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Atomic Structure

PARTICLE           MASS              CHARGE

Proton                     1                         +1

Neutron                   1                          0

Electron                1/2000                  -1

ATOM-  Nucleus = P+N with e surrounding in orbitals

MASS NUMBER(A)(TOP)- P+N

ATOMIC NUMBER(Z)(BOTTOM) - Number P (e= same number of p)

ISOTOPES- Atoms with same P but different number of N

Ions have different amount of P/E. formed by GAIN OR LOSS of E.

- = GAINED e              + = LOST e

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Atomic Structure (3.1.1.2)

MASS SPECTROMETRY (4)

Electrospray Ionisation- Sample dissolved in polar solvent --> nossel, high pressure --> high voltage applied to sample removing an E to form a + charge= GAS of + ions

Acceleration- ions accelerated between + and - plates on electric field= kinetic energu gained(all ions SAME KE) Depends on MASS. Low= faster

ion drift- particles= constant KE as travel through drift region.    

NOTE Different ions, different masses, different velocities = different TOF

Detection- Detects current created when ion hit it. Records how long took for travel through mass spectrum. Data calculates mass/charge values

Produces chart which relative isotopic abundance can be found.

USES- drug detection, forensic science

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Atomic Strcuture 3.1.1.3

1st ionisation energy- Energy needed to remove 1 e from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

X(g) ---> X+ (g)  + e-

Period (across IE increase)- Inrease in P as go across = increase in nuc charge= Greater attraction between NUC and E

Group- (Down IE decrease)-  as shielding increase = less effective nuclear attraction = less energy needed.

Successive ionisation(IE INCREASE) - E removed from progressively more + species = greater attraction = more energy required

Change in energy level (IE increase) -  Less shielding = more effective nuclear attraction - increase in energy required

Period 3- Mg (3s2) --Al (3p1) = 3p orbital= higher energy= futher away, created fropm partial shielding effect. This overrides increased nuclear charge= less energy required

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Amount of Substance 3.1.2.1- 2.4

Mole- 6.02 x 1023

Avogardo constant- 6.02 x 1023 - number of particles in 1 mole of a subatnce

Relative atomic mass(Ar)- Avergae mass of an atom of an element on a scale where an atom of carbon -12 is exactly 12

Relative Molecular mass(Mr)- average mass of a molecule on a scale where an atom of carbon-12 is exactly 12

Carbon 12-  Mr is measured on the relative mass scale with carbon once again being the reference atom

Relative formula mass can be used in ionic equations as the formula is taken in the simplest ratios

Emprical Formula- Formula which gives the simplest whole number ratio of atoms of each element in a compound

Molecular Formula- Formula which shows the actual number of atoms of each element in a compund

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Amount of Substance 3.1.2.4- 2.5

Relationship between empirical formula and molecular formula- Need emprical formula and relative Mr of the compound to calculate MF

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Bonding 3.1.3.1

IONIC( e.g NaCl)

  • Electrostatic forces of attracition between + and - ions in giant lattice= STRONG. 
  • CONDUCTS-  in liquid as e = free to move = carry flow of charge
  • NON malleble -> repulsion= cant slide
  • Soluble- H2O pulls ions away from lattice = dissolve
  • HIGH BP- strong e.s forces beytween ions = energy required to overcome

COVALENT- Includes 2 atoms sharing e so both have full outer shell

  • Single bond- shared pair of e both nuceli electrostatically attracted to
  • Double bond- atoms share 2 pairs of e
  • Triple bond- atoms share 3 pairs of e

Simple Covalent Compounds (e.g I2)

  • Strong C.B within molecules but compund = weak intermolecular forces
  • DONT conduct- no free e to carry charge
  • VARIES in soulubility as depends on polarity
  • LOW BP- Forces = easily broken
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Bonding 3.1.3.1

Giant Covalent/Macromolecular(2)-

  • Graphite- Flat hexagon sheets
  • C.B 3 bonds each 4th = delocalised e
  • sheets bonded by VDW's = easily broken = slide over
  • HIGH BP- Strong CB 
  • CONDCUCT- delocalised e 
  • INSOLUBLE- CB too difficult to break  
  • Diamond - Carbon CB to 4 other Carbon atoms= Tetrahederal shape 
  • HIGH BP- strong CB
  • CANT conduct - all e localised 
  • INSOLUBLE- CB too difficult to break

Metallic (e.g Mg)- Giant lattice of + ions and delocalised e 

  • electrostatic forces of attraction between the 2 
  • conduct electricity in BOTH states = delocalised e 
  • MALLEABLE(shape/bend)- ions same size can slide over eachother 
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Bonding 3.1.3.1-4

Dative Covalent Bond- CB formed when 1 atom provides both of shared e.

They form when 1 atom in the bond has a lone pair and the other doesnt have any available e to share (Shown as an --> representing where the e come from and where theyre going to)

Ions formed from elements 

  • Group 1 - LOSE 1 = 1+
  • Group 2 - LOSE 2 = 2+
  • Group 6 - GAIN 2 = 2- 
  • Group 7 - GAIN 1 = 1-

Compound Ions 

  • Ammounium-  NH4+
  • Carbonate-     CO3 2-
  • Hydroxide-     OH-
  • Nitrate-           NO3-
  • Sulfate-          SO4 2-
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Bonding 3.1.3.4-

Charge Clouds- Molecules = different shapes, depends on e pairs. E can be in bonding pairs or lone pairs 

Electron Pair Repulsion- e are - so repel eachother until far apart as possible. Shapes of clouds affect how much repels other charge clouds. Lone pair repel more than bonding pairs

  • LP-LP       Most
  • LP-BP
  • BP-BP      Least
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Polarisation 3.1.3.6-

Electronegativity- Power of an atom to attract the pair of electrons in a C.B

  • This ^ will be unsymetrical between 2 different elements as 1 will be more EN than the other so the e will be pulled towards that atom 
  • Greater ^that difference = more polar the bond
  • polar bond- difference in EN between the 2 causes a DIPOLE
  • Dipole- difference in charge between the 2 atoms caused by a shift in e density in bons
  • Polar bonds = Permenant dipole = charge disturbuted unevenly throughout whole molecule
  • more than 1- PB arranged symmetrically dipoles cancel out = means becomes non polar 
  • more than 1- PB arranged in same direction- dont cancel = stays same 
  • ALL Diatomic Gases = Non polar = atoms equal EN
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intermolecular forces 3.1.3.6

Van Der Waals Forces-

  • Cause all atoms and molecules be attracted eachother(overrall effect)
  • E in charge clouds move quick (constant- made/destroyed any time) - likely to be more to one side than the other = temp dipole
  • this occur in neighbouring atom = domino effect
  • larger molecules= strong VDW     closer m = Stronger VDW
  • larger molecules, stronger VDW, higher bp

Permenant dipole-dipole attraction- sub made up of this there is weak electrostatic frces of A on +/-  charges on neighbouring molecules

  • electrostatically charged rod next to polar liquid it will move towards rod - molecules move to attract themselves to it charge of R = irrelevant
  • more polar liquid is = stronger electrostatic attraction between the 2 = greater deflection

Hydrogen Bonding- ONLY occurs when H is bonded to one of NOF 

  • As V EN so draw bonding e away from H. 
  • Lone pair needed
  • Molecules with H bonding = higher bp/mp
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Importance H bonding 3.1.1.6

Ice- 

  • As liquid H20 cools = ice 
  • during ^ molecules make more H bonds, arrange into regular lattice structure 
  • H bonds long = distance between = longer in ice = less dense than liquid H20 = UNUSUAL - most sub more dense as S than L
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Energetics 3.1.4.1

  • Endothermic- Reaction that absorbs energy - /\H= +
  • Exothermic- Reaction that releases energy- /\H= -
  • Enthalpy Change- Heat Energy transferred in a reaction at constant pressure
  • Standard Enthalpy of Combustion- Enthalpy change when 1 mole of a substance is competlely burned in O2 under standard condtions with all R+P in their standard states
  • Standard Enthaply of Formation- Enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard condtions 
  • Hess' Law- Total enthaply change for a reaction is independent to its route taken 
  • Mean Bond Enthalpy- Average value for the bond enthalpy of a particular bond over the range of compounds its found in 
  • Standard Conditions- 100Kpa in pressure. Temp- 298K
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Kinetics 3.1.5.1

Rate of Reaction- change in the amount of reactant or product over time.

Reactions can only occur when collisions with sufficient energy take place between particles energy

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Equilibria 3.1.6

Reversible reactions = go both ways. 

  • Forward + reverse reactions occur equal rates
  • at EQB conc of reactants and products remain constant 
  • Kc equation can be deduced from this 

Le Chateliers Principle- If a reaction at EQB is subjected to change in conc, pressure or temp the position of EQB will move to counteract the change

  • Changing Conc- e.g increase conc of R EQB tries to get rid of it by making more P = EQB shift to right 
  • Changing Pressure- ONLY affects Gases. Increase P shifts EQB to side with fewer molecules to reduce the pressure 
  • Changing Temp- increase heat = EQB shift to endothermic side. Decrese in temp = EQB to shift to the exothermic side 
  • INDUSTRIAL COMPROMISE( Money&Product)- Temp--> Lower temp better yield but slower rate of R = compromise.      Pressure-->High P in forward reaction = faster rate of R but Higher P = more expensive (better equipment etc) = compromise between 2 

Dynamic EQB- in a RR when F/R reactions are occuring at equal rates and when conc of the R and P are constant

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Oxidation, Reduction and Redox Reactions 3.1.7

Oxidation- Process of é loss 

Oxidising Agent- Accepts é but gets reduced 

Reduction- Process of é gain

Reducing Agents- Donates é and gets oxidised

(combine HE to make Redox)

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Periodicity 3.1.2.1.1

Period 3

Atomic Radius- 

  • Decreases across a period (EN increases)
  • Increases down a group (EN decreases)

Nuclear Attraction

  • Increases across a period (more effective)
  • Decreases down a group (less eff)

Melting Points

  • Increase from Na- Si decrease from Si-Ar
  • Na, Mg, Al = Metals = Mp increase- metal-metal bonds = stronger as, as cross P metal ions increasing + charge and decreasing number of delocalised é
  • P, S, Cl- Molecular Subs - Mp depends on strength of VDW- weak= easy overcome so LOW mp.
  • Ar= small Mp exsits as monoatomic = weak VDW
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Intro to Organic Chem 3.3.1.1

Compounds can be represented in 

  • Molecular formula- states actual number atoms in each atom c2h6
  • Structural- Shows carbon-carbon with functional groups in formula e.g CHCH
  • Displayed- shows how arranged nd bonds
  • Empirical- divide down molecular to get simplest whole number ratio of atoms in that compound
  • Skeletal- Shows bonds of C skeleton only 
  • General Formula- Algebraic formula that can describe any memeber of a family of compounds
  • Homologous series- Family of compounds that have the same functional group and general formula. (Consecutive memebers differ by -Ch2-
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nomenclature 3.3.1.1

(the naming of Organic Compounds)

IUPAC System for naming compounds

  • Count longest chain of carbons 
  • put substiuents in alphabetical order 
  • put functional groups in priority and order from there

Halogens         Alkyl Groups      Alkenes      other functional groups 

LOWEST P                                                          HIGHEST P

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Isomerism 3.3.1.2-3

Structural Isomerism- Isomers same molecular formula but different structural formula. (3 types

  • Chain i- I with the same Functional Groups but different arrangements in carbon skeleton
  • Position i- Same Skeleton and same atoms- atoms attached to different C
  • Functional Group i- Same atoms arranged into different F.G

Stereoisomerism- I with same Structural Formula but atoms arranged differently in space 

E-Z isomerism= form of S.I.    Occurs as a result of restricted rotation about the planar carbon-carbon double bond.

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Alkanes 3.3.2.1

SATURATED Hydrocarbons 

Petrolum/Crude Oil- mixture mostly alkanes- be seperated by F.D

Cracking- Breaking long-chain alkanes into smaller HC. (2 types)  

1. Thermal cracking

  • High Temp and Pressure
  • produduces alkenes( used to make polymers)

2. Catalytic Cracking

  • High Temp and slight pressure
  • zeolite catalyst
  • produces aromatic hydroC 
  • and alkanes used in motor fuels 

Econmic Reasons for C- people want light fractions as, its got more demand = more valuble than heavier which isnt

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Alkanes 3.3.2.1

Fractional Distillation-     TOP=cool     Bottom= Hot 

  • Crude Oil vapourised
  • the V CO goes to bottom - rises through trays
  • largest HC BP too high- bottom- gooey residue
  • bp increase as get bigger = each fraction condense different temp = drawn off different temps 
  • HC lowest BP dont condense = drawn off as gases at top 

ALKANES AS FUEL

found in fossil fuels- burning small amouts release LOTS energy. Burnt in power stations, central heating systems, cars. Release polluntants when burnt 

Complete Combustion -  Burn fuels with plenty o2. Products CO2 and H2O

Incomplete Combustion- Burn fuels in not enough O2, Products CO, H2O

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Alkanes 3.3.2.4

Pollutants (3)- 

  • Nitrogen Oxide NOx, - caused when high temp and pressure in car engine cause N + O atoms from air to react together 
  • Unburnt Hydrocarbons- React with Nitrogen oxides in presence of sunlight form ground-level ozone O3 = smog compoenent= irritant 
  • Sulfur Dioxide- When S burnt in some FF = reacts form SO2 gas. When gets into atmosphere dissolves in the mostiure and converted to sulfuric acid = cause acid rain.

How reduce

  • Catalytic converters in cars remove P from exhaust (complete comb of HC)
  • Sulfur dioxide removed from Power Stations by Flue Gases 
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Halogenoalkanes 3.3.2.1

HalogenoAlkane- Alkane with at least one Halogen atom in place of a Hydrogen atom 

Halogens react with Alkanes in Photochemical rections to form Halogenoalkanes 

Polar- More EN than the carbon atom

Nulepphilic sub(3)- cynanide (makes nitriles), hydroxide (make alcohol), ammonia (make amines)

C-h bond- enthalpy influences the rate of reaction as decides reactivity. Reaction to occur bond = break = weaker = faster substitution

NUC SUB CONDTIONS - React halogenoA under Water under relfux 

Nucleophilic Elimination- warm with HA with hydroxide ions, dissolved in ethanol 

Ozone- Naturally formed in upper atmosphere = beneficial absorbs UV radiation

CFC's = unreactive, non-flammable, non-toxic used fire extingusers. Research shows do damage to ozone layer = banned.  HCFC's and HCF's used as temp alternatives 

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Halogenakanes

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Alkenes 3.3.4.1

Unsaturated Hydrocarbons- Not maximum amount of H and have a double bond. THIS bond is a double covalent bond high e density 

Bromine Test- Shake alkene with orange Br H2O solution goes from orange to colourless due to double bond 

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Alcohols 3.3.5.1

alcohols produced industriallly by hydration of alkenes in presence of acid catalyst 

  • steam hydration of ethene (from cracking fractions of crude oil) used to produce ethanol. needs solid phosphoric acid catalyst (reactions reversiable)
  • 300 Degress and 60 atm 

Fermentation of glucose 

  • yeast produces enzymes which convert glucose into ethanol and CO2
  • solution = 15% ethanol, yeast dies. FD used to increase concentration of the ethanol ( needs purified) 

(table comparing 2 on seperate sheet)

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Additon Polymers 3.3.4.3

polymers- long chain of monomers joined together 

Additon polymers are formed from alkenes and substitued alkenes( swamp H atom for chlorine)

Additon P = UNreactive - main carbon chain usually non-polar 

monomers within chain = stong CB however, VDW's = intermolecular forces between polymer chains 

Poly(chloroethene)

  • c-cl = polar 
  • permenant dipole dipole forces in polymer chain
  • PVC = hard brittle 
  • ADD chemicals = plasticers to polymers to modify 
  • makes P bendier. P get in between polymer chain and push apart = chains slide = more flexible 
  • Plasticised PVC can be used as electrical cable instulation, flooring tiles 
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Alcohols- Biofuels 3.3.5.1

Biofuel- fuel made from biological materials that have recently died e.g sugar from sugar chain fermented to produce ethanol

Advantages-

  • From renewable energy sources 
  • more sustainable 
  • burning biofuel only releases same amount of C02 that the plant took in as it was growing = carbon neutral

Disadvantages- 

  • 'food vs fuel' debate 
  • deforestation to secure land = destroy habitat Trees cut usually burnt = more co2 released 
  • fertillisers added = pollute waterways, release nitrous oxide 
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Oxidising alcohols 3.3.5.2

1 = oxidised, distillation, heat K2Cr2O7 solution and sulfuric acid in test tube = aldehyde 

oxidise (excess oxidising agent) aldehyde under reflux in acidic conditions (R allows to increase temp of reaction to boiling without losing volatile solvents, reactants or products) = carboxylic acid 

2= reflux secondary A with acidified dichromate = ketone 

3= DONT react with acidifed potassium dichromate as dont have spare hydrogen on the functional group to react it with. Only way to oxidise is to burn it 

  • Fehlings solution 
  • Blue Cu 2+ complexes of copper ii sulfate. = STAYS with Ketone.     Brick red Cu2O precipate when warmed with Aldehyde. 
  • Tollens' Reagent 
  • Colourless (Ag(NH3)2) + Complex.  = Silver coats inside apparatus = silver mirror with Ketone.                 Silver when warmed with Aldehyde.       
  • Test for Carboxylic Acids
  • limewater - co2 bubbled through, calcium carbonate precipate formed = LW Cloudy
  • Alkenes- bromine water orange to colourless if -ene present 
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Organic Analysis 3.3.6.1

Alkenes can be formed from alchols by acid catalysed elimination (dehydration) reactions 

Test for primary secondary and tertiary alcohols 

  • add acidifed potassium dichromate solution to test tube 
  • warm mixture in water bath 
  • primary and secondary = colour foes from Orange TO green 
  • Tertiary = STAYS Orange 

Aldehydes and Ketones- Fehlings and Tollens's 

Alkanes, Halogenoalkanes, Cycloalkanes, branched alkanes, alkenes- Bromine Water 

Carboxylic Acids- Limewater 

Alcohols- Add, sodium metal to it and if A it will fizz = H2 Gas

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Organic Analysis 3.3.6.2-3

Mass Spectrometer can be used to work out molecular formula- M/Z value = Mr value so calculate the Mr of first  few e.g alkanes to work out which Mr would go with which molecular formula 

High resolution mass spectrometry- measure atomic and molecular mass accurately = identify compounds that have same Mr to nearest whole number 

Infared spectroscopy

  •  CB in molecules absorb IR radiation = increase vibrational energy 
  • bonds between different atoms and in different places absorb different frequencies of IR radiation = different wavenumbers
  • fingerprint region 1000-1550 allows identification of a molecule by comparison of spectra 
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Organic Analysis 3.3.6.2-3

Mass Spectrometer can be used to work out molecular formula- M/Z value = Mr value so calculate the Mr of first  few e.g alkanes to work out which Mr would go with which molecular formula 

High resolution mass spectrometry- measure atomic and molecular mass accurately = identify compounds that have same Mr to nearest whole number 

Infared spectroscopy

  •  CB in molecules absorb IR radiation = increase vibrational energy 
  • bonds between different atoms and in different places absorb different frequencies of IR radiation = different wavenumbers
  • fingerprint region 1000-1550 allows identification of a molecule by comparison of spectra 
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