis H₂O H₂O H₂O H₂O OH₂ means can behave as an acid and as a base. the chemical breakdown of a compound due to reaction with water. These electron pairs are all being pulled away from the oxygens towards the 3+ ion. That causes the electron pairs in the O-H bonds to be pulled even doser to the oxygen than normal. That makes the hydrogen atoms even more positive than they normally are when they are attached to oxygen. 2. Acidity of aqua ion In general the acidity of M³+ will be greater than that of M²+; i.e.Fe³+ has a greater charge density than Fe²+, there are more positive charges on the same ion and the ion is smaller. So the Fe³+ aqua ion (pka 2.2) is more acidic than the Fe2+ (pKa 5.9) aqua ion. [M(H₂O)]³+ (aq) + OH (aq) -> [M(H₂O),(OH)]²+ (aq) + H₂O(l) Reaction with OH (same with NH3 M³+ aqua ion M²+ aqua ion Reaction with CO3²- M³+ aqua ion M²+ aqua ion 2. Aluminium [M(H₂O), (OH)]²+ (aq) + OH (aq) -> [M(H₂O)4(OH)₂]+ (aq) + H₂O(l) [M(H₂O)4(OH)₂]+ (aq) + OH (aq) -> [M(H₂O)3(OH)3] (s) + H₂O(l) [M(H₂O)]²+ (aq) + OH (aq) -> [M(H₂O),(OH)]+ (aq) + H₂O(l) [M(H₂O),(OH)]+ (aq) + OH (aq) -> M(H₂O)5(OH)₂ (s) + H₂O(l) 2[M(H₂O)]³+ (aq) + 3CO3²-(aq) -> 2[M(H₂O)(OH)3](s) + 3H₂O (I) + 3CO₂(g) Acting as a base. [M(H₂O)]²+ (aq) + CO3² (aq) -> MCO3(s) + 6H₂O (1) CO3²- does not remove the proton from the complex, it is not a strong enough base. In general carbonates of M³+ do not exist, M²+ ones do. Aluminium hydroxide is the most often quoted example of an amphoteric hydroxide. Acting as an acid [Al(H₂O)]³+ (aq) + OH(aq) → [Al(H₂O)5OH]²+ (aq) + H₂O (1) [AI(H₂O), OH]²+ (aq) + OH (aq) → [AI(H₂O)4(OH)₂ ]+ (aq) + H₂O (1) [AI(H₂O)4(OH)₂ ]+ (aq) + OH(aq) → [AI(H₂O)3(OH)3] (s) + H₂O (1) This ion will dissolve. What you see occurring is the precipitation of neutral aluminium hydroxide and then the re-dissolving of it. The aluminium hydroxide has acted as an acid. [AI(H₂O)3(OH)3](s) + OH(aq) → [Al(OH)4] (aq) + 3H₂O(l) [AI(H₂O)3(OH)3] (s) + 3HCI (aq) → [Al(H₂O)]³+ (aq) + 3Cl- (aq) [AI(H₂O)3(OH)3](s) + H+(aq) → [AI(H₂O)4 (OH)₂]+ (aq) [AI(H₂O)4 (OH)₂]+ (aq) + H+(aq) → [AI(H₂O), (OH)]²+ (aq) [AI(H₂O), (OH)]2+ (aq) + H+(aq) → [AI(H₂O)]³+ (aq) Unit 7b: Reactivity of complex ions 1. Ligand exchange Ligands can be swapped around, this is called ligand exchange. It usually results in a colour change. This happens when: The new ligand can form stronger bonds with the metal ion The new ligand is more concentrated with a neutral ligand Complete substitution with a neutral ligand Partial substitution (Cu²+) with other charged ligands (Cu²+) with a multidentate ligand Metal aqua-ion [Fe(H₂O)]2+ (aq) [Fe(H₂O)6]³+ (aq) Yellow [Cu(H₂O)6]²+ (aq) [Al(H₂O)6]³+ (aq) Colourless With OH or NH3(aq) [Fe(H₂O)4(OH)₂] (s) [Fe(H₂O)3(OH)3] (s) Red-Brown [Cu(H₂O)4(OH)₂] (s) [Al(H₂O)3(OH)3] (s) White With excess OH- (aq) Ammonia first will acts as a base and then as a ligand. the coordination number doesn't change and the shape is still octahedral but it is now a distorted octahedral, this is because water has longer bonds with Cu since it has less affinity. No change, precipitate remains No change, precipitate remains No change, precipitate remains [Al(OH)4] (aq) No change, precipitate remains Replacing water as a ligand [M(H₂O)]²+ (aq) + NH3(aq) [M(NH3)6]²+ (aq) + 6H₂O (1) the coordination number doesn't change and neither does the shape. [Cu(H₂O)6]²+ (aq) + 4NH3(aq) [Cu(NH3)4(H₂O)₂]²+ (aq) + 4H₂O(l [Cu(H₂O)]²+ (aq) + 4Cl (aq) [CuCl4]² (aq) + 6H₂O (1) Change in the coordination number and in the shape See chelate effect Unit 6 Change in the coordination number and in the shape forming more stable complex. With excess NH3(aq) No change, precipitate remains [Cu(NH3)4(H₂O)₂]²+ (aq) Deep-blue No change, precipitate remains With Na₂CO3(aq) FeCO3(s) [Fe(H₂O)3(OH)3] (S) Red-Brown CUCO3(s) Blue-green [Al(H₂O)3(OH)3] (s) White Colours of Vanadium oxidation states VIV III ||