Chemistry

Standard conditions for cell potential stuff

• An ion concentration of 1.00 mol dm^-3

• A temperature of 298 K

• A pressure of 1 atm

Standard eletrode potential def

• Potential difference when a half-cell is connected to a standard hydrogen electrode under standard conditions.

Standard cell potential def

• Potential difference between 2 half-cells under standard conditions of 1 atm, 298k, and all solutions at 1.00 mol/dm³

Lattice energy

• Enthalpy change when 1 mol of an ionic compound is formed from its gaseous ions under standard conditions

• Exothermic

Atomisation

• Enthalpy change when 1 mol of gaseous atoms is formed from its element under standard conditions

• Endothermic

Electron affinity

• Enthalpy change when 1 mol of electrons is added to 1 mol of gaseous atoms to form 1 mol of gaseous negative ions under standard conditions

• First exothermic, then endo

Enthalpy of solution

• Enthalpy change when 1 mol of an ionic substance dissolves in sufficient H2O forming a very dilute solution under standard conditions

• can be exothermic or endothermic

Enthalpy of hydration

• Enthalpy change when 1 mol of specified gaseous ion dissolves in sufficient water forming a very dilute solution under standard conditions.

• Exothermic

• Hsol - Hlattice dissociation = H hydration

Entropy

• Number of possible arrangements for the particles and its energy in a given system.

Rate of reaction

• Change in concentration of a reactant or product per unit time

Rate equation

• K only affected by temperature

• Order only found experimentally

Half life

• Time taken for concentration of limiting reactant to half

• 0.693/k

Order of reaction from half life

• Zero order = Half lives decrease, takes less time for reactant conc to half

• First order = half life constant

• Second order = half life increases with time, takes more time for reactant conc to half

Mode of action of heterogeneous catalysts

• Adsorption of reactants on catalyst surface

• Reactant physically adsorbed onto surface by weak forces

• Chemically adsorbed by stronger bonds

• Causes bonds to weaken in atoms of reactants

• Desorption, bonds between products and catalyst weaken, products break away.

Iron in haber process

• N & H diffuse onto iron surface

• Adsorbed onto iron surface due to bonds between iron and reactants

• Those bonds weaken N-N & H-H bonds

• N and H reacnt = NH3

• Desorption, NH3 & iron bonds break and NH3 diffuse away

Catalytic converters

• Removes oxides of nitrogen and CO, NOx → N2, CO → CO2

• Honeycomb, small beads, platnium palladium rhodium

• Adsorption of NOx & CO

• Weakening of NO and CO bonds

• Bonds form between adjacent N-N

• CO and O form CO2

• CO2 and N2 desorption, diffuse away.

Homogeneous, peroxydisulfate - iodine reaction

• S2O8 + 2I- → 2SO4(2-) + I2

• Negative reactants = repulsive = lots of energy

• Fe3+ reduced to Fe2+ by I-

• Fe2+ oxidised back to Fe3+ by S2O8(2-)

• No repulsive forces by reacting wt +ve Fe ion, Ea lowered.

NO2 as catalyst

• SO3 + H2O → H2SO4

• NO2 + SO2 → SO3 + NO

• NO + 1/2 O2 → NO2

Nitrates & Carbonates thermal stability trend

• More thermally stable down group

• Lower charge density = less polarisation = keeps No3/Co3 bond stable/doesnt distort its bonds.

  • Polarisation = drawing electrons towards itself.

Solubility of sulfates and hydroxides

• Solubility of OH- increases down group

• Solubility of SO4(2-) decreases down group

• BaSO4 completely insoluble

Enthalpy of lattice energy

• ΔHlatt decreases down group

• Cations larger down group = more space between +ve and -ve ions

• Weaker attractive forces

• Less energy released when formed from gaseous ions

• ΔHlatt becomes less exothermic

• For explaining solubility down group 2
  
  

Enthalpy of hydration

• Decreases down group 2

• Larger cations = weaker ion dipoles

• Less energy released when gaseous group 2 ions hydrated

• ΔHhyd less exothermic

ΔHsol of hydroxides

• OH- = small

• ΔHlatt falls faster than ΔHhyd

• ΔHsol more exothermic down G2

ΔHsol of SO42-

• SO42- = large ion

• ΔHlatt falls slower than ΔHhyd

• ΔHsol becomes less exothermic down group

• More exothermic ΔHsol = more soluble

Characteristics of benzene

• Has delocalised e-/delocalised pi bond, gets attacked by electrophiles.

• Conjugated-alternating double bonds & single bonds.

• Resonance hybrid, similar to no2 co3- no3-

• No sp3 hybridisation/all carbons sp2 hybridised.

• All c-c bonds are same length ‘

• c-h bonds are s-sp2 hybridised

• c-c bonds have sp2-sp2 overlap and p-p overlap.

• Planar

• aromatic

• Prefers electrophilic reactions

• Decolourises bromine water

• Has pi electrons above and below ring.

Conditions for being aromatic

• Cyclic hydrocarbon

• Planar molecule

• Conjugated

• No sp3 bonds

• Obeys hekules rule - 4n + 2 (n = no of pi electrons)

Benzene nomenclature

Benzene vs intermediate

Nitration of benzene

• Reagents - Conc HNO3 & conc H2SO4

• At 25 to 60 C

Friedel-crafts alkylation

• Reagents - ROCl & dry alcl3

• Heat

Halogenation of benzene

• Reagents - Cl2/Br2 & AlCl3/AlBr3/FeCl3/FeBr3

Oxidation of benzene side chain

• Reagents - Hot conc acidified KMno4 or alkaline KMno4 then acidification

Reduction of benzene

• Reagents - Pt/Ni and heat

Methyl benzene reaction wt chlorine

Characteristics of substituent groups

Phenol characteristics

• 8 delocalised electrons

• Oxygen in OH group has lone pair delocalised into benzene ring (O activates ring)

• Phenols more reactive/susceptible to electrophiles bc lone pair of Ox bonded wt benzene ring, increases e- density and means reactions need less conditions to occur.

• C=O bond gets stronger (wont react with PCl5)

• Donation of e- by the oxygen outweighs it e- pulling effect.

Electrophilic substitution of phenols

• Phenol + Dil HNO3 → 4 - nitrophenol

• Phenol + conc HNO3 → 2,4,6 - trinitrophenol

• Phenol + 3Br2 - > 2,4,6 tribromophenol + 3HBr, white ppt

• Phenol + 3Cl2 → 2,4,6 trichlorophenol + 3HCl, white ppt

Acid nature of phenols

• Acidic bc react wt water to give H3O+, phenol = weak acid.

• Decreasing acidity - Phenol > water > alcohol.

• Decreasing stability - Phenoxide > hydroxide > alcoxide.

• Only COOH reacts with Na2CO3.

• C6H6 + NaOH → C6H6O-Na+ + H2O

Esterification of phenol

• Phenol + Ch3COCl wt naoh → phenyl ethanote + HCl

Oxidation of carboxylic acids special

• HCOOH/HOOCCOOH + [O] → CO2 + H2O

• Warm tollens reagent silver mirror

• Warm Fehling’s solution, brick red

• H+/K2Cr2O7 + heat, orange to green

• H+/KMnO4 + heat, purple to colourless

Formation of acyl chlorides

• RCOOH + PCl3 + heat → RCOCl + H3PO3 white fumes

• RCOOH + PCl5 → ROCl + POCl3 + HCl(g)

• RCOOH + SOCl2 + heat → RCOCl + SO2 + HCl

• Reverse is hydrolysis

• easiest to hydrolyse

• RCOCl > chloroalkane > chlorobenzene

Acyl chloride to COOH draw

• 
  

Esterification acyl chloride

• RCOCl + ROH → RCOOR + HCL

• RCOCL + phenol dissolved in naoh → RCOOR + HCL

nucleophile acyl chloride

• RCOCL + nuc-h → RCOnuc + HCL

• RCOCL + NH3 → RCONH2 + HCL

Amide formation

• RCOCL + R’NH2 → RCONHR’ + HCL

Amine info

• All react similar to ammonia due to lone pair, lp makes ammonia reactive and:
  Act as base
  Act as nucleophile(donate lp)

• Trigonal pyramidal shape 107 degrees.

• Forms H-bonds wt another amines, high bp.

Formation of amines

• Nuc sub - Rx + NH3 → RNH2 + HX

• Reduce - RCN + 4(H) → RCH2NH2

• H = h2 & Ni + heat or LiAlH4 in dry ether + heat

Reduction of amides

• CONH + [H] → CH2NH

• RCONH2 + 4[H] → RCH2NH2 + H2O

Phenylamines formation

• Reduction

• Nitrobenzene + Sn + conc hcl and heat → Phenylamine

Amines as nucleophiles

• RCOCL + R’NH2 → RCONHR’ + HCl

Amines acting as base

• RCH2NH2 + H+ → RCH2NH3+

Basicity of amines

• RNH2 > NH3 > phenylamine

• Depends on how readily the amine donates its lp.

• R group is e-donating and inductive effect

Phenylamine reactions

• Phenylamine + 3Br2 → 2,4,6 tribromophenylamine

• Phenylamine + NaNO2 at 5 C + dil hcl → diazonium salt + heat + h2o → phenol

Coupling reaction

Common ion effect

• A salt is less soluble in a solution which already contains the ions present in the salt.

• If a solution containing one of the ions of the salt is added to a saturated solution, some of the salt will precipitate

Colour

• Visible light hits transition metal ion, electrons excited to higher energy levels.

What affects delta E

High oxidation state = high E

More repulsion between ligands = high E

Copper and chromium electron configs

• Cr = [Ar] 4s1 3d5

• Cu = [Ar} 4s1 3d10

How to draw ion config transition

• Take out of 4s orbital FIRST before taking out of 3D because filled 4s is higher in energy than 3d.

Why variable oxidation states

• Due to small energy difference between 4s and 3d orbitals