AQA A LEVEL CHEMISTRY PAPER 1 Question and Answers

AQA A LEVEL CHEMISTRY PAPER 1 Question and Answers TOF steps 1) ionisation 2) acceleration 3) ion drift 4) detection 5) analysis Electron impact Sample vaporised and electron gun fires high energy electrons at it which knock off 1 electron from each particle, making them 1+ ions NB- can knock off more than one e or break molecular ion electrospray ionization Sample dissolved in volatile solvent then injected through needle to give fine mist which is attached to positive end of high voltage power supply, particles gain proton NB- Mr of substance is actually one less than shown due to extra H+ Acceleration (TOF) positive ions accelerated using electric field so they all have the same kinetic energy Ion drift (TOF) particles with small mass have larger velocity do ions start to separate with lightest ions reaching detector first Detection (TOF) positive ions hit negatively charged plate and gain an electron which forms a current, the larger the current the higher the abundance Analysis (TOF) -computer uses data to produce mass spectrum which shows mass m / charge z ratio -mr or ar is furthest right peak (small peaks larger than mr are due to isotopes) -may be large peaks at lower mr due to fragmentation Electron spin Property of electron (CW or ACW) Represented by up and down arrows Orbitals Defined regions of space around nucleus where electrons most likely to be found, each orbital holds 2 electrons Hund's Rule Electrons prefer to occupy orbitals on their own and only pair up when no empty or bait ask of same energy are available Electron configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 4d10 exceptions to electron configuration chromium and copper, only take one electron in 4s orbital Why does ionization energy decrease down a group? Atoms get bigger so electrons further away from nucleus, greater shielding Why does ionization energy increase across a period? Atoms get smaller, nuclear charge increases, similar shielding Dip in ionisation energy groups 2-3 Electrons take up higher orbital (s to p) which makes ionisation energy lower as higher orbitals have higher energy Dip in ionisation energy groups 5-6 Electron- electron repulsion in orbital makes electron easier to lose Relative atomic mass The average mass of an atom of an element/ 1/12th of the mass of an atom of carbon-12 Empirical formula The simplest whole number ratio of atoms of each element present in a compound Percentage yield actual yield/theoretical yield x 100 Atom economy (Molecular mass of desired products/ Molecular mass of all products) x 100 Electronegativity The power of an atom to attract electron density in a covalent bond towards itself electronegativity trend increases across a period, decreases down a group Enthalpy change Heat energy change measured at constant pressure Hess's Law The Enthalpy Change for a chemical reaction is independent of the route taken Mean bond enthalpy Energy required to break one mole of covalent bonds between two given atoms, averaged across a range of compounds Activation energy the minimum energy colliding particles must have in order to react dynamic equilibrium Forward and backward reactions take place at same rate so concentrations of products and reactants remain constant Only reached in closed system Le Chatelier's Principle When a system at equilibrium is disturbed, the equilibrium shifts in the direction that reduces the disturbance Haber process conditions - 200 atm : high pressures favour forwards reaction but are costly - 450°C : low temperature favours towards reaction but rate of reaction too slow at low temps - iron catalyst Effect of concentration on Kc no effect Solubility of group 2 hydroxides Solubility increases down the group Solubility if group 2 sulphates Solubility decreases down group Use of Mg Extract titanium from its ore TiCl4 + 2Mg - Ti + 2MgCl2 Use of Mg(OH)2 Medicine as cure for indigestion Use of Ca(OH)2 To neutralise acidic soils Use of CaO and CaCO3 Flue gas desulphurisation Use of BaSO4 Used as a tracer for x-rays Important that it is insoluble as it is safe to ingest Halogen bond energy Energy if halogen- halogen bond decreases down group due to larger atomic nuclei NB fluorine is an exception as radii so small that there is repulsion between lone pairs Cl- with H2SO4 Produces NaHSO4 and HCl Not redox as in change in oxidation state because Cl- not strong enough reducing agent Test for HCl with ammonia solution, white cloud produced Br- with H2SO4 Produces SO2 S goes from +6 to +4 Test for SO2 if potassium dichromate goes from orange to green I- with H2SO4 Produces H2S S goes from +6 to -2 Test for H2S with lead ethanoate which goes from white to black Bronsted-Lowry acid proton donor Bronstead-Lowry base proton acceptor Lewis acid electron pair acceptor Lewis base electron pair donor Partial pressure mole fraction x total pressure mole fraction Moles of one gas/ total miles of all gases Effect of changing pressure on Kp No effect on Kp as partial pressures of top and bottom change to keep Kp same More negative E° Loses electrons (oxidation) Good reducing agent More positive E° Gains electrons (reduction) Good oxidising agent Negative electrode Oxidation (anode) Positive electrode Reduction (cathode) Salt bridge -connects circuit, free moving ions conduct charge between half cells -potassium nitrate used as it is unreactive with electrodes and electrode solutions -wire not used as it would set up its own electrode system with the solutions Use of high voltmeter in electro cell system Stops current in the circuit as to measure Emf Use of platinum electrode Used if system doesn't include metal that can act as an electrode Pt used as it is unreactive and can conduct electricity Standard Hydrogen Electrode (SHE) conditions - 1M solutions - 100kpa H2 gas - 298K - no current Hydrogen fuel cell advantages and disadvantages + less pollution + greater efficiency (less heat energy wasted) - expensive - safety issues storing and transporting H2 - H2 produced by electrolysis of water (uses electricity) What happens if current flows in an electrochemical cell? -reactions occur separately at each electrode -voltage falls to zero as reactants used up Strong acid an acid that dissociates completely in aqueous solution Weak acid an acid that only slightly dissociates in aqueous solution equivalence point of a titration The point at which the unknown solution has exactly reacted with the known solution. Neither is in excess Sudden vertical part of pH curve Buffer A solution that resists changes in pH when small amounts of acid or base are added. Reaction sodium with oxygen Yellow flame to produce white solid Reaction of Mg with oxygen White flame to give white solid smoke Reaction of Al with oxygen White flame to give white solid smoke Reaction of Si with oxygen White flame to give white solid smoke Reaction of P with oxygen White flame to give white solid smoke Reaction of S with oxygen Blue flame to give acidic choking gas Complex Central metal ion surrounded by ligands Ligand Atom/Ion with a lone pair that forms co-ordinate bonds to metal. Coordinate bonding Shared pair of electrons in the covalent bond come from only one of the bonding atoms coordination number The number of coordinate bonds attached to the central metal ion Monodentate Forms one coordinate bond per ligand Bidentate Forms two coordinate bonds per ligand Multidentate A type of ligand that can form many coordinate bonds with ions What can cause a colour change? (Transition metals) 1) oxidation state 2) coordination number 3) ligand heterogeneous catalyst A catalyst that is in a different phase from that of the reactant substances. homogeneous catalyst a catalyst in the same phase as the reactants