AQA module 5 - complete syllabus overview

**GJS** · 21-06-2006, 06:33 PM

Hope these notes help

There is an attachment of all the different ligand substitutions/oxidations aswell

Thermodynamics

Enthalpy Change (DH)

Everything is assumed to be under standard conditions in there standard states
298k, 100KPa, 1.00 mol dm^-3

Enthalpy of Formation (DHf)
The enthalpy change when one mole of a substance is formed from its constituent elements

Ionisation Enthalpy (DHi.e)
The enthalpy change when one mole of gaseous atoms form one mole of unipositive ions

Enthalpy of Atomisation (DHat)
The enthalpy change when one mole of gaseous atoms are formed from an element
½ *(DHdiss) = (DHat)

Bond Dissociation Enthalpy (DHdiss)
The enthalpy change when a particular bond in a covalent molecule is broken to make gaseous atoms

Enthalpy of Lattice Formation (DHL)
The enthalpy change when one mole of an ionic solid is formed from its gaseous ions
(DHL) = -ve lattice dissociation

Enthalpy of Hydration (DHhyd)
The enthalpy change when one mole of gaseous ions become completely dissolved in water forming one mole of aqueous ions

Enthalpy of Solution (DHsol)
The enthalpy change when one mole of an ionic solid completely dissolves in water

Born Haber Cycles

There are two routes round a Born Haber Cycle, Route 1 & Route 2
Route 1 = Route 2
As anion size decreases the lattice formation becomes larger
Charge and size effect (DHL)

Enthalpies of Solution

There are two stages 1) lattice dissociation 2) hydration
(DHsol) = (DHLD) + (DHhyd)

Mean Bond Enthalpies

Are averages so inaccurate, taken from a range of different compounds
Energy Change = Sbonds broken - Sbonds made

Entropy

If (DH) is correct then spontaneous endothermic reactions could not occur, this is explained by Entropy
Entropy - measure of disorder
Reactions tend to increase disorder
Entropy - gases> liquids> solids
NaHCO3 + Acid à CO2 the gas increases entropy
An increase in number of molecules also increases entropy e.g a + b à 2c + 2d
Standard Entropy - S - Jk^-1 mol^-1 at 298k, 100KPa
DS = SSproducts - SSreactants
Feasibility - spontaneous reactions
DG = DH - TDS
If DG less than 0 reaction is feasible
It DG = 0 reaction in equilibrium and we can find the minimum temperature T = DH/DS
DG gives no indication of rate

Periodicity

Period 3 Elements Reactions (water/oxygen/chlorine)

2Na + 2H20 à 2NaOH +H2
Mg + H2O à MgO + H2
Mg only reacts with steam, as smaller and higher charge than sodium, so requires higher activation energy

4Na + O2 à 2Na2O
2Mg + O2 à 2MgO
4Al + 3O2 à 2Al2O3
Si + O2 à SiO2
4P +10O2 à P4O10
S + O2 à SO2

2Na + Cl2 à 2NaCl
Mg + Cl2 à MgCl2
2Al + 3Cl2 à 2AlCl3
Si + 2Cl2 à SiCl4
2P + 5Cl2 à 2PCl5

Acid/Base Properties of P3 Oxides

Na2O - Giant Ionic Lattice - Ionic - Basic - pH = 14
MgO - Giant Ionic Lattice - Ionic - Basic - pH = 7/8
Al2O3 - Giant Ionic Lattice - Ionic - Amphoteric
SiO2 - Macromolecular - Covalent - Insoluble - Acidic
P4O10 - Simple Molecular - Covalent - Acidic - pH = 2
SO2 - Simple Molecular - Covalent - Acidic - pH = 3
SO3 - Simple Molecular - Covalent - Acidic - pH = 1

Metal oxides are basic so react with acids to make a salt and water
Covalent oxides are acidic so react with bases to make a salt and water
SiO2 -strong covalent bonds between atoms - non-polar - therefore insoluble

Na2O + H20 à 2NaOH
MgO + H20 à Mg(OH)2
P4O10 + 6H20 à 4H3PO4
SO2 + H20 à H2SO3
SO3 + H20 à H2SO4

Explaining trends in pH
Low electronegativity on X e.g. metal oxide X---O----H à X+ -O----H
High electronegativity on X e.g. non-metal oxide X---O----H à X-----O- H+

Properties of P3 Chlorides

Charge density - ratio of charge to size - increases across the group - this means hydrolysis of water increases across the group as the ions are able to distort water molecules more

NaCl - Ionic - Giant Ionic Lattice - No hydrolysis - pH = 7

MgCl2 - Ionic - Giant Ionic Lattice - Slight hydrolysis - pH = 7/6
[Mg(H20)6]^2+ + H20 « [Mg(H20)5(OH)]^+ + H3O^+ equilibrium lies too the left

AlCl3 - Covalent - Simple Molecular (exists as a dimer) - is covalent as the Al3+ can polarise the Cl- ion - Hydrolysed - pH = 3
[Al(H20)6]^3 + + H20 « [Al(H20)5(OH)]^2+ + H3O^+ equilibrium lies too the right

SiCl4 - Covalent - Simple Molecular - Hydrolysed - forms gel structure first Si(OH)4, then forms a white ppt - pH = 1
SiCl4 + 2H20 à SiO2 4HCl

PCl5 - Covalent - Ionic Structure due too [PCl4]+[PCl6]-, one molecule donates a Cl- - Hydrolysed - Exothermic/Effervescence (gas evolved) - pH = 0
PCl5 + 4H20 à H3PO4 + 5HCl

Non-metal chlorides are simple molecular, only weak van der waal’s forces exist between molecules, this takes little energy to break the intermolecular bonds - resulting in low melting points

Redox Equilibria

Electrode Potentials

Oxidation - loss of electrons
Reduction - gain of electrons
Oxidising Agent - is reduced - oxidises things
Reducing Agent - is oxidised - reduces things

Electrochemical Cells

You cannot measure a single cell you must measure it against another cell
Also you cannot get a true e.m.f as an e.m.f only exists when no current flows
In half reactions with a non-metal we use an inert electrode e.g. Pt
SEP - Standard Electrode Potential - 298k/100KPa/1.00 mol dm^-3
SHE - Standard Hydrogen Electrode
Other electrodes measured against SHE
SEP of SHE = 0.00V
Emf = Eright - Eleft = Ereduction - Eoxidation
Spontaneous reactions have to have a +ve E.m.f
SHE is not used because - the reaction is slow - electrode is cumbersome - difficult to maintain a constant pressure
Therefore we use a secondary standard electrode e.g. calomel - and then calibrate this against the SHE
Factors that effect electrode potential - conc of ions - temp - pressure of gas
Use le chateliers principle to work out what changes in these factors cause on equilibrium

Writing Half Cells
| -represents boundary between two phases
|| - represents salt bridge
Reduction reaction on the right of the cell
Oxidised species next to the salt bridge
e.g. Pt|H2,H+||Cu2+|Cu or Zn|Zn2+||H+,H2|Pt

Electrochemical series - list of SEP’s
Reducing agent on right
Most -ve are the strongest reducing agents

**GJS** · 21-06-2006, 06:33 PM

Transition Metals

General Properties Of TM’s

TM characteristics are due too incomplete d-robitals
These characteristics are - complex formation/ coloured ion formation/ variable oxidation states/ catalytic activity

Complex Formation

Ligand - species which can donate a pair of electrons too a central metal ion (lewis base)
Complex Ion - central metal ion surrounded by ligands which can donate a pair of electrons in a co-ordinate bond
Co-ordination number - number of co-ordinate bonds

Ligands

Unidentate - donates one pair of electrons - e.g. H20, NH3, OH-
Bidentate - donates two pairs of electrons - e.g. C2O2^2-, NH2CH2CH2NH2
Multidentate - donates more than two pairs of electrons - e.g. EDTA^4-

Haem - Fe(||) complex with a multidentate ligand

Shapes of Complex Ions

Octahedral - small ligands -H20, NH3, OH-
Tetrahedral - larger ligands - Cl-
Linear - e.g. Ag+ - [Ag(NH3)2]^+/ [Ag(CN)2]^- / [Ag(S2O3)2]^3-

Formation of Coloured Ions

Metal ions can be identified by there ions
Colour change can be due too - oxidation state - ligand - co-ordination number
Colour arises due too electronic transitions from ground to excited state when the change in energy (DE) is within the visible light spectrum

UV & Visible Light Spectrophotometry

Shine light through a range of different concentrations
Create a graph of absorption against concentration
The greater the concentration the greater the absorption
Add a suitable ligand to an unknown conc too intensify the colour
Measure absorbance of the unknown and compare against the graph to determine the conc

Variable Oxidation States

Vanadium - reduce with zinc and acid
(VO2)^+ à VO^2+ à V^3+ à V^2+
Yellow à Green à Blue à Mauve

Dichromate - reduce with zinc and acid
(Cr2O7)^2- à Cr^3+ à Cr^2+
Orange à Green à Blue

(MnO4)^2- + 8H+ + 5e- à Mn^2+ + 4H2O
5Fe^2+ à 5Fe^3+ + 5e-
Remember in titration’s 5:1 ratio

(Cr2O7)^2- + 14H+ + 6e- à 2Cr^3+ + 7H2O
6Fe^2+ à 6Fe^3+ + 6e-
6:1 ratio

Oxidations of Co(||)

In alkaline solution - oxidising agent: air
[Co(OH)2(H2O)4] à [Co(OH)3(H2O)3]
Colour change - blue/green ppt à pink/beige ppt

In alkaline solution - oxidising agent: H2O2
2[Co(OH)2(H2O)4] + H2O2 à [Co(OH)3(H2O)3] + 2H2O
Colour change - blue/green ppt à pink/beige ppt

In ammonical solution - oxidising agent: air
[Co(H2O)6]^2+ + 6NH3 à [Co(NH3)6]^2+ + 6H2O
Colour change - pink à yellow
[Co(NH3)6]^2+ à [Co(NH3)6]^3+
Colour change - yellow à brown

Oxidations of Cr(|||)

In alkaline solution
[Cr(H2O)6]^3+ + 3OH- à [Cr(OH)3(H2O)3] + 3H2O
Colour change - green à grey/green ppt
Dissolve in excess NaOH
[Cr(OH)3(H2O)3] + 3OH- à [Cr(OH)6]^3- + 3H2O
Colour change - grey/green ppt à dark green
Oxidise with hydrogen peroxixde
[Cr(OH)6]^3- + 3H2O2 à 2(CrO4)^2- + 2OH- + H2o
Colour change - dark green à yellow
Add acid
2(CrO4)^2- + 2H+ à(Cr2O7)^2- +H2O
Colour change - Yellow à Orange

Catalysis

Heterogeneous

Heterogeneous - different phase
Reactions occur on the surface of a catalyst
They provide an alternative reaction route, with a lower activation energy
Increase the conc of reactants on the surface of the catalyst therefore increasing the rate
They adsorb reactants onto their active sites
Active site - where reactants bind to/ where reactions occur

Adsorbtion

Strength of adsorbtion determines catalysts activity
W - too strong
Ag - too weak
Pt/ Ni just right
The story of Goldie Locks and the three catalytic transition metals a personal fave
We use a ceramic support and coat it with catalyst, this reduces costs as less catalyst is used
Also we can shape the catalyst into a honeycomb grip to maximise surface are e.g. using Rh in a catalytic converter

Examples

Catalytic Converter - Rh/ Pd - 2NO + 2CO à N2 + 2CO2 -Lead Poisons catalytic converter
Contact Process - V2O5 - 2SO2 + O2 à 2SO3
Haber Process - Finely divided Fe - N2 +3H2 à NH3 - sulphur poisons Fe
Catalysts can become poisoned by impurities
If the active sites become blocked it reduces efficiency
This can have cost implications especially in industry

Homogeneous

Homogeneous - same phase
Catalysts act as intermediates
(S2O8)^2- + 2I- à 2(SO4)^2- + I2
Reaction is slow without catalyst as both particles negative so repel each other
(S2O8)^2 + 2Fe^2+ à 2(SO4)^2- + 2Fe^3+
2Fe^3+ + 2I- à 2Fe^2+ + I2

Auto catalysis - product of a reaction acts as a catalyst
Again initial rate of reaction slow due too negative particles repelling
Mn^2+ acts as an auto catalyst
2(MnO4)- + 8H+ +5(C2O4)^2- à 2Mn^2+ + 10CO2 + 4H2O
4Mn^2+ + (MnO4)- + 8H+ à 5Mn^3+ + 4H2O
2Mn^3+ + (C2O4)^2- à 2Mn^2+ + 2CO2
Mn^3+ can make fruitful collisions with (C2O4)^2-

Other Applications

Catalysts have variable oxidation states this can be shown by the oxidations and reductions of V2O5 in the contact process
V2O5 + SO2 à V2O4 + SO3
+5 à +4
2V2O4 + O2 à 2V2O5
+4 à +5

Fe(||) allows haemoglobin too carry O2 in the blood as a ligand
CO forms a stronger bond with Fe(||) than O2
This decreases blood O2 levels “so don’t smoke”

Pt(||) - used in cisplatin - anticancer drug

[Ag(NH3)2]^+ - tollens reagent - distinguishes between aldehydes and ketones
RCH2CHO + 2Ag+ + 2H2O à RCH2COOH + 2H+ + 2Ag

[Ag(S2O3)2]^3- - used in photography
AgBr + 2(S2O3)^2- à [Ag(S2O3)2]^3- + Br-

[Ag(CN)2]^- - used in electroplating

Inorganic Compounds In Aqueous Solution

Lewis Acid - electron pair acceptor
Lewis Base - electron pair donator
Aqua ions can be in a solid state e.g. FeSO4.7H2O/ Co(NO3)2.6H2O

Acidity & Hydrolysis

M^2+ - larger than Mn^3+ - less charge than Mn^3+ - therefore less charge density - less polarising - less hydrolysis - weakly acidic

OH-

M(||)
[M(H2O)6]2+ + H2O à [M(H2O)5(OH)]^+ + H3O^+
Equilibrium lies to the left
If a base is added to remove H3O^+
[M(H2O)5(OH)]^+ + H2O à [M(H2O)4(OH)2] + H3O^+

M(|||)
[M(H2O)6]3+ + H2O à [M(H2O)5(OH)]^2+ + H3O^+
Equilibrium lies to the right pH = 3
If base added to remove H3O^+
[M(H2O)5(OH)]^2+ + H2O à [M(H2O)4(OH)2]^+ + H3O^+
[M(H2O)4(OH)2]^+ + H2O à [M(H2O)3(OH)3] + H3O^+

Strong acids reverse reactions

NH3

Hydroxide ppt forms first as NH3 acts as a base
NH3 + H+ à (NH4)^+
Excess ammonia ligand substitution can occur

Na2CO3

M(||)
Carbonate ppt
[M(H2O)6]^2+ + (CO3)^2- à MCO3 + 6H2O

M(|||)
Hydroxide ppt
2[M(H2O)6]^3+ + 3(CO3)^2- à 2[M(H2O)3(OH)3] +3CO2 + 3H2O

Amphoteric Nature of Hydroxides

Dissolve in acids or bases
Al3+ or Cr3+
[Al(OH)4(H2O)2]^- à [Al(H2O)3(OH)3] à[Al(H2O)6]^3+
Colourless à white ppt à colourless
[Cr(OH)6]^3- à [Cr(H2O)3(OH)3] à[Cr(H2O)6]^3+
Dark green à grey/green ppt à green

2(CrO4)^2- + 2H+ à(Cr2O7)^2- +H2O
Yellow à orange

Ligand Substituiton

H2O & NH3 are similar sized and uncharged, ligand substitution occurs without any change in size
[Co(H2O)6]^2+ + 6NH3 à [Co(NH3)6]^2+ + 6H2O
Pink à yellow

[Cr(H2O)6]^3+ + 6NH3 à [C3(NH3)6]^3+ + 6H2O
Green à purple

Ligand substitution can be incomplete
[Cu(H2O)6]^2+ + 4NH3 à [C3(NH3)4(H2O)2]^3+ + 4H2O
Blue à deep blue

Cl- is a larger, charged ligand
As a result co-ordination number decreases as less Cl- can fit around the central metal ion
[Co(H2O)6]^2+ +4 Cl- à [Co(Cl)4]^2- + 6H2O
Pink à blue
[Cu(H2O)6]^2+ +4 Cl- à [Cu(Cl)4]^2- + 6H2O
Blue à yellow/green

Substitution of a unidentate ligand with a bidentate/ multidentate ligand results in a more stable complex
This is due too an increase in entropy
This is known as the chelate effect

good luck for monday

if anyone needs help just ask

**pathologicalliar** · 21-06-2006, 06:49 PM

wow those notes are great *runs to the printer* thanks

...i hate to ask after youve given so much but what exactly is section B all about? Is it just synoptic questions? Thanks

**GJS** · 21-06-2006, 07:05 PM

section B can be anything, however it usually has some relation to the content in module 5, i think you have to be ready for everything, there is only a very little from AS that you need to know, most you have already redone at A2, or can figure out with A2 knowledge

no probs with the notes, was quite helpful for me writing then