Chemistry

chem :)

Structural Isomers

Molecules with the same molecular formula but different arrangement of atoms.

Stereoisomerism

Same structural formula but different spacial arrangements.

Optical Isomerism

Non-superimposable mirror images about a chiral centre. It is a form of geometric isomerism

Testing for optical isomers

Each enantiomer rotates plane polarised light in opposite directions. Shine plane polarised light through the solution containing the isomers and rotate a polarised eye piece to test which direction light has been rotated.

Racemic Mixture

Solution containing a 50/50 mixture of two enantiomers

Standard Conditions

298K, 101kPa (1 atm), 1M

Standard Enthalpy Change of Reaction

The enthalpy change of a reaction when equation quantities of materials react under standard conditions in their standard states.

Standard Enthalpy Change of Formation

Enthalpy change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions.

Standard Enthalpy Change of Combustion

Enthalpy change when 1 mole of a compound undergoes complete combustion with oxygen in standard conditions, with everything in their standard states.

Exothermic Reactions

More energy is released when bonds are formed than energy required for bond breaking.

Endothermic Reactions

More energy is required for bond breaking than energy released from bonds forming.

Bond Enthalpy

Energy required to break 1 mole of GASEOUS bonds into GASEOUS atoms.

Lattice Dissociation Enthalpy

Enthalpy change when 1 mole of an ionic lattice is broken to form gaseous ions.
Always +ve.

Lattice Formation Enthalpy

Enthalpy change when 1 mole of an ionic lattice is formed from its gaseous ions.
Always -ve.

Effect of Ionic Charge on Lattice Enthalpy

Ions with greater charge have stronger attraction and therefore have larger lattice enthalpies.

Effect of Ionic Radius on Lattice Enthalpy

Ions with smaller atomic radii have greater attraction as the distance between centres is reduced and therefore have larger lattice enthalpies.

Standard Atomisation Enthalpy

Enthalpy change when 1 mole of gaseous atoms are formed from an element in its standard state.
Always +ve.

Ionisation enthalpy

Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions.

Electron affinity

Energy required to add 1 mole of electrons to 1 mole of gaseous atoms to form 1 mole of gaseous ions.

Enthalpy Change of Solution

Enthalpy change when 1 mole of an ionic compound is dissolved in water to give an infinite dilution.

Infinite dilution

A solution with enough solvent that adding any more will have no change in concentration.

Hydration Enthalpy

Enthalpy change when 1 mole of gaseous ions dissolve in water to give an infinite dilution. E.g. Na+(g) -\> Na+(aq)

Standard Enthalpy Change of Neutralisation

Enthalpy change when solutions of acid and alkali react together under standard conditions to produce 1 mole of water.

Electrophile

An electron pair acceptor

Nucleophile

An electron pair donor

Transition Metal

Elements which form ions with an incomplete d subshell.

Give an example of an octahedral complex ion

\[Cu(H2O)6]2+

Give an example of a tetrahedral complex ion

\[CuCl4]2-

Give an example of a square planar complex ion

Cis-platin

Give an example of a linear complex ion

\[Ag(NH3)2]+

What colour is [Cu(H2O)6]2+(aq)

Pale blue solution

What colour is [Cu(NH3)4(H2O)2]2+(aq)

Dark blue solution

What colour is [CuCl4]2-(aq)

Yellow solution

What colour is [Cr(H2O)6]3+(aq)

Violet

What colour is [Cr(NH3)6]3+(aq)

Purple

What colour is Cu(OH)2(s)

Blue precipitate

What colour is Fe(OH)2(s)

Green precipitate

What colour is Fe(OH)3(s)

Brown/orange precipitate

What colour is Mn(OH)2(s)

Brown precipitate

What colour is Cr(OH)3(s)

Green precipitate

What colour is [Cr(OH)6]3-(aq)

Dark green solution

Aqueous transition metal + NH3 (not in excess)

Forms transition metal-hydroxide precipitate

\[Cu(H2O)6]2+(aq) + NH3 (in excess)

\[Cu(NH3)4(H2O)2]2+(aq) + 4H2O(l)

\[Cu(H2O)6]2+(aq) + 4Cl-

\[CuCl4]2-(aq) + 6H2O(l)

\[Cr(H2O)6]2+(aq) + NH3 (in excess)

\[Cr(NH3)6]3+(aq) + 6H2O

Cr(OH)3(s) + OH-(aq) (in excess)

\[Cr(OH)6]3-(aq)

Which transition metal ions form hydroxide ligands but not NH3 ligands when reacted with excess ammonia

Fe2+, Fe3+ and Mn2+

Testing for cabonate ions

Add an acid, if carbonate ions are present a gas will be produced. Gas can be bubbled through limewater, CO2 will form a cloudy white precipiate (CaCO3).

Testing for sufate ions

Add barium ions, Ba will react with SO4^2- to form the insolube precipitate BaSO4 which is white and heavy (will sink).

Testing for halides

Add AgNO3(aq), a precipitate will be produced if halide ions are present.

What colour is AgCl(s)?

White, soluble in dilute NH3.

What colour is AgBr(s)?

Cream, soluble in conc. NH3.

What colour is AgI(s)?

Yellow, insolube in NH3.

Sequence for testing for anions

1. Carbonate - halides and sulfate don't produce gas in reaction with dilute acid.

2. Sulfate - barium also forms a white precipitate with carbonate ions.

3. Halides - silver forms precipitates with carbonate and sulfate ions.

Testing for ammonium ions

Add hydroxide ions, NH4+ will react to form NH3(g) though it is highly soluble. Heat the mixture and hold damp litmus (pH) paper above it, if ammonia was produced the litmus paper will turn blue.

Testing for carbonyls

Add 2,4-DNP (2,4-dinitrophenylhydrazine), will form a bright orange precipitate with aldehydes and ketones but not carboxylic acids. This also changes the melting point making it easier to distinguish between carbonyls.

Differentiating ketones and aldehydes

Tollens - in the presence of an aldehyde a silver mirror is produced.
Silver ions are reduced to form Ag(s) and the aldehyde is oxidised into a carboxylic acid; ketones cant be oxidised.

Differentiating phenol and carboxylic acids

Carbonate ions will react with carboxylic acids but not phenol.

How nitrogen oxide radicals break down ozone

Radicals produced by lightning.
Propagation 1: NO¬ + O3 -\> NO2¬ + O2
Propagation 2: NO2¬ + O -\> NO¬ + O2
Overall: O3 + O -\> 2O2
(this is the same as CFCs/Cl radicals)

Examples of disposal of polymers

Landfill - :(
PVC recycling - hazardous when burnt or put in landfill. Modern techniques dissolve the polymer and reprecipitate it. Requires polymers to be sorted.
Waste polymers as fuel - burnt to generate steam and turn turbines.
Feedstock recycling - processes which turn waste polymers into monomers, gases and oil, similar to crude oil. Polymers can be unsorted!

Biodegradable polymers

Broken down by microorganisms. Esters are typically biodegradable.

Bioplastics

Plastics produced from cellulose, starch, plant oils and proteins.

Photodegradable polymers

Contain bonds that are weakened by absorbing light to start the degradation.

Main factor affecting ionisation energy

Atomic radius - greater atomic radius means less nuclear attraction and therefore less energy is required to remove an electron.

Fuel cells

Uses energy from the reaction of a fuel cell with oxygen to produce a voltage.

Advantages of fuel cells

Can operate continously, provided oxygen and fuel are supplied.

Fuel cells do not need to be recharged.

Dynamic Equilibrium

The rate of the forward and reverse reactions are equal.

Concentrations of reactants and products do not change.