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A2 chemistry 5.3.2
Edexcel chemistry - transition metals
| Question | Answer |
|---|---|
| What are the rules for filling electron shells? | Aufbau principle - fill subshells from low to high energy. Pauli exclusion principle - each orbital contains a max of 2 electrons with opposite spin. Hund's rule- electron pairing won't take place in degenerate orbitals until each orbital is singly filled |
| What is the order of filling/ losing the 3d and 4s electrons and why? | 4s are filled before 3d, excluding Cr and Cu. 4s are lost before 3d: losing 4s electrons makes ion smaller than if it had lost 3d electrons, so ΔHLE and ΔHhyd are more exothermic |
| What is the pattern in successive ionisation energies of d-block elements? | Not a significant difference in the amount of energy required to remove final 4s and first 3d electron (similar energy levels), 1st large jump is removal of first 3p electron (closer to nucleus, less shielding, greater force of attraction) |
| What is the pattern in ionisation energies along d-block period? | 1st/ 2nd: IE doesn't increase by much - nuclear charge increases but not effective nuclear charge, more shielding. 3rd: IE increases: nuclear charge increases, minimal change in sheilding, greater force of attraction |
| What is the pattern in ionisation energies down a group? | IE decreases: distance between nucleus and outer electron increases, increase in shielding, weaker force of attraction - outweighs increase in nuclear charge |
| Why may there be fluctuations in successive ionisation energies/ along a period? | Pairing of electrons: requires less energy ot remove a paired electron (experiences repulsion) than an unpaired electron with parallel spin |
| Define: transition metal | An element that forms at least 1 stable ion with a partially-filled d-subshell |
| List the properties of transition metals | Variable oxidation states, coloured compounds, form complex ions, good catalysts, high tensile strength, high melting points, high density |
| Why do transition metals have high tensile strength/ high melting points? | Lots of valence electrons (4s and 3d) so sea of delocalised electrons is electron-dense, strong metallic bonding requires a lot of energy to overcome |
| What is the most common oxidation state for transition metals and why? | +2, loss of 2 4s electrons |
| Why does the +2 oxidation state become more stable across the period? | 3rd IE increases (nuclear charge increases, minimal change in shielding, force of attraction increases) so energy required to remove 3rd electron isn't compensated for by ΔHLE/ ΔHhyd |
| Which element has the most number of possible oxidation states and why? | Mn: contains 5 unpaired electrons, can promote a 4s electron to an empty 4p orbital which allows 7 bonds to form |
| What types of bonding exist for different oxidation states? | +2/+3 = ionic. > +3 = covalent, often as a polyatomic anion, bonded to O/F (most electronegative elements) |
| Why do transition metals show variable oxidation states? | 4s/3d subshells have similar energy;similar stability is gained by promoting/ losing different numbers of e-s. No large jumps in successive IEs: EXTRA ENERGY ΔHLE/ ΔHhyd compensates for energy reqd to remove e- , covalent bonding compensates for energy re |
| Why does Zn not show variable oxidation states? | Zn2+ has a stable d10 configuration: energy required to promote or remove one of these electrons wouldn't be recovered by bond formation |
| Define: complex ion, coordination number | Central metal ion surrounded by ligands (lewis base) which use lone pair to form dative covalent bonds with ion. coordination number: number of dative covalent bonds between central metal ion and all ligands |
| Define dentate nature of ligand + egs | Number of dative covalent bonds ligand forms with central metal ion. Mono e.g. Cl-, CN-, H2O. Bi: 2 lone pairs in correct location, 2 dative covalent bonds, folds into a ring e.g. 1,2-diaminoethane (en), ethanedioate ion. Poly e.g. EDTA4- (6 bonds) |
| Why are transition metals particularly likely to form complex ion? | Small, high charge density, polarising, attract ligand electrons |
| What are the names of metals for anionic complexes? | Cuprate, ferrate, vanadate, aluminate, cobaltate, nickelate, chromate, manganate |
| How do we predict the shapes of complex ion? | VSEPR: electron pairs repel each other to a position of maximum separation and minimum repulsion |
| If the coordination number is 4, how do you know which shape the complex will be? | Central metal ion has a full d-subshell: tetrahedral e.g. [CrCl4]- Central metal ion has a partially-filled d-subshell: square planar e.g. [Pt(NH3)2Cl2] |
| How can you achieve stereoisomerism in metal complexes? | Optical isomerism: octahedral complex with 3 bidentate ligands/ 2 bidentate + 2 monodentate ligands. Geometric isomerism: square planar complex - cis (same isomer 90˚ apart), trans (same isomer 180˚ apart) |
| Draw + name 5 d-orbitals | 3 lie between axis (dxy, dxz, dyz), 2 lie along axis (dx2-y2, dz2) |
| Describe d-orbital splitting in an octahedral complex | In free metal ion, d-orbitals are degenerate. 6 ligands approach ion along axis, greater repulsion between lone pair in ligand and electrons in axial d-orbitals, energy of electrons in axial d-orbtials increases/ in orbitals between axis decreases, Δo |
| Describe d-orbital splitting in a tetrahedral complex | In free metal ion, d-orbitals are degenerate. 4 ligands approach ion between axis, greater repulsion between lone pair in ligand and electrons in dxy, dxz, dyz orbitals, energy of electrons in d-orbtials between axis increases/ axial decreases, Δt |
| Why do complex ions have a colour? | d-orbitals are split by ligands. Shine white visible light on complex. Wavelength of visible light is absorbed (energy of photon=Δ), colour absorbed/ observed diametrically opposite.D-D: e- promoted to high-energy d-orbital, excited state, unstable, e- dr |
| What factors affect colour? | Strength of ligand: strong-field ligands interact more strongly with d-orbitals in ion, greater splitting (CN->RNH2>NH3>H2O>OH->F->Cl->SO42-, poly>bi>mono dentate. Oxidation state/ charge density of central metal ion: higher charge density... |
| Cont. | = more polarising, attracts electron density more, greater splitting. More splitting in octahedral than tetrahedral complexes. More splitting = more energy required to promote e- to high energy d-orbital, colour of light absorbed (+observed) shifts toward |
| What are the colours of solid and aqueous chromate (VI)? | Solid: yellow Aqueous: yellow |
| What are the colours of solid and aqueous iron (II)? | Solid: pale green. Aqueous: pale green. |
| What are the colours of solid and aqueous nickel (II)? | Solid: green Aqueous: green |
| What are the colours of solid and aqueous zinc (II)? | Solid: white Aqueous: colourless |
| What are the colours of solid and aqueous chromium (III)? | Solid: green Aqueous: green |
| What are the colours of solid and aqueous Mn (II)? | Solid: pink Aqueous: pale pink |
| What are the colours of solid and aqueous iron (III)? | Solid: brown Aqueous: brown/ yellow |
| What are the colours of solid and aqueous copper (II)? | Solid: blue/ green Aqueous: blue |
| What are the colours of solid and aqueous dichromate (VI)? | Solid: orange Aqueous: orange |
| What are the colours of solid and aqueous manganate (VII)? | Solid: purple Aqueous: purple |
| What is a heterogeneous catalyst? | Different phase to reactants, gaseous reactants adsorb onto solid catalyst active sites by bonding weakly to accessible 3d/4s electrons, increases rate: brings reactants into close contact+ orientation, weakens reactant intramolecular forces |
| Examples: | Palladium, platinium, rhodium in catalytic converters: convert CO, NO and unburnt hydrocarbons via surface activity. Vanadium catalyses contact process (H2SO4 production) through variable oxidation states: V2O5 + SO2 --> 2VO2 + SO3. 2VO2 + 1/2O2 --> V2O5 |
| What is a homogeneous catalyst? | Same phase as reactants: catalyst reacts with reactant to form an intermediate compound, reacts with other reactant to reform catalyst. Transition metals show VARIABLE OXIDATION STATES so can READILY REDUCE/OXIDISE OTHER SPECIES, energy differences betwee |
| What is deprotonation? | Where a base removes a proton from a HYDRATED d-block metal ION, forms a METAL HYDROXIDE! |
| What affects the extent of deprotonation? | Charge density/ oxidation state of metal ion (higher charge density attracts ELECTRON DENISTY towards it and AWAY FROM O-H BOND, LIGAND H is more delta + and more readily lost as an H+ ion), base strength: NH3 and OH- > H2O |
| Describe the properties of H2O as a base for deprotonation | Weak base, only deprotonates hydrated d-block metal compexes with oxidation state of +3 or higher, doesn't form a precipitate, further deprotonation is unlikely, forms H3O+, higher charge density = more acidic |
| Describe the properties of OH- as a base for deprotonation | Strong base, small amount of NaOH deprotonates all hydrated complexes, forms ppt + H2O, amphoteric hydroxides (chromium, zinc) act as acids, undergo further deprotonationin excess OH- so ppt dissolves, [Cr(OH[6]3-, [Zn(OH)4]2- |
| Effect of small amount/ excess OH-: cobalt | Small amount: blue ppt, turns pink on standing Excess: no reaction |
| Effect of small amount/ excess OH-: copper | Small amount: blue ppt Excess: no reaction |
| Effect of small amount/ excess OH-: zinc | Small amount: white ppt Excess: ppt dissolves to form a colourless solution |
| Effect of small amount/ excess OH-: nickel | Small amount: green ppt Excess: no reaction |
| Effect of small amount/ excess OH-: manganese (II) | Small amount: off-white ppt, darkens on exposure to air as it is oxidised to MnO4 Excess: no reaction |
| Effect of small amount/ excess OH-: iron (III) | Small amount: brown ppt Excess: no reaction |
| Effect of small amount/ excess OH-: Mg2+/Ca2+/Sr2+/Ba2+ | Small amount: white ppt Excess: solubility INCREASES down group |
| Effect of small amount/ excess OH-: iron (II) | Small amount: green ppt, turns brown in air as it is oxidised to iron (III) hydroxide Excess: no reaction |
| Effect of small amount/ excess OH-: chromium (III) | Small amount: green ppt Excess: ppt dissolves ot form a green solution |
| Describe the properties of NH3 as a base for deprotonation | Strong base, deprotonates all hydrated d-block metal complexes, ppt + NH4+, hydroxides of copper (II), chromium (III), cobalt (II), zinc (II), nickel (II), silver (I) dissolve, ligand exchange, 4 ammine ligands |
| Effect of small amount/ excess NH3: iron (II) | Small amount of NH3: green ppt, turns brown on exposure to air as oxidised to iron (III) hydroxide. Excess NH3: no reaction |
| Effect of small amount/ excess NH3: zinc (II) | Small amount of NH3: white ppt. Excess NH3: ppt dissolves to form a colourless solution |
| Effect of small amount/ excess NH3: cobalt (II) | Small amount of NH3: blue ppt. Excess NH3: BROWN solution. |
| Effect of small amount/ excess NH3: manganese (II) | Small amount of NH3: off-white ppt., darkens on exposure to air as it is oxidised to MnO2. Excess NH3: no reaction |
| Effect of small amount/ excess NH3: chromium (III) | Small amount of NH3: green ppt. Excess NH3: ppt. dissolves slowly to form a green solution |
| Effect of small amount/ excess NH3: nickel (II) | Small amount of NH3: green ppt. Excess NH3: pale blue solution |
| Effect of small amount/ excess NH3: iron (III) | Small amount of NH3: red-brown ppt. Excess NH3: no reaction |
| Effect of small amount/ excess NH3: copper (II) | Small amount of NH3: blue ppt. Excess NH3: deep blue |
| Effect of small amount/ excess NH3: silver (I) | Small amount of NH3: brown ppt., not often seen as ammonia complex forms so readily Excess NH3: colourless |
| What are the colours of the vanadium oxidation states? | V2+ = violet. V3+ = green. (VO)2+ = blue. (VO2)+ = yellow. VO3- = colourless |
| How would you increase the oxidation state of vanadium? | oxidise with HNO3/ O2 in air |
| How would you decrease the oxidation state of vanadium? | Dissolve NH4VO3 in water, VO3- + H2O <-> (VO2)+ + 2OH-. Add HCl to drive position of equilibrium to the right, yellow. Reducing agent must have E* lower than E* of vanadium compound but possibly higher than vanadium product to avoid further reduction |
| What is the colour of [Cr(H2O)6]2+, [Cr2(CH3COO)4(H2O)2], [Cr(H2O)6]3+, [Cr(H2O)4(OH)2], [Cr(NH3)4(H2O)2]3+, [Cr(OH)6]3-, Cr2O72-, CrO42-? | Blue, red, violet, green, green, green, orange, yellow |
| How would you stabilise Cr(II) + method + equation? | Chromium ethanoate complex (red): 2[Cr(H2O)6]2+ + CH3COO- --> [Cr2(H2O)2(CH3COO)4] (s) + 10H2O. Mix K2Cr2O7, Zn, HCl in a flask, connected to test tube of CH3COONa via delivery tube, screw cap seal open. Orange --> green --> blue, close screw cap seal... |
| Cont | Pressure from H2 build up forces Cr2+ solution out of flask into CH3COONa solution, ligand exchange, red ppt |
| What are the colours of all Cu+ complexes? | White solids, colourless solutions - electron configuration is [Ar]3d104s0, although there is d-orbital splitting there are no empty d-orbitals to promote electrons into so no d-d transitions, doesn't absorb visible light. Except: Cu2O = brick red ppt |
| How do you convert CuCl to [CuCl2]- and back again? | forwards: add excess HCl. backwards: add ice cold water |
| How do you convert CuI and CuCl into [Cu(NH3)2]+? | Add NH3 |
| Describe the stability of Cu(I), when does it exist, what happens normally in solution? | Unstable, only exists in complexes, precipitates, high temperatures. Disproportionates: 2Cu+ --> Cu2+ + Cu |
| How do you convert [CuCl2]- to [CuCl4]2-? What kind of reaction is this + colour change? What is the reverse? | Boil with HCl, oxidation from Cu(1) to Cu(II), colourless --> yellow. Reverse: reduce by boiling with Cu |
| What colour is CuCl2? What oxidation state is Cu in? How do you convert it to [CuCl4]-? | Brown solid, but absorbs moisture from air --> blue. Boil with HCl to form [CuCl4]- |
| What colour is CuSO4 and what complex ion does it contain? | Blue, [Cu(H2O)6]2+ |
| What happens when you add excess HCl to [Cu(H2O)6]2+? | ligand exchange, reversible reaction: [Cu(H2O)6]2+ + 4Cl- <-> [CuCl4]- + 6H2O. blue --> yellow, so appears green. Drive position of equilibrium by adding/ removing acid |
| How is [Cu(H2O)6]2+ deprotonated? | NH3/ OH-, forms [Cu(H2O)4(OH)2] (s), blue |
| What happens upon adding excess NH3? | Ligand exchange reaction, [Cu(NH3)4(H2O)2]2+, blue ppt dissolves to form a deep blue solution |
| What happens upon heating [Cu(H2O)4(OH)2] (s)? | Decomposition, forms black solid CuO |
| How is solid, anhydrous CuSO4 formed + what colour is it + why? | recrystallize [Cu(H2O)6]2+ --> [Cu(H2O)4].SO4.H2O (blue solid), then heat to remove ligands, no d-orbital splitting, white |
| How is Cu (0) oxidised to Cu (II)? | Add Cl2: CuCl2 (brown/ blue). Add conc H2SO4: CuSO4 (blue). Add HNO3: Cu(NO3)2 (blue) |
| How is CuI formed? | Mix CuSO4 and KI: Cu2+ + 2I- --> CuI + 0.5 I2. White solid |
| List some uses of transition metals | Catalysts: iron (haber process), vanadium (V) oxide (contact process), nickel (hydrogenation of alkenes/ unsaturated triglycerides), catalytic converters in cars, ethanoic acid production. Cis-platin = chemotherapy drug. Polychromic/ photochromic glasses |
| What are the methods of producing ethanoic acid? | Oxidation of ethanol (low atom economy) or CO + methanol (addition reaction) |
| What different catalysts have been used to make ethanoic acid? | BASF: cobalt/ iodide catalyst, 300˚c, 700atm. Monsanto: rhodium/iodide catalyst, 180˚c, 30atm - rhodium is expensive, catalyses side reactions,remove impurities. Cativa: iridium/iodide catalyst - iridium is cheaper, fewer side reactions, 99% yield, faster |
| What are some uses of ethanoic acid? | Vinegar - preservative. Esters e.g. ethyl ethanoate - solvent for decaffeinating tea+ coffee, nail varnish, pear flavouring. Ethenyl ethanoate monomer -> poly (ethenyl ethanoate) - PVA glue. Poly (ethenyl ethanoate) -> poly(ethenol) - liquitabs, laundry |
| How does cis-platin behave as a drug? | Cis-[Pt(NH3)2Cl2] - chemotherapy drug - binds to guanine via ligand exchange reaction, nitrogen atom replaces chloride ion, so cell cant divide so immune system destroys cell |
| How do polychromic/ photochromic glasses work? | Contain small amounts of colourless AgCl, CuCl. Cu+ + Ag+ <-> Cu2+ + Ag. In UV, Cu+ reduces Ag+ to Ag, darkens glass, reducing UV exposure + light intensity. Without UV, backwards reaction, glasses are transparent |
| How would you prepare a water-soluble inorganic salt? | Make salt, filter through filter paper in a funnel to remove unreacted reactant, heat with Bunsen burner to remove some excess water, when crystals start to form remove heat and allow to crystallise + cool in an evaporating basin |
| Why may the % yield be above or below 100? | Above: not enough water evaporates. Below: incomplete reaction, competing side reactions, reactants/ products lost in practical preparation e.g. left in crucible |
| How would you prepare an inorganic salt that is insoluble in water? | Make salt, filter under reduced pressure in a Buchner funnel, discard filtrate, dry solid between filter paper and leave to dry completely |
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