Chemistry Unit 1

Chemistry Unit 1

1. Orbitals

   • shell= energy level= quantum number
   • Aufbau principle
     • lowest first
   • Pauli exclusion principle
     • Max 2 electrons
   • Hund's rule
     • occupy separate orbitals before sharing
   • opposite spins to minimise repulsion
   • 1s(2), 2s (2), 2 P (6), 3s (2), 3P (6), 4S (2) 3D (10), 4 P(6) 5S(2)
     • First number= energy level
   • P orbitals - x,y and z axxis
   • Line emission spectra- lines, not continuous- electrons must have fixed energy levels (only certain quanta of energy released between levels)

2. Periodicity

   • Ionisation energy
     • 1st= one mole singly charged gaseous cations
     • large decrease= change in principle quatum number
     • Trend down group= decrease
       • outer electrons further away
         • Electrostatic attraction between nucleus and electron weaker
       • Protons and shielding balance
     • Trend across period= increase
       • Increased nuclear charge
       • Shielding same
     • Group 3= less energy for p orbital than s
     • group 6= electrons paired in orbital= repulsion= easier
     • E(g) ==> E+(g) + e-
     • Graphs- lowest point ( on left) outermost electron- removed first
     • Successive I.E.s evidence for quantum shells & 1st I.E. of successive elements evidence for subshells
     • Group reactions faster than group 2- less energy to remove just one (think maxwell-boltzmann)
   • Electron affinity
     • 1st= energy released for one mole of gasous 1- anions.
     • Same trends as ionisation energy
   • Metallic properties
     • Decreases across period
       • Harder to remove electrons (smaller radius)
   • Atomic radius
     • down groups= increases
       • across periods= decreases
         • More protons with same shielding
     • Not fixed- found by halving distance between 2 nuclei
   • Melting point
     • Highest in centre of table
     • noble gases lowest
   • Valency
     • number electrons used in bonding
     • 2 unused valence electrons= lone pair

3. Bonding

   • Metallic
     • Depends on= ion charge and size
       • High MP: high charge & small ion
     • Diagrams: make sure no. electrons equal to total positive ion charge
       • Lattice
   • Ionic
     • Ionic radius
       • cations= smaller anions= larger
         • Anions more electrons= more shielding= weaker electrostatic force of attraction
       • Increases down a group
       • Decreases across set of isoelectronic ions e.g. N 3- to Al 3+ as proton number increases but no. electrons stays the same
     • Trend in ionic radii= down group increases
       • Across isoelectronic radii decreases
     • Properties/ evidence for existence
       • High MPs
         • Strong electrostatic attractions
       • Hard and brittle
         • Evidence for lattice formation- when layers more and similar ions next to each other- repulsion- split/ fracture crystal
       • Conduct electricity when molten
         • When solid ions in lattice & not free to move
       • Soluble in polar solvents
         • Stabilizes the separated ion
       • Electron density maps
       • Migration of ions e.g. electrolysis
         • e.g. aqueous copper chromate (VI)- yellow chromate to positive anode and blue Cu2 ions to negative cathode
         • Potassium permanganate: potassium to cathode, manganese + oxygen to anode (see purple)
         • Copper chloride: copper to cathode (see green), chloride to anode
     • Giant Lattice
     • Coord number= number of + ions surrounding - ion (or oppsoite)
     • Polarisation= distortion of electron cloud
       • Cation distorts anion
         • Effects properties: -not crystalline -lowers MP -poor water solubility -may react w/ water
         • Cation has polarizing power
         • Anion is polarizable
         • Changing anion size: large radius- easily polarized due to electron shielding- greater covalent nature- greater differene between theoretical and experimental values (Born- haber cycles). Vice versa for cation (and high charge also increase polarizing power)
           • When become more electron sharing than electron transfer- covalent bonding
             • Seen on electron density maps
       • Calculation of attractive force based on assumption that ions carry full charge+ distance between then is sum of ionic radii
     • Equations
       • Make sure: balanced, no spectator ions
   • Covalent
     • 2 s orbitals or 1 s and 1 p= sigma bond
       • 2 p orbitals overlapping= pi bond
         • If atoms too close nuclei+ inner electrons repel other atoms'= bond has certain length
     • Electron spin must be opposite
     • Down group radii increases- bond weaker
       • As bond length increases strength decreases
     • In double bond second not as strong as first (overlap further from nuclei)
       • So double bond breaking doesn't need double energy
     • Dative (both electrons from same atom)
       • A.K.A. co-ordinate bond
       • Be careful when calculating charge of molecule
     • Simple molecular/ lattice
       • Low MPs & BPs
         • Only have to break weak IMFs
       • Don't conduct electricity
       • Soluble
     • Giant covalent struture
       • High MPs & BPs
       • Don't conduct electircity
       • Insoluble
       • Allotropes e.g. graphite and diamond
         • Graphite- 1 free electron- create IDIDs. Lower density than diamond
         • Buckminsterfullerene- 60 carbon atoms, each bonded to 3 others (2 single and 1 double bond)
         • Carbon nanotubes- components in electrical transistors
   • Electron deficient

4. Calculations

   • Empirical formula
     • Element, Mass or %, Atomic mass, Mass/Ar, /by smallest, Ratio
   • Moles
     • Avagadro's constant
     • moles= mass/mr
     • No. particles= no. moles * A.N.
     • ! mole of gas (room temp. and pressure)= 24 dm cubed
     • moles= volume (cm)/ 24000 (volume of one mole)
     • moles= concentration * volume (/ 1000)
   • ppm= parts per million (by mass)
   • mol (dm to minus three) * 1000= milli mol
     • milli mol * 1000= micro mol
       • micro mol * 1000= nano mol
   • Atom economy= percentage total mass of reactants that end up useful product
     • Assuming 100% yield
       • useful product/ total reactant * 100
     • Taking yield in account
       • % yield/ total reactant * useful product
   • Ionic equations
     • 1) Balanced equation 2) Split into ions 3) Cut out spectator ions
       • Remember state symbols

5. Energetics

   • Q=m*c*change in T
     • Assume: c= 4.18 J/g/C, 1g water= 1 cube cm, no heat lost to surroundings
   • Definitions for standard enthalpies of reactions/ formations/ combustion/ neutralisation/ atomization
     • Remember: one mole, standard conditions (100 kPa and 298 K)
   • Hess's law
     • Total enthalpy change is independent of route taken
     • Combustion- arrows downwards
     • Formation- arrows upwards
     • Remember: multiply energy by no. moles
   • Exothermic
   • Endothermic
   • Enthalpy profile diagrams
   • Thermochemical equations
   • calorimetry
     • if different concentrations/ volumes: find no. moles & discard higher moles
   • Bond enthalpies
   • Born- Haber cycles
     • Lattice enthalpy
       • Affeted by : charge on ion & ionic radii
     • -Formation- route 1 -Atomisation (both) + Ionisation (I.E. and E.A.) - route 2
     • Assume: compound fully ionic= theoretical values
       • Close to true value: compound is fully ionic
       • Different to true value: partial covalent nature
       • Theoretical values calculated from electrostatic theory
     • Limitations: based on average bond enthalpies (changes depending on surrounding atoms) and only gaseous species
   • Thermometric titrations for enthalpy neutralisation
     • If given temp. acid and base separetely find average for solution
     • Use moles of acid in calculations (not in excess)
     • Slow reactions: plot temp. change and extrapolate
       • Remember: earlier results more accurate (less heat loss + heat spread over smaller volume)

6. Organic chemistry

   • Alkanes
     • Chlorination
       • Conditions: UV light, excess alkane (avoids di- tri- chloro...)
       • Overall: CH4+ Cl2 ==> CH3Cl+HCl
       • Initiation: homolytic fission
         • Remember: curly half arrows +dot for free radicals
       • Propagation: CH4+Cl* --> *ClH3+HCl
         • *CH3+Cl2--> CH3Cl+Cl*
       • Minor termination: *CH3+*CH3--> Ch3CH3
       • Termination: CH3+Cl*-->CH3Cl
       • Further substitution: CH3Cl+Cl2--> CH2CL2+HCl
         • CH2Cl2+Cl2--> ChCl3+HCl
           • ChCl3+Cl2--> CCl4+HCl
         • Condtions: UV light, excess chlorine
       • (could be done with others e.g. flourination)
       • Form of free radical substitution
     • Unreative compounds
       • bonds non-polar= strong
     • Combustion in car makes NO (needs high temp. & N froma air)
     • Hydrocarbons
       • Aliphatic
       • Alicyclics (closed ring)
       • Arenes (ring stabilised by delocalisation of electrons
     • Crude oil
       • Fractional distillation
         • Refinery gas
           • Lightests, gaseous alknaes (1-4 carbons)
         • Gasoline
           • Knocking
             • Explosion in car engine too early- lose power
             • Too many straight chain alkanes
         • Kerosene
           • C11 and C12
           • Aircraft engines
         • Diesel/ gas oil
         • Residue
           • V. viscous
           • Fuel for power stations
           • further fractioning in a vacuum (avoid high temp.)- bitumen to surface roads
       • Catalytic cracking
         • Zeolite catalyst
       • Octane rating
         • Ratio bracnhed to straight (bracnhed combustion more controllable)
         • ignition tested aganst known one
       • Catalytic reforming
         • Make branched hydrocarbons and arenes
         • Platinum catalyst
     • Insoluble
   • Alkenes
     • Oxidation
       • with: acidified Potassium manganate
         • Solution purple --> colourless
       • Makes a di....ol
     • Electrophilic addition
       • With halogens
         • Creates di....
         • dipole induced in halogen molecule
         • e.g. Br2
           • Heterolytic fission
           • Test for alkenes
       • with hydrogen halides
         • Markovnikov's rule (hydrogen to carbon with most Hs already)
           • Creates most stable compound
         • Electrophile
       • Bromine water
         • Test for alkenes
         • Makes bromo- alcohol (addition of OH and BR)
     • Addition
       • with hydrogen
       • Nickel catalyst (at 150-200 degrees)
       • Platinum catalyst at room temp.
       • Both reacting species absorbed onto catalyst (used to bring them together- NOT electrostatic attraction)
     • Polymerisation
       • Opens double bond
       • Started by free radical (ends with one on each side)
       • LDPE
         • High pressure- 1400 atm
         • 170 degrees C
         • Lots of branching- not as compact
       • HDPE
         • Lower pressure- 2 atm
         • 70 degrees C
         • Little branching- tightly packed
       • Poly-tetra-flouroethene a.k.a. teflon (cooking +clothes coatng
         • F v. electronegative- electrons can't be moved- low friction
     • Carbo cations
       • Higher= more stable (molecule's + charge slightly less)
     • Uses electrons from pi bond
       • Say 'electron pair in double bond attack +ve side...)
   • Isomers
     • Stereo-isomerism
       • Geometric
         • Entgegen/ zusammen
           • Different groups: look at atomic no. of atoms attached to C=C (nothing extra)
           • Not alkanes- single bond can rotate
       • Optical (non super-imposable mirror images)
     • Structural
       • Chain (longest one different)
       • Positional (functional group of different carbon)
       • Functional group (different)
   • General formula
     • Alkane: CnH2n+2
     • Alkene: CnH2n
     • Alcohol: CnH(2n+2)O
     • Carboxylic acids: CnH2nO2
   • Homologous series
   • Functional group
     • Alcohol: C-O-H
     • Aldehyde: double C=O bond and C-H bond (same carbon)
     • Ketone:C=O
     • Carboxylic acid: C=O and C-O-H (same carbon)
   • Hazard and risk
     • Risk assessments
     • Reduce risk
       • Smaller scale
       • Specific precautions/ alternative techniques e.g. using a lower concentration
       • Alternative method with less hazardous substances

7. Mass spectrospocy

   • Vaporisation
     • Heated
   • Ionisation
     • Bombard with high speed electrons- knocks electrons off (too fast to stick)
   • Acceleration
     • Negatively charged plates
   • Deflection
     • Electromagnet- various strengths
     • Lightest deflected most
   • Detection
     • AS electrons move- induce electric current
     • Then amplifier then chart recorder
   • Vacuum pump- gets rid of impurities to avoid collisions with sample
   • Uses:radioactive dating, in sport to detect anabolic steroids, in pharmaceutical industry to identify chemicals and impurities in drugs
     • Biologists can identify proteins etc. e.g. to understand HIV acts in a host cell
       • Anesthesiologists can measure metabolic gas exchange of patient during surgery- can tell if cells getting enough oxygen/ getting rid of CO2

8. Reactions

   • Acids
     • metal + acid ==> salt + hydrogen
     • Metal oxide (base) + acid ==> salt + water
     • Metal carbonate + acid ==> salt + water + CO2
     • Metal hydroxide + acid ==> salt + water
   • Halides
     • Add HNO3+ AgNO3
       • Chloride: White PPt (dissolves in excess aqueous ammonia)
       • Bromide: cream PPt dissolves in conc. ammonia
       • Iodide: yellow ppt insoluble in ammonia
     • Displacement
       • Cl2> Br2 > I2 as oxidising agent
       • e.g. Cl2 + 2I- ==> 2Cl- + I2 (all aq)
     • Add H2SO4
       • Chloride: steamy fumes of HCL
       • Bromide: steamy fumes HBr, brown vapour Br2
       • Iodide: steamy fumes HI, purple vapour and dark solid I2