Unit 4 A2 Chemistry Edexcel

1. +
Chemistry
Unit 4

2. + Reaction Rates
It‟s the change in the amount of reactants or products per unit time
1st Order:
• Rate = k[X]1
• Constant Half Lives
2nd Order:
• Rate = k[Y]2
• Increasing Half Lives
0 Order:
• Rate = k[Z]0
Continuous Rate Method
Initial Rate Method
1o Halogenoalkane + Hydroxide = SN2
3o Halogenoalkane + Hydroxide = SN1
If it‟s in the Rate
Equation then it‟s the
Rate Determining Step

3. + Nucleophilic Substitution
SN2
SN1
Iodine + Propanone
Slow Fast
Fast
Fast Fast
Slow

4. + Arrhenius Equation
ln(k) = X + Constant
-Ea 1
R T
R = Gas Constant (8.314)
T = Temperature (Kelvin)
1/temp (k)
ln( 1/time)
x
y
Gradient =
X
Y

5. + Heterogeneous Catalysts
• In a different state to the reactants
• Large surface area as they‟re usually powder or a mesh
• Easily separated from products & excess reactants
• Can be poisoned:
• Adsorbs too strongly to surface of catalyst and doesn‟t allow other
reactants to adsorb to the catalyst
• e.g. – Nickel in Hydrogenation of Vegetable Oil
– Platinum in catalytic converters in cars
How do they work?
- Reactant adsorbed onto surface of catalysts at the active site
- Interaction between reactant & catalyst
- Reaction occurs from the interaction
- Products are desorbed – breaks off catalyst
Adsorb – forms a
temporary bond when
something sticks to a
surface
Homogeneous Catalyst is
when catalyst is in the same
state to reactants

6. + Entropy
Entropy change of a reaction is measure of order or disorder
The order within is a substance is how the quanta of energy are arranged
Reaction will occur if overall entropy is increasing, from order to disorder
If entropy is +ve then reaction will tend to occur
Ordered
Disordered
But doesn‟t exist
More disorder = more +ve SΘ
Solid  Ordered
Liquid  Disordered
Gas  Very Disordered
More Complex/Moles  More Disordered

7. +

8. + Stable or Inert
Stable if there is no tendency for the reaction to „go‟
If system is +ve but surrounding are –ve we can „force‟ the reaction by
increasing the temp
C (Diamond)  C (Graphite)
+ve entropy change
But Ea is too high – Kinetically Inert
= - ve

9. + Solubility
Endo
Endo
Exo
Standard enthalpy change of Solution:
• Enthalpy change when one mole of a compound
dissolves to form solution containing 1moldm-3 under
standard conditions.
Standard enthalpy change of Hydration:
• Enthalpy change when one mole of gaseous ions is
hydrated under standard conditions to form a solution in
which the concentration of ions is 1moldm-3.
Factors affecting solubility:
• Ionic Charge > Ionic Radii
Lattice Energy & Hydration Energy increase as Charge increases
Group 2 Compounds less soluble as solution energy is more endo
Group 1 Compounds more soluble as solution energy is less endo

10. + Equilibria (Kc)
Dynamic Equilibria:
- The forwards reaction and back reaction are at the same rate so
there‟s no overall change in yield of products or reactants in a closed system

11. + Kp

12. + Effects on Equilibria
Adding a catalyst:
• Equilibrium constants not affected
• Position of equilibria not affected
• Speeds up forward & backward reaction at same rate
Change in concentration:
• Equilibrium constants not affected
• Adding reactant shifts equilibria right
• Adding product shifts equilibria left
Change in pressure:
• Equilibria shifts to side with fewest molecules
• Equilibrium constants not affected
Increase in temperature:
• Endothermic = +ve shifts right, more product produced
• Exothermic = -ve shifts left, more reactant produced
• Kc & Kp:
• Increases if endothermic
• Decreases if exothermic
• Total entropy change = RlnK
(R = Gas Constant)
K
Total
Entropy
Progression of Reaction
>10-10 >-191 Doesn‟t Go
>10-5 >-96 Reversible pushed to Left
1 0 Equilibrium
<105 <+96 Reversible pushed to Right
<1010 <+191 Goes to Completion

13. + Uses in Industry
They alter conditions to produce maximum yield
Requiring least amount of energy
Often looking for new more environmentally friendly catalysts
e.g. Ethene + H2O  Ethanol
Sped up by using catalyst (Silica soaked in H3PO4)
Remove product as it‟s being formed

14. +
Arrhenius:
- Acids are H+ producers
- Bases are OH- producers in H2O
- Only used in aqueous solutions
Bronsted-Lowry:
- Acids are proton donors
- Bases are proton acceptors
Acid Theories
A base has a lone pair of
electrons which can form a
dative covalent bond with a H+
Lewis:
- Acid is an electron pair acceptor
- Bases is an electron pair donor
HA + H2O H3O+ + A-
Conjugate Pairs
H2O & HA are Acids & Bases as they give & accept H+
H+ = H3O+
HA is the acid & A- is it‟s conjugate base
H2O is the base & H3O+ is it‟s conjugate acid
Amphoteric Substances:
• It acts a base or an acid
H3O+  H2O  OH-
AcidBase

15. + pH Strong Acids or Strong Bases
pH = - log 10 [ H+ ] Kw @298k = 1x10
-14
[ H+ ] = 10-pH

16. + pH Strong Acid & Strong Base
Excess [ H+ ]
Excess [ OH- ]

17. + pH Weak Acids
pKa = -log10( Ka )
Larger pKa = Weaker Acid

18. + pH Buffers
Substances that resist change to pH when small amounts of acid/alkali are added
Assumptions:
• [CH3COO-] = [H+]
• [CH3COO-] = [CH3COONa]
• [CH3COOH] same at equilibrium &
at start
[CH3COO-] from Salt
Acidic Buffers = Weak Acid + it‟s Salt

19. + Isomerism/Optical Isomers
Only occurs when chiral carbon present:
- Carbon with 4 different groups attached
Mirror
Enantiomers/Optical Isomers
Racemic Mixture:
Contains equal amounts of each enantiomer

20. + Carbonyls
Ethanal PropanoneCarbonyl Group

21. + Reactions of Carbonyls
Reaction with Dichromate:
- Aldehydes can be oxidised
- Orange  Green
- Ketones can‟t be oxidised
3RCHO + Cr2O7
- + 8H+  3RCOOH + 2Cr3+ + 4H2O
Reaction with Tollens:
- Aldehydes +ve Silver Mirror Forms
Reaction with Benedicts:
- Aldehydes + ve Blue (Cu2+)  Red Precipitate (Cu+)
Reaction with Brady‟s:
(2,4 DNP or 2,4 DiNitroPhenylhydrazine)
- Carbonyls +ve Orange Precipitate
Reaction with Iodine:
- Methyl group adjacent to C=O +ve Pale Yellow Precipitate, Antiseptic smell
Melting point used to identify Carbonyl compound
Dissolved
in
Methanol
& conc
H2SO4
AgNO3
dissolved
in NH3(aq)
Dissolved
in NaOH
dil H2SO4
LiAlH4 (in Dry Ether) to go
from Carb Acid to 1o or
Aldehyde
Presence
of Alkali
Triiodomethane

22. + HCN Reactions
Propanone + HCN
Ethanal + HCN
2 methyl 2 hydroxypropanenitrile
2 hydroxypropanenitrile
Nucleophilic Addition
In a lab HCN made by reacting KCN(s) + H2SO4

23. + Carboxylic Acids
H-Bonding in Pure Ethanoic Acid (Dimer Shape)
Sodium
Ethanoate
Identifying Carboxylic Acids:
Add Sodium Carbonate – effervescence if +ve
Weak acid –
Partially Dissociate
Very soluble in H2O
Longer chain, less soluble
Formed from:
- Oxidising a 1°
- Hydrolysis of Nitrile (Reflux
with dil HCl and distil off)
Reaction with PCl5:
CH3COOH + PCl5  POCl3 + HCl + CH3COCl
Ethanoyl Chloride
Add NH3 white
smoke = +ve
Acid + Alcohol  Ester + H2O

24. + Esters
Acid Hydrolysis:
- Reflux with dil HCl or H2SO4
Transesterification
Base Hydrolysis:
- Reflux with dil Alkali (e.g. NaOH)
Dicarboxylic Acid + Diol Alcohol  Polyester

25. + Acyl Chlorides Ethanoyl Chloride
Reaction with H2O:
CH3COCl(l) + H2O(l)  CH3COOH(l) + HCl(g)
Reaction with Alcohol:
CH3COCl(l) + CH3CH2OH(l)  CH3COOCH2CH3(l) + HCl(g)
Ethyl Ethanoate
Reaction with Ammonia:
CH3COCl(l) + NH3(aq)  CH3CONH2(aq) + HCl(g)
Ethanamide
Reaction with Ethylamine:
CH3COCl(l) + C2H5NH2(aq)  CH3CONHCH3CH2(aq) + HCl(g)
N-Ethyl-Ethanamide
N-“substituted”-Amide
-amide

26. + Soap & Triglycerides
Fats solid at RTP
Triglycerides have lower melting point due to less regular shape
Hydrogenation:
- Nickel catalyst @ 150°C
- Unsaturated  Saturated FA
- Solidifies fats

27. + UV & Microwave Radiation
UV can initiate reactions:
• In the form of electromagnetic radiation
• Wavelength between Visible and X-ray (400nm-10nm)
• e.g. Cl-Cl = Clo + Clo homolytic fission
Free-radical Substitution:
• Initiation = Breaks homolytically (sufficient energy in sunlight)
• Propagation = Cl* + CH4 = HCl + CH3*
= CH3* + Cl2 = ClCH3 + Cl*
• Termination = Cl* + Cl* = Cl2
= CH3* + Cl* = CH3Cl*
= CH3* + CH3* = C2H6
Microwaves are used to heat things
• Longer than IR also used for communications (1mm-1m)
• Most molecules in food are polar (water, fats, sugars)
• Microwaves are passed through, they create an electric field
• Any polar molecules align with the electric field
• Whilst rotating they collide with other molecules generating heat energy
Exo

28. + Mass Spectroscopy
Vapourisation:
Sample put into vacuum – analysed as a gas
Ionisation:
High energy e-s knock off other e-s (+vely
charged sample)
Electron gun used
Acceleration:
-vely charged plate pulls sample up the tube
Deflection:
Magnetic field introduced – lighter atoms = deflect
more
All same charge, so only mass varies
Detection:
Atoms hit charged plate – small charge is created
Uses: Drugs Testing &
Carbon dating
Particle Charge Mass
Proton +1 1
Neutron 0 1
Electron -1 1/1840
Parent Ion Peak = Mr
Only Ions show up

29. + NMR – Nuclear Magnetic Resonance
• Any molecules with odd number of nucleons (protons & neutrons) has nuclear
spin
• This causes a weak magnetic field
• NMR looks at how this weak field reacts when you put it in a much larger
external field
• When field is applied the protons align themselves with or against the field
• When aligned protons are at a lower energy level than opposing protons, and
they can absorb radio waves of the right frequency, they flip to a higher energy
level
• The opposing protons can absorb the radio waves and flip to a lower energy
level
• There tends to be more aligned protons, so there‟s an absorption of energy
overall. NMR measures this absorption
• Proton environments can affect the amount of absorption
Chemical Shift:
• Every energy peak is relative to the peak of tetramethylsilane at „0‟
• Height of peak is no of protons
2 Proton
Environments
2 Proton
Environments

30. Chemical Shift

31. + NMR
High Resolution NMR:
• Each peak is broken down into smaller peaks, this is due to neighbouring
magnetic field interacting with each other (spin-spin coupling)
2 Peaks so 2
Proton
Environments
Peak at 9.5ppm
due to R-CHO
(aldehyde)
Peak at 2.5ppm
due to R-COCH3
(carbonyl)
Doublet
Quartet
Tetramethylsilane
Uses:
• To ensure pharmaceutical products are pure
• Studies internal structures of the body

32. + IR – Infrared Spectroscopy
• A beam of IR radiation goes through the sample
• Bonds absorb the IR energy, increasing the vibrational energy
• Different bonds absorb different IR wavelengths
• Bonds in different places of a molecule also absorb different wavelengths
Uses:
• Able to detect when one functional group has been
changed to another in a reaction
• The degree of polymerisation that has occurred
• Detects weaknesses in polymer is reacted with O2

33. + Chromatography
• Mobile Phase – where the molecules can move
• Stationary Phase – where the molecules can‟t move
Gas Chromatography:
• Stationary phase is a viscous liquid (e.g. oil) which coats a coiled tube
• Mobile phase is N2(g) as it‟s unreactive
• Sample is injected into tube as gas
• Each compound adsorbs to the stationary phase differently
• So each compound takes a different amount of time to be recorded (retention
time)
• Recorder produces a graph – area/height shows amount of each compound
HPLC – High Performance Liquid Chromatography:
• Stationary phase is small particles of a solid in a tube
• Liquid mobile phase is usually a polar mixture (e.g. methanol & H2O)
• Liquid is forced through tube under high-pressure
• Sample added to liquid phase and forced through tube as a solution
• Mass spectrometer is used to analyse each compound as it‟s collected
• Mixture separates as it adsorbs to the solid differently
• UV shone through liquid stream at end
• UV absorbed by mixture as it comes through. Graph is produced
• Can be used when sample is heat sensitive or high boiling point
Uses:
• Checks chemical
equipment for
impurities
Uses:
• Routinely check purity
of products in a
continuous process