8961106
Description
Flashcards by Kayla Price, updated more than 1 year ago
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Created by Kayla Price
almost 7 years ago
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| Question | Answer |
| Where are transition metals found on the periodic table? | d-block |
| What is a transition metal? | A d-block element that can form one or more stable ions with incompletely filled d-orbitals |
| How many electrons can a d subshell hold? | 10 (has 5 orbitals) |
| How are chromium d-orbitals filled and why? | One electron in each orbital (4s1 3d5) as this is more stable than filling the 4s orbital. |
| Why are 4s orbitals filled before 3d orbitals? | They are at a lower energy level |
| How are copper d-orbitals filled and why? | 4s1 3d10 because this is more stable so a 4s electron is promoted to a higher energy orbital. |
| Why are scandium and zinc not transition metals? | Scandium only forms one ion, Sc3+, which has an empty d-subshell and zinc only forms one ion, Zn2+, which has a full d-subshell. |
| List all the stable ions that Vanadium can form | Vanadium(II) V^2+, Vanadium(III) V^3+, Vanadyl(IV) VO^2+, Vanadate(V) VO2+ |
| What makes an ion likely to form a compound or complex? | The energy given out when the compound or complex forms is greater than the ionisation energy for the ion |
| Describe the pattern for the successive ionisation energies for a transition metal | 4s and 3d orbitals are all at similar energy levels, so it takes a similar amount of energy to remove an electron from them. The is a large increase when all the 3d subshells are removed as there is there is a decrease in shielding |
| Describe the pattern of successive ionisation energies for Chromium | Large increase for second ionisation energy as electron is being removed from a half filled 3d subshell so it is very stable, then smaller increases as repulsion decreases |
| What is a complex ion? | a metal ion surrounded by dative covalently bonded ligands |
| What is a ligand? | an atom, ion or molecule that donates a pair of electrons to a central metal atom or ion |
| What is a monodentate ligand? Give examples | A ligand with one lone pair (H20, NH3, Cl-, OH-) |
| What is a bidentate ligand? Give examples | A ligand with two lone pairs of electrons (1,2-diaminoethane) |
| What is a multidentate ligand? | One with more than two lone pairs of electrons (EDTA4+) |
| How would you calculate the oxidation number of a central metal ion in a complex? | Total oxidation number (the overall charge of a complex ion is its oxidation number) - sum of charges of ligands |
| What is the coordination number? | Number of dative covalent (coordinate) bonds formed with the central metal ion |
| What are the most common coordination numbers? | 6(smaller ligands) and 4(larger ligands) |
| What is the shape of a six-fold coordination complex ion? | octahedral (bond angles 90) |
| What is the shape of a four-fold coordination complex ion? | tetrahedral (bond angles 109.5) occasionally square planar shape |
| How do transition metals form coloured complexes? | When ligands bond to the transition metal ions the 3d orbitals split into two different energy levels. Electrons therefore require energy to jump to higher orbitals (excited states), which they get from visible light). The frequencies of light that aren's absorbed are transmitted as colour. |
| Ti^3+ | Purple |
| Ti^2+ | Violet |
| VO2+ | Yellow |
| VO^2+ | Blue |
| V3+ | Green |
| V2+ | Violet |
| Cr2O7^2- | Orange |
| Cr3+ | Green |
| MnO4^- | Purple |
| MnO4^2- | Green |
| Mn2+ | Pale pink |
| Fe2+ | Pale green |
| Ni2+ | Green |
| Fe3+ | Yellow |
| Co2+ | Pink |
| Cu2+ | Pale blue |
| Give the reduction half equation for VO2^+ and the oxidation states of vanadium | VO2^+ + 2H+ + e- <--> VO^2+ + H2O VO2^+ = +5 VO^2+ = +4 |
| Give the reduction half equation for VO^2+ and the oxidation numbers of vanadium | VO^2+ + 2H+ +e- <--> V^3+ + H2O VO^2+ = +4 V^3+ = +3 |
| Give the reduction half equation for V3+ and the oxidation numbers for vanadium | V3+ + e- <--> V2+ V3+ = +3 V2+ = +2 |
| Give the reduction half equation for V2+ and the oxidation numbers for vanadium | V2+ +e- <--> V V2+ = +2 V = 0 |
| Which oxidation state of vanadium has the most positive electrode potential? | +5 |
| Which oxidation state of vanadium has the highest reducing power? | +2 because it has the most negative electrode potential, so is most likely to oxidise and produce electrons that go on to reduce another species |
| List the chromium oxidation states in order of stability | Most stable +3, +6, +2 Least stable |
| What are the 2 ions chromium form with oxygen? What makes them good oxidising agents? | Chromate(VI) ions CrO4^2- and dichromate(VI) ions Cr2O7^2- Easily reduced to Cr3+ because it is very stable |
| What is the colour of aqueous Cr3+ ions? | Violet but the water ligands are usually substituted by impurities so appears green |
| What is the colour of aqueous dichromate ions? | Orange |
| What is the colour of aqueous chromate ions? | Yellow |
| What is the colour of aqueous Cr2+ ions? | Blue |
| How would you reduce dichromate ions? Give the equation | Add zinc and dilute acid: Cr2O7^2-(aq) + 14H+(aq) + 3Zn(s) --> 3Zn2+(aq) + 2Cr3+(aq) + 7H2O(l) |
| How would you reduce Cr3+ ions? Give the equation | Add zinc in an inert atmosphere (Cr2+ ions very unstable so oxidises in the air) 2Cr3+(aq) +Zn(s) --> Zn2+(aq) + 2Cr2+(aq) |
| How would you oxidise Cr3+ ions? Give the equation | Add hydrogen peroxide in alkaline solution: 2Cr3+(aq) + 10OH-(aq) + 3H2O2(aq) --> 2CrO4^2-(aq) + 8H2O(l) |
| How would you create an equilibrium between chromate ions and dichromate ions? Give the equation | Add acid: 2CrO4^2-(aq)+ 2H+(aq) --> Cr2O7^2-(aq) + H2O(l) |
| How is chromium hydroxide made? Give the equations | By reacting Cr3+ (green solution) with aqueous sodium hydroxide or aqueous ammonia which forms a chromium hydroxide precipitate (grey-green). [Cr(H2O)6]3+(aq) + 3OH-(aq) --> [Cr(OH)3(H2O)3](s) + 3H2O(l) [Cr(H2O)6]3+(aq) + 3NH3(aq) --> [Cr(OH)3(H2O)3](s) + 3NH4+(aq) |
| Give equations to show the amphoteric characteristic of chromium hydroxide | When acid is added the OH- ligands protonate forming a green solution of [Cr(H2O)6]3+: [Cr(OH)3(H2O)3](s) + 3H+(aq) --> [Cr(H2O)6]3+(aq) When excess alkali is added the H2O ligands deprotonate, forming a solution of [Cr(OH)6]3-: [Cr(OH)3(H2O)3](s) + 3OH-(aq) --> [Cr(OH)6]3-(aq) +3H2O(l) |
| What is the difference between a ligand exchange reaction and an acid-base reaction? | In ligand exchange reactions the whole ligand is swapped for a different type of ligand, whereas in acid base reactions the ligands are chemically modified by the addition or removal of H+ |
| Write the equation for adding excess ammonia to chromium hydroxide, name the type of reaction and give the colour change | [Cr(OH)3(H2O)3](s) + 6NH3(aq) --> [Cr(NH3)6]3+(aq) + 3OH-(aq) + 3H2O(l) Grey green to purple |
| Describe the process of forming chromium(II) ethanoate [Cr2(CH3COO)4(H2O)2](s) | Sodium dichromate(VI) is reduced using zinc in acid to form a green solution of Cr3+ ions and then further reduced by the zinc to form a blue solution of Cr2+ ions. Sodium ethanoate is mixed with the Cr2+ solution and a red precipitate of chromium ethanoate forms. |
| What apparatus is used to prepare chromium(II) ethanoate? Describe how the procedure is carried out | 1) HCl is slowly added to a flask containing sodium dichromate and zinc mesh. The zinc will reduce the dichromate ions and react with the acid to produce H2(g), which escapes through the rubber tube. 2) As soon as the solution turns clear blue, the rubber tube is pinched shut so no more H2(g) can escape. 3) The build up of pressure in the flask will force Cr2+ solution through the glass tube into a flask of sodium ethanoate 4) A red precipitate of chromium(II) ethanoate will form 5) The precipitate is filtered and washed with water, then ethanol, then ether (while still keeping the solid in an inert atmosphere) |
| What is a ligand exchange reaction? | When two different ligands undergo substitution |
| What happens when ligands of a similar size undergo a ligand exchange reaction? | Coordination number of the complex ion doesn't change or the shape |
| Give the colour change when aqueous chromium(III) ions undergo a ligand exchange reaction with excess aqueous ammonia | Dark green --> purple |
| What happens when larger charged ligands exchange with small uncharged ligands? | There's a change in coordination number and shape as less of the larger ligands can fit around the central metal ion |
| Give the colour change when chloride ions are added to aqueous copper(II) ions | pale blue --> yellow |
| Give the colour change when chloride ions are added to aqueous cobalt(II) ions | Pale pink --> blue |
| Give the colour change when excess ammonia is added to aqueous copper(II) ions? | Pale blue --> deep blue |
| How does carbon monoxide poisoning happen? | An oxygen or water molecule in haemoglobin can be replaced in a ligand exchange reaction by CO, forming carboxyhaemoglobin. The strong dative bond that forms between the CO and iron means the reaction is irreversible so oxygen can no longer get transported. |
| Give the equations and colour change for the reaction between aqueous copper(II) ions and hydroxide ions/ammonia | pale blue --> blue preciptate |
| Give the equation and colour change for the reaction between copper(II) hydroxide and excess ammonia | Blue precipitate --> deep blue |
| Give the equations and colour change for the reaction between aqueous iron(II) ions and hydroxide ions/ammonia | pale green--> green precipitate (which darkens on standing as the precipitate is oxidised by water and oxygen in air to form iron(III) hydroxide |
| Give the equations and colour change for the reaction between aqueous iron(III) ions and hydroxide ions/ammonia | Yellow --> orange precipitate that darkens on standing |
| Give the equations and colour change for the reaction between aqueous cobalt(II) ions and hydroxide ions/ammonia | Pale pink --> blue precipitate which turns brown on standing |
| Give the equation and colour change for the reaction between cobalt(II) hydroxide and excess ammonia | Blue precipitate --> yellow brown (goes brown on standing as oxidised to [Co(H2O)6]3+) |
| Why are transition metals good catalysts? | Because they can change oxidation state easily - which means they can transfer electrons to speed up reactions |
| How do heterogeneous catalysts work? | By adsorbing reactants at active sites on their surface which allows them to react more easily |
| What happens during the contact process? Include equations | A vanadium(V) oxide catalyst changes its oxidation state reversibly, first reducing to vanadium(IV) oxide as it oxidises sulfur dioxide to sulfur trioxide. Then oxidising back to vanadium(V) oxide by reacting with oxygen. |
| How do homogeneous catalysts work? | By combining with the reactants to form an intermediate, which then reacts to form the product and reform the catalyst. The activation energy needed to form the intermediate is lower than that needed to form the products directly from the reactants. |
| Describe what happens during the reaction between peroxodisulfate ions and iodide ions - include equations | As both ions are negatively charged, they are very unlikely to collide and react so and iron(II) catalyst is used. This gets oxidised by the peroxydisulfate ions to form iron(III) sulfate ions. The intermediate Fe(III) ions then oxidise the I- ions to iodine and the catalyst is regenerated. |
| What is autocatalysis? | When a product acts as a catalyst |
| Describe the reaction between ethandioic acid and permanganate ions - include equations | Mn2+ catalyses the reaction by first reacting with MnO4- to form Mn3+ ions. The Mn3+ ions then react with C2O4^2- to form carbon dioxide and reform the Mn2+ catalyst. |
| How do catalytic converters work? | A platinum-rhodium catalyst is used to convert nitrogen oxides and carbon monoxide to nitrogen and carbon dioxide. It works by adsorbing the reactant molecules onto the surface of the solid, activating the molecules so they react more easily (the bonds between the reactants are weakened making them easier to break and reform), then the product molecules are desorbed from the surface. |
| Why is the strength of adsorption important when finding the most appropriate catalyst? | The catalyst needs to not adsorb molecules so strongly that the reactive sites of the surface of the metal are inactivated or adsorb molecules not strongly enough to weaken bonds enough. |
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