Chemistry AS

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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