Chemistry AS

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atomic structure

  • proton m = 1, c = 1
  • neutron m = 1, c = 0
  • electron m = 1/1840, c = -1
  • isoptopes have same number of protons and electrons but different number of neutrons
  • S orbtial is spherical, P is dumbell
  • electrons negatively charged, repel eachother so spin in opposite directions
  • 1s1 2s2 2p6 3s2 3p6 4s2 3d10
  • group 1 and 2 = s, transition metals = d, all others = p
  • ions = atoms with a charge, +ve if lost e-, -ve if gained e-
  • ionisation energy - how easily electrons are lost to form +ve ions
  • first ionisation energy - energy required to remove one mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions
  • factors: 1) atomic radius  2) nuclear charge  3) shells and shielding
  • trends: increase across period, sharp decrease from end of one peroid to start of next, decrease down group, drops slightly with each new subshell
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mass

  • relative atomic mass used to cmpare Ar of elements
  • rfm is compared to 1/12th of the mass of carbon 12
  • VIADD - vapourisation (turned to gas) ionisation (e- knocked off forming +ve ions) acceleration (electric field) deflection (magnetic field, heavier ions - less) detection
  • shows relative isotopic masses and abundances
  • furthest right peak = molecular mass
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ionic bonding

  • between +ve metal and -ve non metal
  • electrons transferred, held by electrostatic attraction
  • ions attract oppositely charged ions around them = giant ionic lattices
  • ammonium NH4+, hydroxide OH-, nitrate NO3-, nitrite NO2-, hydrogen H-, carbonate CO3-, carbonate CO3 2-, sulphate SO4 2-, sulphite SO3 2-, phosphate PO4 3-
  • high melting/ boiling point - particles held tightly
  • hard/ brittle - layers slip when repulsion occurs
  • soluble - dipoles overcome electrostatic forces
  • conduct in liquids - ions are charged and free to move, no electron density between ions
  • electrolysis: green coppper II chromate VI --> yellow chromate at anode, blue copper at cathode
  • polarisation - distortion of electron cloud
  • cation distorts anion - has the polarising power, anion = polarisable
  • cation: smaller = better as charge condensed, larger charge = better as more attraction
  • anion: larger = better, electron cloud further from nucleus so held loser
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covalent bonding

  • electrons shared not transferred
  • strong intramolecular forces, weak intermolecular forces, electrostatic attraction
  • directional - only between atoms
  • concentrated charge
  • dative covalent = one atom supplies both electrons
  • giant covalent = crystals, strong bonds, hard, high melting point, dont conduct, fixed electrons
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metallic bonding

  • atoms in solid metals and alloys held together
  • atoms are ionised, electrons are fixed in lattice as tightly as possible
  • sea of delocalised electrons, electrostatic attraction
  • strength: more electrons = stronger due to higher electron density, bigger radius = weaker, cloud must cover larger area
  • good conductors: electrons free to move and carry charge
  • flexible: layers can slide over each other
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shapes of molecules

  • EPRT - electron pair repulsion theory
  • lone pair lone pair repulsion > lone pair bond pair repulsion > bond pair bond pair repulsion
  • linear - 180 - 2B
  • trigonal plannar - 120 - 3B
  • tetrahedral - 109.5 - 4B
  • trigonal bipyramidal - 120 / 90 - 5B
  • octrahedral - 90 - 6B
  • pyramidal - 107 - 3B 1L
  • angular - 104.5 - 2B 2L
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carbon

  • graphite
    • high melting point - covalent bonds
    • malluable - layers slide over each other
    • conducts - free electrons suspended in layers
    • used as lubricant
  • diamond
    • high melting point - covalent bonds (3200)
    • strong - each carbon forms 4 bonds
    • doesnt conduct - all electrons used in bonding
    • used in tools
  • buckminster fullerene
    • new allotype found 1985
    • C60 - 20 hexagons, 12 pentagons
    • each carbon has 3 bonds so can conduct
  • nanoscience
    • few hundred atoms, very strong, good conductors
    • could be used as catalysts as they have a large SA:V 
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electronegativity

  • ability of an atom to attract an electron pair in a covalent bond to itself
  • smaller atoms are more electronegative
  • increase across period and up groups
  • F>O>Cl/ N
  • greater the difference between electronegativies - the greater the ionic character, more similar - greater covalent character
  • Na+Cl- = ionic, H(d+) Cl(d-) = polar covalent, Cl-Cl = covalent
  • 100% ionic - complete electron transfer from metal to non metal - never happens
  • ionic always has covalent character as there is a degree of e- sharing
  • covalent always has ionic character - electrons not shared equally so degree of transfer
  • non polar - bonded electrons equally shared, may have p. bonds but no net dipole, atoms have same electronegativity
  • polar - electrons not equally shared, greater electronegativity difference = greater polarity
  • permanent dipoles - dipoles cancel out if molecule is symetrical, otherwise polar molecule
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intermolecular forces

  • van der waals / london
    • exist between all molecules, weakest type - caused by attractions between small dipoles
    • electrons constantly moving - causes instantaneous dipole which induces dipole on neighboring molecule
    • always changing as e- move very fast
    • more electrons = stronger, longer = stronger as more SA
  • dipole - dipole
    • 2nd weakest, act with london
    • permanent dipoles attract opposte dipoles
    • bigger molecules = stronger - more space between dipoles so more sticky
    • long thin shape = stronger - greater SA:V
    • more electrons = stronger
  • hydrogen
    • strongest, 1/10th strength covalent bond bond forms straight line
    • H must be bonded to N/O/F (between H and lone pair of electrons on NOF)
    • NOF hae low electron densities and high electronegativities
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water

  • ice - less dense than water
    • H bonds hold water molecules apart in open lattice structure
    • tetrehedral shape -  collapses when melts
  • high melting / boiling points
    • strong hydrogen bonds and london and dipole - dipole also present
  • each H2O can form four bonds, two lone pairs and two oxygens which is perfect ratio as there are enough lone pairs to bond with each H
  • proof: bp of groups 6 and 7 decrease down group but NOF is high, must be extra forces
  • solvents
    • like dissolves like
    • bonds made must be stronger than those broken
    • forces in two liquids must be equal for them to mix
  • water
    • energy to overcome electrostatic forces is supplied when positive ions are attracted to O- and negative ions are attracted to H+
    • ions are hydrated
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infra-red spectrometry

  • bonds in molecules vibrate, increases when IR is absorbed
  • IR light passed through compounds to gain spectrum
  • absorbtion peaks formed as energy is taken in
  • frequency of peaks used to match to known bonds
  • bonds can:
    • symetric stretch
    • asymetric stretch
    • bend
  • groups:
    • C-H: organic 2800 - 3100 strong sharp
    • O-H: alcohols 3200 - 3550 strong broad
    • O-H: carboxylic acids 2500 - 3300 medium broad
    • N-H: amines 3200 - 3500 strong sharp
    • C=O: aldehydes/ ketones/ acids 1680 - 1750 strong sharp
    • finger print region 1000 - 1550
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mass spectrometry

  • VIADD
  • m --> m+ + e- ions formed (positive)
  • fragmented ions
    • some ions fragmented by bond fission
    • can take place across any bond
    • results give mass and molcules and fragments
  • different isomers give different masses
  • most abundant is base peak - most stable
  • only ions show up - free radicals are lost
  • common fragments:
    • CH3 - 15
    • C2H5 - 29
    • C3H7 - 43
    • OH - 17
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group 1 and 2

  • general properties
    • hydroxides that are alkaline
    • g1 - alkali metals, g2 - alkali earth metals
    • S block metals -  g1:1 atom in S orbital, g2: 2 atoms in s orbital
  • physical properties
    • group 1
      • soft - can cut with knife
      • low melting and boiling points - decrease down group
        • weaker metallic bonding, larger ions so sea of electrons further away
      • low densities - decrease down group
      • form colourless compounds
    • group 2
      • high densities - increases down group
        • mass increases faster than volume, strong metallic bonding
      • high melting and boiling points decrease down group
      • colourless compounds formed
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group 1 and 2 cont.

  • flame colours
    • group 1 
      • lithium - red
      • sodium - yellow
      • potassium - lilac
    • group 2
      • magnesium - white
      • calcium - brick red
      • strontium - crimson red
      • berilium - apple green
    • colours - thermal energy exites e- to a higher level, then falls back down and excess energy released as light, colour depends on difference between energy levels
  • ionisation energy: decrease down groups as radius increases, more shielding as there are more shells, cancels out higher nuclear charge
  • reactivity: reactive metals, strong oxidising agents, increases down the group
    • g1: lose e-   m --> m+ + e- 
    • g2: lose 2 e-   m--> m2+ + 2e-
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group 2 metals reactions

  • oxygen
    • vigorous reaction - increases down group
    • 2Mg (s) + O2 --> 2MgO (s) 
  • chlorine
    • gives chloride when heated in Cl
    • Mg (s) + Cl2 (g) --> MgCl2 (s)
  • water 
    • vigorousity increases down group
    • effervesence and goes cloudy
    • elements lower than Ca are all vigorous reactions giving hydroxide and water
    • reacting slowly with cold water gives hydroxide
      • Mg (s) + 2H2O (l) --> Mg(OH)2 (aq) + H2 (g)
    • reacting quickly with stream gives oxide
      • Mg (s) + H2O (g) --> MgO (s) + H2 (g)
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group 2 metal oxide reactions

  • water
    • forms alkaline solution / hydroxide
    • MgO (s) + H2O (l) --> Mg2+ (aq) + 2OH- (aq)
    • solubility increase down group
      • bigger charge density so less attraction to OH- ions so larger diposture ions
      • split away easier meaning greater concentration of OH- in the water and more dissociate so solution is alkali
    • Ca(OH)2 neutralises acidic soils
      • Ca(OH)2 (s) + 2H+ (aq) --> Ca2+ + 2H2O (l)
    • Mg(OH)2 used as an antiacid in idigestion tablets
      • Mg(OH)2 (s) + 2H+ (aq) --> Mg2+ + 2H2O (l)
  • acid: metal oxides are bases so will neutralise acids forming a salt and water
    • MgO (s) + 2HCl (aq) --> MgCl2 (aq) + H2O (l)
  • solubility of sulphates
    • decreases down group - BaSO4 is insoluble, used as test for sulphate ion
    • BaCl2 added to X in dilute nitic acid, if X has sulphate dense white ppt forms, used for x rays to show gut movement: Ba2+ + SO4(2-) --> BaSO4
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thermal stability S block compounds

  • thernal decomposition - splitting by heat
  • carbonates
    • more stable down group so more difficult to decompose
    • polarising abiity of metal cation decreases so less distortion to surrounding negative ions so bonds are stretched so weaker
    • the more shells = less polarising power
      • MCO3 (s) --> MO (s) + CO2 (g) 
    • group 1 have larger cations so less polarising except lithium which is very small
      • LiCO3 (s) --> Li(2)O (s) + CO2 (g)
  • nitrates
    • group 1
      • give out O2
      • Li decomposes as group 2 so NO2 and O2
        • MNO3 (s) --> MNO2 (s) + 1/2O2 (g)
    • group 2
      • give out NO2 and O2
      • more stable down the group - same as carbonates
        • 2M(NO3)2 (s) --> 2MO (s) + 4NO2 (g) + O2 (g)
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solubility S block compounds

  • hydroxides more soluble down group
  • sulphates less soluble down group
  • carbonates less soluble down group
  • resonant forms of CO3(2-):
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group 7 elements and compounds

  • general properties
    • diatomic molecules
    • boiling points and van der waals forces increase down group
    • electronegativity and reactivity decrease down group
  • F - colourless gas
  • Cl - yellow/ green gas, yellow/ green with H2O and hydrocarbon
  • Br - red liquid, red with H2O (partially soluble) and with hydrocarbon - very soluble
  • I - grey solid/ purple gas, pale yellow with H2O, pink/ red with hydrocarbon
  • reactivity as oxidising agents
    • very strong, oxidising power decreases down group
    • halogens are reduced, each atom gains 1e-, electron attracted to outer shell by nuclear charge
    • F2 is strongest
    • bigger radius, more shielding = less oxidising down group - less force on e-
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group 7 reactions

  • displacement
    • demonstrates reactivity of halogens
    • halogens displace those below them, can be identified by their colours
    • Cl2 + 2Br- --> 2Cl- + Br2
    • colours:
      • Cl2 - H2O pale green, hexane pale green
      • Br2 - H2O orange, hexane orange
      • I2 - H2O brown, hexane purple
  • metals
    • oxidise many metals to form ionic chlorides
    • 2Na (s) + Cl2 (g) --> 2NaCl (s)
  • ions
    • oxidise some ions to higher oxidation states
    • 2Fe(2+) + Cl2 --> 2Fe(3+) + 2Cl-
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halides tests

  • halide ions
    • add silver nitrate (AgNO3) and colourless ppt forms
    • Cl (AgCl) - white ppt, soluble in aqeuous ammonia
    • Br (AgBr) - cream ppt, insoluble in aqeous ammonia, soluble in conc
    • I (AgI) - yellow ppt, insoluble in aqeuous and conc ammonia
    • Cl and Br also darken in UV light - I doesnt 
  • hydrogen halides
    • colourless gases, very soluble in water, form strong acidic solutions - increase down group: HCl +aq --> H+ Cl-
  • reactions with H2SO = oxidising agent / acid
    • NaCl - misty fumes, HCl = -1 displacement of Cl-
    • NaBr - misty fumes HBr = -1 displacement of Br-, brown vapour Br2 = 0 oxidation of Br-,colourless gas SO2 = +4 reduction of H2SO4
    • NaI - misty fumes HI = -1 displacement of I, purple vapour I2 = 0 oxidation of I-, colourless gas SO2 = +4 reduction of H2SO4, yellow solid S = 0 reduction of H2SO4, bad egg smell H2S = -2 reduction of H2SO4
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group 7 cont.

  • disproportionation - the same element is oxidised and reduced
  • Cl2 in H2O: Cl reduced to Cl- and oxidised to ClO-
    • Cl2 (0) + H2O --> HClO (-1)  + HCl (+1)
  • Cl2 in dilute aqeous alkalis
    • reacts with halogens at room temperature - basis for bleach
    • Cl2 (0) + 2NaOH --> NaCl (-1) + NaClO (+1) + H2O
  • Cl2 in concentrated aqeous alkalis eg. hot sodium hydroxide
    • 1) I2 + 2NaOH --> NaI + NaIO + H2O
    • 2) 3NaIO (+1) --> NaI (-1) + NaIO3 (+5)
  • fluorine: gas, highly electronegative and oxidising, fluorocarbon very stable
  • astatine: solid, low electronegativity, least oxidising, astocarbon least stable
  • acid - base titrations
    • acid is standard solution
    • alkali into conical flask via pipette
    • concordant titres - must be with 0.2cm
    • indicators:
      • methylorange: yellow (alkali) --> orange
      • phenolphthalein: pink (alkali) --> colourless
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mole

  • used by chemists to count atoms, symbol = n
  • carbon 12 is standard - number of moles in 12g carbon 12 is 1
  • 1 mole = 6.02 x10(23) atoms - avagadros number
  • moles = mass / rfm   ---   mass = moles x rfm  ---   rfm = mass / moles
  • moles = concentration x volume  ---   v = m / c   --- c = m / v
  • 1dm = 1xm / 100
  • gas
    • if temperature and pressure are the same then volume of gases will be the same
    • 1 mole of gas occupies 24dm under standard conditions
    • volume = moles / 24  ---  moles = volume / 24
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formula and equations

  • empiracal: simpliest whole number ratio of atoms in each element in a compound
  • molecular: actual number of atoms in each element of a molecule
  • displayed: molecule and structure drawn in full
  • structural: bonds not shown
  • skeletal: carbons drawn as lines, each end = 1 carbon
  • equations
    • chemical reactions involve the rearrangement of atoms and or ions
    • qualitive - what atoms / ions are rearranging
    • quantitative - how many atoms / ions are rearranging
  • percentage yield
    • mass of the actual product as a % theorectical
    • p y = actual / theoretical x 100
  • atom economy
    • higher economy means fewer waste materials
    • addition reactions have 100% atom economy
    • a e = Mr of desired / Mr of all products x 100
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redox

  • oxidation - loss of electrons so charge increase
  • reduction - gain of electrons so charge decreases
  • if one species gains, the other must lose
  • OILRIG - oxidation is loss, reduction is gain
  • oxidising agent - accepts electrons from another reactant
  • reducing agent - donates electrons to another reactant
  • oxidation numbers:
    • uncombined elements = 0
    • hydrogen = +1 (metal hydrides = -1)
    • fluorine = -1
    • oxygen = -2 unless with fluorine - then calculate using F values (peroxides = -1)
    • charge given = that shown
    • chlorine = -1 unless with O or F - becomes +
    • bromine = -1 unless with O or F or Cl - becomes +
    • iodine = -1 unless with O or F or Cl or Br - becomes +
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half equations

  • electron transfer: 
    • oxidation - Mg --> Mg2+ +2e-
    • reduction - Cl2 +2e- --> 2Cl-
    • combined: Mg + Cl2 --> MgCl2
  • can combine to form full equations
    • Ag+ and Zn:
      • Ag+ + e- --> Ag
      • Zn --> Zn2+ + 2e-
  • electrons must balance across both equations
    • Ag equation x 2: 2Ag + 2e- --> 2Ag
  • half equations added together
    • 2Ag+ + 2e- + Zn --> 2Ag + Zn2+ + 2e-
  • cancel electrons out
    • 2Ag2+ + Zn --> 2Ag + Zn2+
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organic chemistry

  • hydrocarbons
    • compounds containing only hydrogen and carbon
    • saturated have single bonds only
    • unsaturated contain one or more double bonds
  • homologous series
    • series of organic compounds with the same functional group
    • same general formula, differing from next compound by CH2
    • similar chemical properties, gradual change in physical properties
    • prepared by similar methods
  • functional groups
    • groups of atoms responsible for the characteristic reactions
    • reactive part of the compound
  • naming carbon chains:
    • meth - 1 carbon
    • eth - 2 carbons
    • prop - 3 carbons
    • but - 4 carbons
    • pent - 5 carbons
    • hex - 6 carbons
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naming organic compounds

  • types of compounds: 
    • alkane - ane - CnH2n+2 - C-H
    • alkene - ene - CnH2n - C=C
    • halogenoalkane - (bromo)ane - CnH2n+1X - C-X
    • alcohol - ol - CnH2n+1OH - C-OH
    • aldehyde - al - CnH2nO - C(-H)=O
    • ketone - one - CnH2nO - R-C(=O)-R
    • carboxylic acid - oic acid - CnH2n+1COOH - C(=O)-OH
  • to name
    • find longest carbon chain
    • identify functional groups and add to start or end
    • start numbering from end which gives lowest possible values
    • write side chains and functional groups in alphabetical order
  • structural isomerism
    • same molecular formula - different structural
    • chain isomer - skeletal formula is different
    • branched isomer - treated as side chain - akyl group
    • positional isomer - functional group can be in different places on chain
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alcohols

  • hydroxyl group - C-OH
  • properties
    • hydrogen bonding so high melting and boiling points, soluble in water
    • less soluble as chain grows - non polar part
    • C and H are electron defficient, O is delta negative
  • primary - OH group attached to carbon bonded to one other carbon
  • secondary - OH group attached to carbon bonded to two other carbonns
  • tertiary - OH group attached to carbon bonded to three other carbons
  • hazard = potential of a substance to do harm, absolute and constant, eg. petrol is toxic and flammable
  • risk = chance that a substance will do harm, variable, petrol in cars vs. petrol on fires, risked can be reduced by identifying hazards 
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alcohol reactions

  • combustion: longer chains give more sooty flame
    • CH3CH2CH2OH (l) + 4O2 (g) --> 3CO2 (g) + 4H2O
  • sodium
    • gives sodium alkoxide, effervesence, white film forms, redox reaction, room temp
    • 2CH3CH2CH2OH + 2Na --> 2CH3CH2CH2O-Na+ + H2
  • phosphorus (V) chloride
    • gives chloroalkane, white flame (HCl) given off, more dense with ammonia fumes
      • HCl + NH3 --> NH4Cl - used to test for OH group
    • nucleophillic subsitution - OH for Cl
    • CH3CH2CH2OH + PCl5 --> CH3CH2CH2Cl + PCl3 + HCl
  • bromide ions
    • KBr and 50% H2SO4, relfux reaction - the acid oxidises, OH replaced with Br
    • CH3CH2CH2OH + KBr --> CH3CH2CH2Br + KOH
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oxidation of alcohols

  • aldehyde (primary)
    • add K2Cr2O7 in conc H2SO4, distillation reaction - H2 lost double bond formed
    • orange --> green (Cr ions 6+ to 3+) Fehlings blue to brick red
    • CH3CH2CH2OH + [O] --> CH3CH2CHO + H2O
  • carboxylic acid (primary)
    • excess oxidising agent so further oxidation, heat under reflux
    • Fehlings stays blue, test for by using Na2CO3 - bubbles of CO2 form
    • CH3CH2CH2OH + 2[O] --> CH3CH2COOH + H2O
  • ketone (secondary)
    • orange to green, heat under reflux
    • Fehlings stays blue
    • CH3CH(OH)CH3 + [O] --> CH3COCH3 + H2O
  • phosphorus ad iodine
    • gives iodoalkane, reflux reaction, substitution
    • CH3CH2CH2OH + PI --> CH3CH2CH2I
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halogenoalkanes

  • CnH2n+1X
  • polar bond: C(d+)-X(d-)
  • polarity decreases from F to I as electronegativity decreases, attacked by nucleophiles = reagents that seek positive centres
  • nucleophillic substitution / hydrolosis
    • produces an alcohol - work only for primary 
    • NaOH (used for OH-) donates electron pair - forms NaBr
    • CH3CH2CH2Br + OH --> CH3CH2CH2OH + Br-
    • rate of hydrolosis increases as the C-X bond gets weaker
    • test - add nitric acid then AgNO3 and a white ppt should form
  • elimination
    • halogenoalkane heated under reflux with KOH in anhydrais
    • 78 degrees, ethanol used as a solvent to prevent nucleophillic substitution
    • OH- from KOH acts as base extracting a proton so the halide follows and an alkene forms
    • CH3CH2Br + OH --> CH2=CH2 + Br- + H2O
  • rates of reaction increase down group due to bond strength - tertiary are fastest due to stable intermediate carbocation
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halogenoalkanes cont

  • nucleophillic substitution to amine
    • ammonia in ethanol, reflux under pressure
    • excess ammonia removes HBr
  • halogenoalkane --> alkene = elimination, KOH in ethanol
  • halogenoalkane --> alcohol = nucleophillic substitution, KOH and H2O
  • halogenoalkane --> amine = nucleophillic substitution, NH3 in ethanol, pressure
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nucleophillic substitution mechanisms

  • primary
    • SN2 - substitution nucleophillic 2nd order (2 species in slow step)
    • CH3Br + OH- --> CH3OH + Br-
    • nucleophile attacks from behind due to lack of space
    • transition state - Br bond weakened
  • tertiary
    • SN1 - substitution nucleophillic 1st order (1 specie in slow step) 
    • C(CH3)3Br + OH- --> C(CH3)3OH + Br-
    • large CH3 causes steric hindrance - OH cant fit
    • CH3 stabalise C+ charge by positive induction effect, electrons pushed toward C so C-Br weak enough to drop
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alkanes

  • CnH2n+2
  • saturated hydrocarbons, unreactive, immisible with water
  • melting points increase with chain length - increasing IM forces, branched chains = lower
  • combustion
    • useful fuels, can be complete or incomplete
    • more carbon means more energy but needs more O2 to burn
    • catalytic converters: CO --> CO2, NOx --> NO2
  • hetrolytic: ions formed, doesnt split equally
  • homolytic: free radicals formed, splits equally
  • free radical substitution
    • initiation: UV light causes homolytic fisson
      • Cl - Cl --> Cl' + Cl'
    • propogation: chain reaction, 2 steps
      • 1) Cl' + CH4 --> CH3' + HCl          2) CH3' + Cl2 --> CH3Cl + Cl'
    • termination: free radicals react
      • 1) Cl' + Cl' --> Cl2          2) Cl' + CH3' --> CH3Cl          3) CH3' + CH3' --> C2H6
    • impure products - get a mixture of products
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crude oil

  • fractional distillation
    • oil heated to 400 degrees, short chains condense at top, long at bottom
    • vacuum so even heavy chains evapourate, must be refined to remove sulphur
    • gas 1-4, naptha 5-10, kerosene 10-16, diesel 14-20, heavy diesel 20, lube oil 20-50, catalytic cracker 20-50, fuel oil 20-70, bitumen >70
    • good nice koalas dont hear lion cubs fighting baddies
  • cracking
    • heavy long chains to short useful chains (less than 12 carbons)
    • aluminium oxide or zeolite catalyst at 500 degrees
    • contiuous process, burns off carbon in the air
    • reactor --> separator --> regenarator --> start again
  • reforming
    • straight chains to rings or arenes
    • metal catalyst platinum 500 degrees and pressure
    • hydrogen by-product
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alkenes

  • CnH2n -ene
  • unsaturated, has C=C, number in name shows double bond position
  • undergo electrophillic addition
  • double bonds
    • single bonds = sigma - overlap of s and p orbitals
    • double bonds = sigma and pi - overlap of p orbitals, pi bond is weaker, split into two parts
  • E / Z isomerism
    • type of stereoisomerism
    • C-C bonds can rotate, C=C cant so groups cant move from one side to the other
    • C=C and two different groups attached to each carbon
    • have different chemical properties - they react differently
    • E (trans) 
      • main groups on opposite sides, one above and one below double bond
    • Z (cis)
      • main group on the same side, both above or below double bond
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reactions of alkenes

  • electrophillic addition with halogens
    • forms dihalogenoalkane, heterolytic fission, orange to colourless
    • double bond repels electrons on Br2 polarising it - heterolytic fission occurs
    • a carbocation is formed - an organic ion containing a positve carbon ion
    • Br- moves to the positive carbon and bonds to form 1,2-dibromoethane
    • C=C + Br-Br --> CH2BrCH2+ --> CH2BrCH2Br
  • electrophillic addition with hydrogen halides
    • HBr - H turns delta positive Br delta negative, H joins one carbon, Br attracted to positive carbon = CH3CH2Br
  • bromine water
    • orange bromine water decolourises in the precence of C=C
    • Br + H2O --> HOBr + HBr
    • C=C + HOBr --> CH2BrCH2OH
  • addition of acidified manganate (VII)
    • cold acidified manganate oxidises C=C to form a diol, purple to colourless
    • ethene --> ethane 1,2 - diol
    • C=C + [O] + H2O --> CH2OHCH2OH
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reactions of alkenes cont

  • reduction with hydrogen / hydrogenation
    • gain of hydrogen / loss of oxygen, forms alkane
    • nickel catalyst at 150 degrees
    • C=C + H2 --> CH3CH3
  • addition polymerisation
    • long chain molecules with high Mr made from joining monomers
    • monomer = unsaturated alkane with C=C, volatile liquids / gases
    • polymers = saturated, solids due to increased van der waals
  • unsymetrical alkenes
    • HBr and propene --> two isometric brominated compounds, can add in two ways forming unequal products
    • HBr - H adds to carbon which as the most H already to produce a more stable cation
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energetics

  • thermochemistry - used to measure and predict energy changes in a reaction
  • system - chemicals being reacted
    • closed: no exchange of matter with surroundings
    • open: exchange can occur
  • surrounding - whole universe
  • chemical reactions - exchange energy between system and surroundings
  • 1st lw - energy may be exchanged between a system and surroundings but total energy remains constant
  • exothermic - energy from system to surroundings, reacting chemicals lose energy (-)
  • endothermic - energy from surrounding to system, reacting chemicals gain energy (+)
  • enthalpy - measure of heat content of a substance at a constant pressure per mole
  • enthalpy = Q = m x c x change in temp (m = mass) (c = specific heat capacity) 
  • c usually 4.18j/g//k = H2O value
  • m is only of liquids, taken as 1g / cm3
  • limitations: heat lost via container to surroundings, incomplete combustion, high accuracy and low reliability, repeats should be taken
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standard enthalpy change

  • of reaction H-r: under standard conditions in standard states, reactions occurs in molar quantities shown in equation
  • of formation H-f: 1 mole of a compound formed in its standard state from elements in their standard states and standard conditions, elements = 0 
  • of combustion H-c: 1 mole of a substance undergoes complete combustion in O2 under standard conditions in standard states, 1 CO2 for every C, 1 H2O for every 2H
  • of neutralisation H-neut: 1 mole of H2O formed from neutralisaton of hydrogen ions and hydroxide ions under standard conditions, exothermic = -57kj/mol strong alkali and acid
  • of atomisation H-at: 1 mole of gaseous atoms formed from its element in standard state
  • bond dissociation enthalpy
    • energy required to break 1 mole of gaseos bonds to form 1 mole of gaseous atoms
    • enothermic - energy required to break bonds, exo = making bonds
    • bond strength depends on the environment
    • bond enthalpy = 2 x enthalpy of atomisation for diatomic gases
    • smaller enthalpy means weaker bonds so easier to break 
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Hess's law

  • enthaply change in converting reactants to products is the same regardless of the route taken if the initial and final conditions are the same 
  • steps to calculate
    • 1) write the equation of the reaction
      • balanced with state symbols
    • 2) use additional information to complete the cycle
      • formation - arrows go up or down
    • 3) apply Hess's law by following the arrows
      • if there are two moles of something you must times energy by 2
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bond enthalpy

  • stored within chemical bonds, indicates the strength of bond within a gaseous molecule
  • definition: energy needed to break and separate 1 mole of bonds in the molecules of a gaseous element or compound so the resulting gaseous species exhert no forces upon eachother
  • must use average bond enthalpy as some bonds are different in different environments
  • bond breaking - exo - needs energy
  • bond making - endo - releases energy
  • enthalpy changes = reactants - products
  • theoretical - assumes perfect roundness, 100% bonding and that complete separation occurs
  • experimental - heat given out to surroundings, assumer density, molecules cant move infinitely apart, dehydration reactions cant be controlled
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chemical equilibrium

  • not all reactions go to completion, some end with a mixture of reactants on the right hand side and products on the left hand side
  • dynamic equilibrium = rates of reaction forward and backward are the same, no apparent change, concentrations are constant, only takes place in a close system where nothing is added or removed
  • chateliers principle - when a change is applied, the system reacts to oppose the effect of the change
    • greater conc of reactants - equilibrium shifts right to produce more products
    • increase total pressure - equilibrium shifts to side with the least gas molecules
    • increase temperature - equilibrium shifts in endothermic direction
  • catalyst - doenst change the position of the equilibrium but speeds up reaction so its reached faster = provides an alternative pathway for the reaction with a lower activation energy, save money and energy and provide a better atom economy - not used up 
  • Maxwell- Boltzmann distribution
    • catalyst - Ea shifts to the left
    • temperature - whole curve shifts right and down with higher temperature so more molecules with Ea
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kinetics

  • collision theory
    • particles must collide before a reaction takes place
    • must approach in a certain way (steric hindrance) and moving fast enough to reach activation energy so not all collisions mean a reaction
  • increase rate
    • more frequent collisions = more speed and more particles
    • more successful collisions = more energy and lower activation energy
    • increase pressure
    • increase temperature
    • increase concentration
    • increase surface area
    • add catalyst
  • homogenous catalyst: some state, offers lower energy intermediate step eg. NO / O3
  • heterogenous catalyst: different state to reactants, powder form, often mounted on frame so easy to remove products
    • 1) absorbtion - bonds made with catalyst weakening bonds so reaction is easier
    • 2) reaction - things held on surface can react more easily
    • 3) desorbtion - products are released
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chemistry in the environment

  • greenhouse effect: natural process, keeps earth warm enough to support life, most IR radiation gets reflected back to space, GHG can trap it, equilibrium maintains steady temperature: 78% N2, 21% O2, 0.9% Ar, 0.04% CO2
  • green house gases: absorbtion of long wave IR causes bonds to vibrate - polar gases
  • climate change: natural cycles cause warming and ice ages, short term can be caused by sunspots and volcanic activity, always experienced climate change, GHG released by humans causing rapid warming - melting ice caps = sea level rise
  • ozone layer: 10-60km thick in stratosphere, formed by the interaction of UV light and O2
    • O2 --> 2O'          O2 + O' --> O3 + heat          O2 +O <--> O3
    • natural concentration maintained, UV absorbed, heat evolved = temp increases
    • ozone - protects against skin cancers, can be depleted by free radicals Cl' and NO'
    • Cl removes O3: O3 + Cl' --> ClO' + O2       ClO' + O --> O2 + Cl'        O3 + O --> 2O2
    • NO removes O3: O3 + NO' --> NO2' + O2   NO2' + O --> O2 + NO'     O3 + O --> 2O2
  • CFCs vs. ozone
    • CFCs banned since 2000, found in fridges and aerosols
    • ideal for use - highly volatile, non toxic, non flammable, no smell, unreactive
    • unreactive CFC diffuse to upper atomosphere, split by UV forming free radicals
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green chemistry

  • chemistry must reinvent itself to be greener
  • more developed countries have more options
  • more efficiet processes needed:
    • eliminate hazardous chemicals
    • higher atom economies
    • non toxic waste through recycling
    • consumer less energy by using renewables
  • anthropogenic climate change: short time scale, human activities eg. burning fossil fuels
  • carbon neutral: CO2 released = CO2 removed during growth so no net CO2 emission
  • carbon footprint: mass of CO2 from production to consumption
  • five key points:
    • renewables
    • alternatives to hazardous chemicals
    • catalysts
    • energy efficiency
    • reduced waste
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