Chemistry, Module Five
xX…TheDitzyBlonde…Xx
What are electrochemical cells?
• Made from two different metals dipped in salt solutions of their own ions and
connected by a wire.
• The wire is the external circuit.
• Half-cells involving solutions of two aqueous ions of the same element use a
platinum electrode.
– Conversion of ions happens on the surface of the electrode.
• Half-cells are connected by a salt bridge made of filter paper soaked in
KNO3(aq).
– K+ and NO3- ions flow through the salt bridge and balance out the charges in the
two half-cells.
• Redox reaction occurs in an electrochemical cell.
– Electrons flow through the wire from the most reactive metal to the least reactive
metal.
• Voltmeter in the external circuit shows the voltage between the two half-cells.
– This is called the cell potential or emf, Ecell.
– Ecell = EᶱRHS - EᶱLHS
– More negative standard electrode potential put on the left, so the cell potential is
always positive.
– Swapping the sides the half-cells are on changes the sign of the cell potential.
What are electrode potentials?
• Standard electrode potential, Eᶱ, of a half-cell is the voltage measured under
standard conditions when the half-cell is connected to a standard hydrogen
electrode.
– Standard conditions are solutions at 1M concentration, temperature of 298K and
pressure of 101kPa.
– Ecell = EᶱRHS - EᶱLHS
– Electrode potential can be positive or negative, depending on which way the
electrons flow.
• Standard hydrogen electrode is hydrogen gas bubbled through a 1M solution
of H+ ions with a platinized platinum electrode submerged in, at 101kPa
pressure.
– Platinized platinum electrode means the surface of foil is finely powdered with
platinum.
– The platinized surface absorbs hydrogen gas, causing an equilibrium to be set up.
– H2(g) ↔ H+(aq) + 2e– Electrode potential of a standard hydrogen electrode is 0.
– Reference cell because it is used to measure the electrode potentials of other
half-cells.
– Standard hydrogen electrode always put on the left.
What is the electrochemical series?
• The more reactive the metal, the more likely it is to lose
electrons to form a positive ion.
– More reactive metals have more negative standard electrode
potentials.
• The more reactive a non-metal, the more likely it is to gain
electrons to form a negative ion.
– More reactive non-metals have more positive standard electrode
potentials.
• Electrochemical series shows standard electrode potentials.
– More positive electrode potentials mean the left-hand substance
is more easily reduced and the right-hand substance is more
stable.
– More negative electrode potentials mean the right-hand
substance is more easily oxidised and the left-hand substance is
more stable.
What is the anticlockwise rule?
• Predicts whether a reaction is likely to happen.
1. Write half-equations out with the most negative standard
electrode potential on top.
2. Draw on anticlockwise arrows to show the direction of the
half-reaction.
3. Swap the top half-equation, and change the sign of the
electrode potential.
4. Combine the half-equations and find the cell potential.
• Electrode potential chart is another way of showing the
anticlockwise rule; upside down y-axis showing electrode
potentials with the more negative electrode potential on
top.
What problems are there with
predicting electrode potential?
• If the electrode potential is positive then the reaction is feasible under
standard conditions.
– Above 0.4V, the reaction goes to completion.
– Between 0 and 0.4V then the reaction will be reversible.
• Changing concentration or temperature of the solution can cause the
electrode potential to change.
– If you increase the concentration of the ions in the top half-equation, the
equilibrium shifts left, making electron loss more difficult and the cell potential
lower.
– If you increase the concentration of the ions in the bottom half-equation, the
equilibrium shifts right, making electron gain easier and the cell potential
higher.
• The kinetics may not be favourable:
– Reactions occurs so slowly that it appears that nothing is happening.
– Reaction has a high activation energy, which stops it occurring.
• Bigger cell potential = bigger total entropy change.
• 𝐶𝑒𝑙𝑙 𝑝𝑜𝑡𝑒𝑛𝑡𝑖𝑎𝑙 ∝ ∆𝑆𝑡𝑜𝑡𝑎𝑙 ∝ ln 𝑘
What are redox titrations?
• Titrations using transition element ions.
• Used to determine the amount of oxidising agent needed to exactly
react with a quantity of reducing agent or vice versa.
• Measure out a volume of reducing agent into a conical flask.
• Add 20cm3 of dilute sulfuric acid to the conical flask.
• Add a solution of oxidising agent to the conical flask while swirling
the conical flask.
• Record the volume of oxidation added when the end point is
reached (rough titration).
• Repeat the titration, adding the oxidising agent dropwise when you
are within 2cm3 of the end point, until you get two or more
readings within 0.1cm3 of each other.
• Find average titre.
What can you calculate from redox
titration?
• Concentration of reducing agent:
1.
2.
3.
4.
5.
Write the balanced equation for the reaction.
Work out the number of mols of oxidising agent added to the conical flask.
Use the stoichiometry of the equation to find the number of mols of the
reducing agent that has reacted.
Scale up if needed.
Work out concentration of the reducing agent.
• Percentage purity:
1.
2.
3.
4.
5.
6.
Write the balanced equation for the reaction.
Work out the number of mols of the oxidising agent added to the conical
flask.
Use the stoichiometry of the equation to find the number of mols of
reducing agent that has reacted.
Scale up if needed.
Work out the mass of reducing agent.
Work out the percentage purity.
What are iodine-sodium thiosulfate
titrations?
•
Used to find the concentration of oxidising agents.
– The more concentrated the oxidising agent is, the more ions will be oxidised.
1.
Potassium iodate(V), KIO3, (oxidising agent) is added to excess acidic potassium
iodide solution, oxidising some of the iodide ions to iodine.
– IO3-(aq) + 5I-(aq) + 6H+(aq)  3I2(aq) + 3H2O(l)
2.
Solution produced is titrated with sodium thiosulfate, Na2S2O3.
– I2(aq) + 2S2O32-(aq)  2I-(aq) + S4O62-(aq)
– Colour change from blue/black to colourless.
3.
4.
5.
6.
7.
8.
Find average titre.
Work out the number of mols of thiosulfate that reacted.
Use the stoichiometry of the second equation to work out the number of mols of
iodine that reacted.
Use the stoichiometry of the first equation to work out the number of mols of
iodate(V) ions that reacted.
Scale up if needed.
Calculate the concentration of potassium iodate(V).
How can iodine-sodium thiosulfate titrations be used to
calculate percentage copper in a copper alloy?
1.
Weighed amount of copper alloy dissolved in conc nitric acid and made up to
250cm3 with deionised water in a volumetric flask.
25cm3 portion added to conical flask
Excess nitric acid is neutralised with sodium carbonate solution, until white
ppate appears.
2.
3.
–
4.
Excess potassium iodide added, which reacts with copper(II) ions.
–
–
5.
2Cu2+(aq) + 4I-(aq)  2CuI(s) + I2(aq)
CuI is seen as a white ppate.
Product mixture is titrated against sodium thiosulfate solution.
–
6.
7.
8.
9.
10.
11.
White ppate removed by adding ethanoic acid.
I2(aq) + 2S2O32-(aq)  2I-(aq) + S4O62-(aq)
Find average titre.
Work out the number of mols of sodium thiosulfate that reacted.
Use the stoichiometry of the equations to find the number of mols of copper.
Scale up.
Work out the mass of copper.
Work out percentage of copper in the copper alloy.
What errors can occur in iodinesodium thiosulfate titrations?
• Starch indicator must be added when most of the
iodine has reacted, otherwise the blue colour will be
very slow to disappear.
• Starch indicator must be freshly made, otherwise it
won’t behave as expected.
• Ppate of CuI makes seeing the colour of the solution
quite difficult.
• Iodine produced in the reaction can evaporate from
the solution, giving a false titration reading, leading to
a percentage of copper lower than the actual value.
– Avoid this by keeping the solution cool.
What are fuel cells?
• Produces electricity by reacting a fuel such as hydrogen with an oxidant such
as oxygen.
1. At the anode (fuel), the platinum catalyst splits H2 into protons and
electrons.
2. Polymer electrolyte membrane (PEM) only allows H+ ions across, and
forces electrons to travel around the circuit to get to the cathode.
3. This created an electric current in the circuit, with voltage about 0.6V,
which is used to power things.
4. At the cathode, O2 combines with H+ from the anode, and electrons from
the circuit to make water, which is the only waste product.
• Cell continues to produce a current as long as there is hydrogen and oxygen
available.
• Hydrogen has to be made to use the fuel cell:
– Natural gas reacted with steam to produce hydrogen gas gives carbon dioxide as
a waste product and needs fossil fuels to heat the process.
– Electrolysis of water to produce hydrogen requires electricity, which is generated
using fossil fuels, but could be generated sustainably using renewable resources.
What fuels can be used in fuel cells
other than hydrogen?
• Hydrogen-rich fuels are fuels with a high percentage of hydrogen in their
molecules, e.g. ethanol or methanol.
– These are converted into hydrogen by a reformer.
• Some fuel cells can use alcohols without have to reform them to produce
hydrogen first.
– Alcohol oxidised at the anode in the presence of water.
• CH3OH + H2O  CO2 + 6e- + 6H+
– H+ ions pass through the electrolyte and are oxidised to water.
• 6H+ + 6e- + (3/2)O2  3H2O
• Advantages of using alcohols:
– Higher hydrogen density than liquefied hydrogen.
– Methanol and ethanol can be made on a large scale using renewable
resources.
– Alcohols are liquid at room temperature so don’t need refrigerated storage.
– Methanol can be made from carbon dioxide, so can be used to reduce carbon
dioxide levels in the atmosphere.
What are the advantages and disadvantages of
fuel cell cars compared to normal cars?
• Advantages:
– Produce less pollution when they are used.
• Only waste product of hydrogen fuel cell is water.
• Engines running on hydrogen-rich fuels produce less carbon
dioxide than petrol or diesel engines.
– More efficient at converting fuel to power.
• Disadvantages:
– Expensive.
– Made using toxic chemicals, which are difficult to
dispose of.
– Limited life-span.
What are fuel cells used for?
• Powering cars.
• Space Shuttle uses 3 fuel cell power plants, each made up of many fuel
cells, which produce up to 12kW of power.
– Waste product of water can be used as drinking water for astronauts.
• In breathalysers, as the amount of ethanol in someone’s breath is
directly related to the amount of ethanol in the blood.
– Used to be measured using the reaction of ethanol with potassium
dichromate(VII), and measuring the colour change using a photocell
system.
– Breathalysers in police stations use IR spectrometry, which is accurate but
not easily portable.
– Ethanol fuel cells can be used, as the amount of alcohol present is
proportional to the current produced when the person’s breath is fed to
the anode of the cell.
• Less susceptible to producing false readings to other substances in the breath.
• Easily portable.
• Accurate.
What are the transition metals?
• D-block is in the middle of the periodic table.
– D-block elements are those that are found in the d-block of the periodic table.
• Transition metals are elements that can form one or more stable ions with a
partially filled d-subshell.
– Scandium is not a transition metal because it only forms one ions, Sc3+, which
has an empty d-subshell.
– Zinc is not a transition metal because it only forms one ions, Zn2+, which has a
full d-subshell.
• Electron configuration is given as [Ar]3dx4sy.
– All orbitals are occupied singly before any orbitals have two electrons.
– 4s subshell fills before the 3d subshell except chromium and copper, which have
only one electron in the 4s subshell as this gives them more stability.
• Special chemical properties:
–
–
–
–
Form complex ions.
Form coloured ions in solution.
Are good catalysts.
Can exist in variable oxidation states.
How are ions of transition metals
formed?
• Electrons are removed, first from the 4s subshell, then from the 3d
subshell.
• Ionisation energy is the energy required to remove an electron from
an element.
• First ionisation energies are similar for d-block elements, showing
they have similar structures, and lose the first electron from the
same shell.
• Second ionisation energies increase steadily along the period, with
exceptions of Cr and Cu, which have a slightly higher IE than
expected because the 2nd electron is removed from the 3d-subshell,
which is much closer to the nucleus than the 4s-subshell.
• Third ionisation energies increase steadily, with a step down in
ionisation energy at iron because form that point onwards the third
electron is being removed from a paired 3d orbital, which is easier
to remove due to repulsion between the electrons.
What are complex ions?
• Complex ions are metal ions surrounded by dative covalently bonded ligands.
• Dative covalent bonds are covalent bonds where both electrons in the shared
pair come from the same atom.
• Ligands are atoms, ions or molecules that donate a pair of electrons to a
central metal ion.
• Coordination number is the number of dative covalent bonds that are formed
with the central metal ion.
– Small ligands such as H2O, NH3 and CN- result in a coordination number of 6.
• Octahedral shape.
– Larger ligands such as Cl- result in a coordination number of 4.
• Tetrahedral shape if all Cl- ligands.
• Square planar shape if a mix of small and larger ligands.
– Some complexes only have 2 dative covalent bonds so are linear.
• (CuCl2)- and [Ag(NH3)2]+
• Ligand substitution is when one ligand is replaced with a different ligand.
– Causes a colour change.
• Oxidation state of metal ion = total oxidation state – oxidation states of
ligands.
How do ligands form bonds?
•
Using lone pairs of electrons:
– Monodentate ligands have one lone pair of electrons.
•
E.g. H2O, NH3, Cl-, CN-
– Bidentate ligands have two lone pairs of electrons.
•
E.g. ethane-1,2-diamine and ethanedioate.
– Polydentate ligands have more than two lone pairs of electrons.
•
•
E.g. EDTA+ has six lone pairs of electrons and haemoglobin has four lone pairs of electrons.
Ligands split the 3d subshell into two energy levels, creating a visible colour.
– Usually 3d orbitals of a transition metal ion all have the same energy.
– Repulsion between electrons in the ligand and the 3d electrons of the metal increases the
energy of the 3d orbitals of the transition metal ion in the complex.
– Some orbitals are increased in energy level more than others, splitting the 3d orbital into two
different energy levels.
– Electrons occupy lower orbitals but can jump to higher energy level orbitals by absorbing energy
from visible light equal to the energy gap.
– The remaining frequencies of light are transmitted, causing the complement colour to be seen.
•
•
•
The central metal ion, the ligands and the coordination number affect the size of the energy gap, and
thus the colour seen.
If there are no 3d electrons, there are no electrons to jump energy levels so no energy is absorbed and
no colour is seen.
If the 3d subshell is full then there is no room in the upper orbitals for electrons to jump to, so no energy
is absorbed and no colour is seen.
What is ligand exchange?
• Swapping of a ligand in a complex ion with a different ligand.
– Causes a colour change.
• If the ligands are similar size then the coordination number doesn’t change.
• If the ligands are different sizes then the coordination number changes and
so does the shape of the complex.
• Ligand exchange can occur partially.
– E.g. [Cu(H2O)6]2+(aq) + 4NH3(aq)  [Cu(NH3)4(H2O)2]2+(aq) + 4H2O(l)
• Shape changes from octahedral to elongated octahedral.
– E.g. [Cr(H2O)6]3+(aq) + SO42-(aq)  [Cr(H2O)5(SO4)]+(aq) + H2O(l)
• Shape changes from octahedral to distorted octahedral.
• Can be reversed, unless the new complex is more stable than the original
complex.
– If the new bonds between the ligand and the transition metal ion are stronger
than the original bonds.
– Bidentate ligands for more stable complexes than monodentate ligands.
– Polydentate ligands form more stable complexes than bidentate and
monodentate ligands.
• This is because the number of molecules increases so the entropy is likely to increase.
What reactions do ligands undergo?
• Addition of sodium hydroxide produces insoluble metal hydroxides by
disproportionation.
– Metal aqua 3+ ions:
• [M(H2O)6]3+(aq) + H2O(l) (equil)[M(H2O)5(OH)]2+(aq) + H3O+(aq)
– Adding OH- ions shifts the equilibrium to the right because H3O+ ions are removed.
• [M(H2O)5(OH)]2+(aq) + H2O(l) (equil) [M(H2O)4(OH)2]+(aq) + H3O+(aq)
– Adding OH- ions shifts the equilibrium right to replace the H3O+ ions that are removed.
• [M(H2O)4(OH)2]+(aq) + H2O(l) (equil) M(H2O)3(OH)3(s) + H3O+(aq)
– M(H2O)3(OH)3(s) is insoluble in water because it is uncharged, therefore is seen as a white ppate.
– Metal aqua 2+ ions:
• Same process occurs as with 3+ ions.
– [M(H2O)6]2+(aq) + H2O(l) (equil) [M(H2O)5(OH)]+(aq) + H3O+(aq)
– [M(H2O)5(OH)]+(aq) + H2O(l) (equil) M(H2O)4(OH)2(s) + H3O+(aq)
• Ammonia dissolves in water:
– NH3 + H2O (equil) NH4+ + OH-
• Adding a small amount of ammonia solution gives the same result as adding sodium
hydroxide because OH- ions are formed in the equilibrium.
• Adding excess ammonia solution can cause ligand exchange to occur, producing a soluble
metal ion complex.
What colour changes occur when complexes react?
Metal aqua ion
With OH-(aq) or
NH3(aq)
With excess OH-(aq)
With excess NH3(aq)
[Cu(H2O)6]2+
Blue solution
Cu(H2O)4(OH)2
Blue ppate
No change
[Cu(NH3)4(H2O)2]2+
Deep blue solution
[Fe(H2O)6]2+
Green solution
Fe(H2O)4(OH)2
Green ppate
No change
No change
[Fe(H2O)6]3+
Yellow solution
Fe(H2O)3(OH)3
Brown ppate
No change
No change
[Cr(H2O)6]3+
Violet solution
Cr(H2O)3(OH)3
Green ppate
[Cr(OH)6]3Green solution
[Cr(NH3)6]3+
Purple solution
[Mn(H2O)6]2+
Pale pink solution
Mn(H2O)4(OH)2
Brown ppate
No change
No change
[Ni(H2O)6]2+
Green solution
Ni(H2O)4(OH)2
Green ppate
No change
[Ni(NH3)6]2+
Blue solution
[Zn(H2O)6]2+
Colourless solution
Zn(H2O)3(OH)3
White ppate
[Zn(OH)4]2Colourless solution
[Zn(NH3)4]2+
Colourless solution
What oxidation states can copper exist as?
• +1 (Cu+):
– Stable in solid copper(I) compounds.
• Can be coloured, such as Cu2O, which is red.
– Unstable in aqueous solution, disproportionating to Cu2+(aq) and
Cu(s)
• 2Cu+(aq)  Cu2+(aq) + Cu(s)
– Can form some stable ligands such as [CuCl2]- forms if there are
excess Cl- ions to stabilise the copper(I) and prevent it
disproportionating.
• Complexes cannot be coloured because the 3d subshell of copper(I) ions
is full.
• +2 (Cu2+):
– Stable in aqueous solution.
• Reduced to Cu metal by more electrophilic metals via a displacement
reaction.
– E.g. Cu2+ + Zn  Cu + Zn2+
What oxidation states can chromium
exist as?
• Chromium is used to make stainless steel and added to steel to make it
harder.
• +2 (Cr2+):
– Least stable oxidation state.
• +3 (Cr3+):
– Most stable oxidation state.
– Violet colour solution when surrounded by 6 water ligands.
– Green colour solution when some of the water ligands are substituted for
impurities in water such as Cl- ligands.
• +6 (CrO42-):
– Chromate(VI) ions.
– Good oxidising agent as it is easily reduced to Cr3+.
• +6 (Cr2O72-):
– Dichromate(VI) ions.
– Good oxidising agent as it is easily reduced to Cr3+ using a reducing agent such
as zinc and dilute acid.
What reactions do different oxidation
states of chromium do?
• Dichromate(VI) ions can be reduced to Cr3+ ions using a reducing agent such
as zinc and dilute acid.
– Cr2O72-(aq) + 14H+(aq) + 3Zn(s)  3Zn2+(aq) + 2Cr3+(aq) + 7H2O(l)
– Colour change from orange to green.
• Cr3+ ions can be reduced to Cr2+ ions in an inert atmosphere using a zinc.
– 2Cr3+(aq) + Zn(s)  Zn2+(aq) + 2Cr2+(aq)
– Colour change from green to blue.
– Cr2+ is so unstable that it oxidises straight back to Cr3+ ions in air.
• Cr3+ can be oxidised to chromate(VI) ions by hydrogen peroxide in alkaline
solution.
– 2Cr3+(aq) + 10OH-(aq) + 3H2O2(aq)  2CrO42-(aq) + 8H2O(l)
– Colour change from green to yellow.
• Dichromate(VI) ions can be converted to chromate(VI) ions using a strong
alkali.
– Cr2O72- + 2OH-  2CrO42- + H2O
• Chromate(VI) ions can be converted to dichromate(VI) ions using an acid.
– 2CrO42- + 2H+  Cr2O72- + H2O
What is chromium(II) ethanoate?
•
•
Cr2(CH3COO)4(H2O)2
Made using chromium(II) chloride solution.
– Chromium(II) chloride is reduced with zinc in acid solution to produce Cr2+ ions.
•
•
2Cr3+ + Zn  2Cr2+ + Zn2+
Colour change from green to blue.
– Sodium ethanoate is mixed with the solution to produce a red ppate of chromium(II)
ethanoate.
•
•
1.
2.
3.
4.
5.
6.
2Cr2+ + 4CH3COO- + 2H2O  [Cr2(CH3COO)4(H2O)2]
The complex is easily oxidised so it has to be made in an inert atmosphere such as nitrogen.
HCl added to a flask containing chromium(II) chloride solution and zinc mesh to
reduce Cr3+ ions to Cr2+ ions.
Hydrogen produced can escape through a rubber tube into a beaker of water.
When the solution turns clear blue, pinch the rubber tube shut so hydrogen can
no longer escape.
Pressure builds up in the flask, forcing Cr2+ ions through the open glass tube into
a flask of sodium ethanoate.
Red ppate of chromium(II) ethanoate forms.
Ppate filtered off and washed using water, then ethanol, then ether.
Why is chromium hydroxide
amphoteric?
• Can react with both acids and alkalis.
– [Cr(H2O)3(OH)3](aq) + 3H+(aq)  [Cr(H2O)6]3+(aq)
– [Cr(H2O)3(OH)3](aq) + 3OH-(aq)  [Cr(OH)6]3-(aq) + 3H2O(l)
• Zn is the only other 3d element that can do
this.
• Not ligand exchange, but ligand is chemically
modified.
How do transition metals catalyse
reactions?
• Good catalysts because they can use their s and d orbitals to form bonds with
reactants and can change oxidation state easily.
• Can act as homogeneous or heterogeneous catalysts.
• Platinum and rhodium are used in catalytic converters in cars to convert nitric
oxide and carbon monoxide to nitrogen and carbon dioxide.
•
– Catalyst used is a mesh or fine powder to increase surface area.
– 2NO(g) + 2CO(g)  N2(g) + 2CO2(g)
1. Reactant molecules bond to the surface of the solid catalyst (adsorption).
2. Bonds between the reactant atoms are weakened and break, forming radicals.
3. Radicals join to make the new molecules.
The Contact process produces large quantities of sulfuric acid using a vanadium
catalyst.
•
2SO2(g) + O2(g)  2SO3(g)
• Vanadium(V) oxide reacts with sulfur dioxide and is reduced to vanadium(IV).
– V2O5(s) + SO2(g)  V2O4(s) + SO3(g)
• Vanadium(IV) oxide is oxidised back to vanadium(V) by oxygen.
– 2V2O4(s) + O2(g)  2V2O5(s)
How can catalyst development be used
to make chemistry greener?
• Allow reaction to occur at a lower temperature or pressure.
– Saves energy and reduces cost.
• Can improve atom economy.
– Reduces amount of raw materials that need to be used.
– Reduces amount of waste products.
• E.g. Making ethanoic acid from methanol and carbon monoxide:
– Originally made by oxidising butane but atom economy was very low.
– Methanol first used to make ethanoic acid in the 1960’s using a catalyst
of cobalt/iodine.
• Methanol is a cheaper, more environmentally friendly raw material than butane.
• Higher atom economy than using butane.
– Rhodium/iodine catalyst developed in 1966, which requires lower
temperature and pressure than cobalt/iodine catalyst.
– Iridium/iodine catalyst developed in 1980’s made reaction faster and
more efficient.
What are transition metals used for?
•
•
•
•
Construction because they are strong and unreactive.
Jewellery because they are malleable and inert, and often look shiny.
Paint pigments/coloured glass.
Chemotherapy:
– Cisplatin is a complex of platinum(II) with two chloride ions and two ammonia
molecules.
– Square planar shape.
– Chloride ions next to each other.
– Prevents cancer cells from dividing by crosslinking the DNA, causing the cells to
die.
– Also prevents normal cells dividing including in the blood, so can supress the
immune system, increasing the risk of infection.
– Can damage the kidneys.
• Polychromic sunglasses:
– Sunglasses embedded with a silver halide which decomposes in UV radiation to
form silver atoms.
– Silver atoms darken the lenses.
– Process reverses when there is no UV, reforming the silver halide, lightening the
lenses.
What are aromatic compounds?
•
•
•
•
•
Derived from benzene.
Also known as arenes.
Named as substituted benzene rings or as compounds with a phenyl group.
Benzene = C6H6
Kekule’s structure and delocalised structure are the two ways of
representing benzene.
– Kekule’s structure consists of carbon atoms in a ring with alternating single and
double bonds.
• X-ray diffraction studies show all carbon-carbon bonds in benzene are the same length,
between the length of C-C and C=C bonds.
• IR studies have shown no carbon-carbon bonds in benzene are C-C or C=C bonds because
they absorb energy at different frequencies.
– Delocalised model is the now accepted model.
• P-orbitals of all 6 carbon atoms overlap to create pi bonds.
• Two ring shaped electron clouds are formed; one above the plane and one below the
plane.
• Electrons in the ring are delocalised, and represented by a circle inside the carbon ring.
• All bonds are the same so are the same length.
How does delocalisation of electrons
give benzene stability?
• Cyclohexane has one double bond.
– Hydrogenation of cyclohexane to hexane gives an enthalpy
change of -120kJmol-1.
– If benzene had Kekule’s structure, you would expect the
enthalpy of hydrogenation to be -360kJmol-1.
• Actual enthalpy of hydrogenation of benzene is 208kJmol-1.
– Less exothermic than the expected enthalpy of
hydrogenation.
– More energy is needed to break the bonds of benzene than
would be required to break the bonds if benzene had
Kekule’s structure.
– This is because the delocalised system gives benzene more
stability.
How does benzene react with oxygen?
• Benzene is a hydrocarbon.
• Burns in oxygen to produce carbon dioxide
and water.
– 2C6H6 + 15O2  12CO2 + 6H2O
• Burns in air to give a smoky flame because
there is too little oxygen so the benzene does
not burn completely.
– Carbon atoms stay as carbon and form soot
particles in hot gas.
How do arenes react with nitronium
ions?
• Electrophilic substitution.
• Benzene warmed with conc nitric acid and sulfuric acids to
produce nitrobenzene.
• Nitrobenzene forms as a layer floating on top of the acid.
• Sulfuric acid acts as a catalyst to make the nitronium ion,
NO2+.
– HNO3 + H2SO4  H2NO3+ + HSO4– H2NO3+  NO2+ + H2O
– H+ + HSO4-  H2SO4
• The nitronium ions acts as an electrophile.
• Temperature must be kept below 55 degrees for
mononitration, as above this, extra substitutions occur.
How do arenes react with sulfur
trioxide molecules?
• Electrophilic substitution.
• Benzene warmed with fuming sulfuric acid at
40 degrees for about 30 mins.
– Fuming sulfuric acid is sulfur trioxide, SO3,
dissolved in sulfuric acid.
• Electrophile is sulfur trioxide, SO3.
How can halogen carriers be used to
make good electrophiles?
• Electrophiles have a strong positive charge.
• Halogen carriers can be used as a catalyst to
make strong electrophiles.
– E.g. aluminium halides, iron halides and iron.
• Halogen carriers accept lone pairs of electrons
from the halogen in the electrophile.
• This causes polarisation of the R-X bond to
increase, forming a carbocation.
– RCl: - - - > AlCl3  R+AlCl4-
What is a Friedel-Crafts alkylation
reaction?
• Halogenoalkane reacts with benzene via
electrophilic substitution to produce
alkylbenzenes.
– Refluxed in dry ether because aluminium chloride is
sensitive to hydrolysis.
• Halogen carrier of aluminium chloride is used as a
catalyst.
– RCl + AlCl3  R+ + AlCl4-
• Alkylbenzenes are more reactive than benzene so
more than one substitution occurs
(polyalkylation).
What is a Friedel-Crafts acylation
reaction?
• Acyl chlorides react with benzene by electrophilic
substitution to produce phenylketones.
– Heated under reflux with dry ether.
• Acyl chlorides are polar so can easily lose chlorine
atoms to form carbocations.
• Halogen carrier such as aluminium chloride is
used as a catalyst.
– RCOCl + AlCl3  RCO+ + AlCl4-
• Phenylketones are less reactive than benzene so
only one substitution occurs.
How do arenes react with halogens?
• Cycloalkene shaken with bromine water.
– Bromine adds to the cycloalkene to give
dibromoalkane.
– Bromine water is decolourised (can be used to test for
alkenes).
• Bromine isn’t a strong enough electrophile to
react with benzene unless a halogen carrier
catalyst is also used.
– Halogen carrier polarises the halogen, creating a
strong electrophile that can attack the benzene ring.
• Br2 + FeBr3  Br+ + FeBr4-
– Bromobenzene is formed.
How does benzene react with
hydrogen?
• Finely divided nickel catalyst used.
• Temperature of 150 degrees and pressure of
10 atm required.
• Cyclohexane is produced.
• C6H6 + 3H2  C6H12
What reactions do methylbenzene and
methoxybenzene undergo?
• CH3 and OCH3 groups donate electrons to the
delocalised ring, increasing the rings electron
density, making it more reactive.
• React faster than benzene.
• Less toxic than benzene so they are safer to
use.
• Products will be slightly different but the
mechanism is the same as when benzene is
used.
What is phenol?
• Benzene ring with an –OH group substituted for an Hatom.
• Phenol reacts by electrophilic substitution when shaken
with bromine water, decolourising bromine water and
forming 2,4,6-tribromophenol, which is an insoluble
white ppate in water that smells of antiseptic.
– Phenol reacts with bromine water because one of the
electron pairs in the p-orbital of oxygen overlaps with the
delocalised ring of benzene, increasing the electron density
of the ring.
• Phenol reacts with dilute nitric acid to produce two
isomers of nitrophenol, and water.
– Easier to nitrate than benzene because it has a higher
electron density in the ring.
What are amines?
• Ammonia with one or more hydrogen replaced with an organic group.
– Aliphatic amines are chained/branched.
– Aromatic amines have a benzene ring.
• Small amines have a slightly fishy smell, larger amines smell very fishy.
• Solubility:
– Small amines are soluble in water because amines can form hydrogen
bonds to water molecules.
– Bigger amines are less soluble in water because there are greater London
forces between the amine molecules.
– Amines form alkaline solutions when dissolved by forming alkyl
ammonium ions and hydroxide ions.
• React with acids to form salts:
– Can act as bases because they have a lone pair on the nitrogen atom so
can form a dative covalent bond with H+ ions (accept protons).
– CH3CH2NH2 + HCl  CH3CH2NH3+Cl-
How do amines form complex ions?
• Copper(II) sulfate solution forms [Cu(H2O)6]2+
complexes with water.
• Adding a small amount of methylamine solution
produces a pale blue ppate of [Cu(H2O)4(OH)2] by
acting as a Bronsted-Lowry base and taking two
H+ ions from the complex.
• Adding excess methylamine solution causes the
ppate to dissolve, producing a deep blue solution
of [Cu(CH3NH2)4(H2O)2]2+ complex ions by ligand
exchange.
How are amines acylated to Nsubstituted amides?
• Butylamine + ethanoyl chloride  N-butylethanamide
– Butylamine + HCl  Butylammonium chloride
– CH2COCl + 2C4H9NH2  CH3CONHC4H9 + [C4H9NH]+Cl– Ethanoyl chloride must be added to a concentrated aqueous solution of
amine.
• Chloroethane + butylamine  ethylbutylamine + HCl
– Lone pair of the nitrogen attacks the slightly positive carbon atom on the
halogenoalkane, displacing the chlorine and making a salt.
– Chlorine removes a hydrogen from the salt to produce the amine.
• Ethanoyl chloride + phenylamine  N-phenylethanamide + HCl
– Phenylamine + HCl  phenylammonium chloride
– Product stained brown with unreacted phenylamine.
• Paracetamol is made by reacting ethanoyl chloride with paminophenol.
– HCl is made as a waste product.
What are amides?
• Derivatives of carboxylic acids.
• Functional group of –CONH2.
• N-substituted amides are amides with one
hydrogen replaced with an alkyl group.
– Carbonyl group (C=O) pulls electrons away from the
rest of the group.
• Made by acylation of ammonia or amines.
– Acyl chloride + ammonia  amide + HCl
– Acyl chloride + amine  N-substituted amide + HCl
– HCl can go on to react with ammonia/amine to form
an ammonium salt.
What are amines?
• Made by reducing a nitro compound such as nitrobenzene.
– Heated with tin and conc HCl under reflux to make an
ammonium salt.
– Sodium hydroxide solution added to remove H+ ions from the
salt, forming an aromatic amine.
– Nitrobenzene + 6[H]  phenylamine + 2H2O.
• Used to make azo dyes:
– Azo dyes contain the group –N=N– Made by reacting aromatic amines with phenols.
– Azo group links two aromatic groups.
• Two aromatic groups give stability to the molecule.
• Azo group becomes part of the delocalized electron system.
• Colours are a result of light absorption by the delocalised electron
system.
How are azo dyes made?
• Diazonium salt made by reacting nitrous acid with phenylamine:
– Nitrous acid made in situ from sodium nitrite and HCl because it is
unstable.
• NaNO2 + HCl  HNO2 + NaCl
– Nitrous acid reacts with phenylamine and HCl to form benzenediazonium
chloride.
– Temperature must be about 5 degrees to prevent phenol being made
instead.
– Diazonium salts are unstable and decompose quickly so the azo dye must
be made straight away.
• Azo dye made from diazonium salts by a coupling reaction with phenol.
– Phenol is dissolved in sodium hydroxide to make sodium phenol.
– Sodium phenol is stood in ice and chilled benzenediazonium chloride is
added.
– Azo dye ppates out of the solution.
– Lone pairs on the oxygen of phenol act as a coupling agent, increasing the
electron density of the benzene ring, giving the diazonium ion electrophile
an electron rich area to attack.
What are amino acids?
• Contains amino group (NH2) and carboxyl group (COOH).
– Amphoteric because the amino group is basic and the carboxyl group is
acidic.
• Alpha amino acids have the amino group and the carboxyl group
attached to the same carbon atom.
– General formula RCH(NH2)COOH.
• Monomers used to make proteins.
• Can exist as zwitterions.
– Dipolar ions with both a positive and a negative charge in different parts
of the molecule.
– Can only form zwitterions near the amino acid’s isoelectric point; the
pH where the average overall charge of the amino acid is zero.
• In conditions more acidic than the isoelectric point, the NH2 group is likely to be
protonated to NH3+.
• At the isoelectric point, the carboxyl group and the amino group are ionised.
• In conditions more basic than the isoelectric point, the COOH group is likely to
lose protons, becoming COO-.
What is paper chromatography?
• Used to identify unknown amino acids.
1.
2.
3.
4.
5.
Pencil line drawn near the bottom of a piece of chromatography paper.
Spot of amino acid mixture put on the line.
Bottom of the paper is dipped into solvent.
As the solvent spreads up the paper, the amino acids move at different rates
so separate.
When the solvent reaches the top of the paper, the paper is removed form
the solvent and a line is made to mark the distance the solvent travelled
(solvent front).
• Amino acids are colourless so they are located using ninhydrin solution.
– Ninhydrin reacts with amino acids to produce ammonia, aldehydes, carbon
dioxide and hydrindantin.
– Hydrindantin reacts with ammonia and ninhydrin to give a purple pigment
called Ruhemann’s purple.
– Amino acid spots turn purple.
• R1 value = distance travelled by spot/distance travelled by solvent.
– Table of R1 values can be used to identify the amino acids in the mixture.
Why are amino acids chiral?
• Alpha amino acids usually have four different groups
attached to a central carbon atom.
• This makes the central carbon a chiral centre, so two
optical isomers exist.
• When plane-polarised monochromatic light is shone
through an aqueous solution of alpha amino acid, the
plane of light is rotated because of the chiral carbon
atom.
• Glycine does not show optical activity because it has
two hydrogen atoms attached to the central carbon, so
it does not rotate the plane of plane polarised light.
What are polymers?
• Long chain molecules made up of small, repeating
units called monomers.
• Alkenes form polymers by addition polymerisation.
– Double bonds open and molecules join to make long
chains.
• Condensation polymers are formed by removing water.
– E.g. polyester, polyamides, polypeptides.
– Each monomer in a condensation polymer must have at
least two functional groups.
– Each functional group reacts with a functional group of
another monomer to form the polymer.
What is polyethanol?
• Polyethanol is a water soluble addition polymer
made from ethanol monomers.
– Water soluble because it can hydrogen bond with
water via the –OH groups.
• Used for water soluble laundry bags in hospitals
to prevent spread of infection, as the polyethanol
bag breaks down at 40 degrees.
• Used to wrap liquid detergents to form liquitabs,
as the polyethanol wrapping dissolves.
What are condensation polymers?
• Proteins are made by reacting amino acids to form peptide links (-CONH-).
– Dipeptide is two amino acids joined by a peptide bond.
– Polypeptides are several amino acids joined by peptide bonds.
• Polyesters are made by reacting dicarboxylic acids diols to form ester links
(-COO-).
– Terylene (PET) is made from benzene-1,4-dicarboxylic acid and ethane-1,2diol.
• Strong, flexible and abrasion resistant.
• Used to make clothing.
• Polyamides formed by reaction of dicarboxylic acids and diamines to form
amide links (-CONH-).
– Nylon 6,6 is made from 1,6-diaminohexane and hexanedioic acid.
• Strong, elastic and abrasion resistant.
• Used to make ropes, clothes, and carpets.
– Kevlar is made from benzene-1,4-diamine and benzene-1,4-dicarboxylic acid.
• Strong and light.
• Used to make bulletproof vests.
What is the difference between
empirical and molecular formulae?
• Empirical formula gives the smallest whole number ratio of atoms
in a compound.
– Calculated using experimental data:
1.
2.
3.
4.
5.
Calculate the total number of mols of the product of the reaction.
Use the Mr to calculate the number of mols of each atom in the compound
that reacted.
Form ratio of atoms in the compound and their mols.
Simplify the ratio by dividing each number by the smallest.
Write empirical formula using the ratio.
• Molecular formula gives the actual numbers of atoms in a molecule.
– Calculated from experimental data:
1.
2.
3.
Follow steps for finding the empirical formula.
When the empirical formula is found, compare the molecular mass of the
empirical formula to the actual molecular mass.
Form the molecular formula by multiply the empirical formula so the
molecular mass of the empirical formula and the actual molecular masses
are equal to each other.
What is refluxing?
• Ensures you don’t lose volatile organic substances.
– Used to improve rate of reaction without losing volatile
substances.
– Volatile substances are substances with low boiling points.
• Mixture heated in a flask fitted with a vertical ‘Liebig’
condenser.
– ‘Liebig’ condenser condenses vapours and recycles them
back to the flask.
• Electrical heating such as hot plates is usually used as
naked flames could ignite flammable substances.
How can a sample be purified?
•
Melting and boiling points indicate purity:
– Impure substances have lower melting points and higher boiling points than expected.
– Melting point measured by adding solid to a capillary tube and placing it in a beaker of oil with a
sensitive thermometer.
•
Slowly heat until the solid melts and then read the temperature off the thermometer.
– Boiling point measured by measuring the temperature the liquid is collected at during
distillation.
•
Washing:
– Removes unwanted side products and unreacted reagents.
•
Solvent extraction:
– Separates product from a mixture by dissolving it in a solvent.
•
Drying:
– Add anhydrous calcium chloride granules to remove water.
– Crystals are formed, which can be filtered off.
•
Recrystallization:
– Removes impurity in a solid.
1.
Add hot solvent until solid just dissolves to produce a saturated solution of impure product.
2.
Solution cooled slowly.
•
3.
Impurities stay in solution because they’re present in smaller amounts so take longer to crystallise.
Remove crystals by filtration, wash crystals, and leave to dry.
What is fractional distillation?
• Separates liquids with different boiling points.
1.
2.
3.
4.
5.
6.
Mixture heated in distillation apparatus.
Liquid in the flask boils.
Vapour goes up the fractionating column.
As the vapours cool they condense.
Vapours that fall below the collecting boiling point condense
and run back down the fractionating column through the glass
beads, while vapours that have a higher boiling temperature
go through the ‘Liebig’ condenser.
Temperature is increased to distil different liquids.
• To stop products decomposing on heating, steam can be
passed over the mixture to lower the products boiling
point, so it can be distilled at a lower temperature.
How is cholesteryl benzoate prepared?
• Used in liquid crystal displays, hair colours and cosmetics.
• Ester of cholesterol and benzoic acid.
• Prepared from benzoyl chloride in a fume cupboard:
1.
Cholesterol dissolved in pyridine.
• Pyridine is toxic.
2.
Benzoyl chloride is added.
• Benzoyl chloride is a lachrymator (makes you cry).
3.
4.
5.
6.
7.
Heat in steam bath for 10 mins.
Cool, then add methanol.
Filter off crystals of the ester that form.
Wash crystals with methanol.
Recrystallize the ester using ethyl ethanoate as the solvent.
How is methyl 3-nitrobenzoate
prepared?
• Prepared by nitrating methyl benzoate.
1. Dissolve methyl benzoate in conc sulfuric acid that
has been cooled in an ice bath.
2. Add 50:50 mixture of conc sulfuric and nitric acids
dropwise with constant stirring, keeping the
temperature below 10 degrees using ice.
3. Stir for 15 mins, then pour over crushed ice in a
beaker.
4. Filter off crystals when the ice is melted.
5. Wash the crystals with water.
6. Recrystallize using ethanol.
What issues do chemists have to consider
when carrying out organic synthesis?
• Stereoisomers:
– Different stereoisomers have different properties.
– Thalidomide is a drug with a chiral centre that was prescribed in the
1960’s.
– One enantiomer helped morning sickness, while the other caused
birth defects.
– Chemist’s can plan the enantiomer produced by understanding the
mechanism of the reaction.
– SN2 nucleophilic substitution can produce a single enantiomer if a
single enantiomer is used as the starting molecule.
• Safety:
– Reagents used in organic synthesis can be dangerous.
– Fume cupboards can be used to remove toxic gases.
– Hot plates can be used to heat flammable reagents so there are no
naked flames near the flammable reagents.
Why is organic synthesis important?
•
Many molecules tested before one is found that is an effective medicine.
– Combinatorial chemistry can be used to speed up the process.
•
•
•
•
•
Hundreds of similar molecules made at once by passing reactants over reagents held on a polymer
support.
Can be automated.
Can be repeated quickly and easily.
Set of compounds created is called a library.
Each chemical in the library is tested to see if it a safe and effective drug.
–
–
•
Analysis of synthesis include IR spectroscopy, mass spectrometry, gas chromatography and UV spectrometry.
Chemical tests can be used to identical functional groups.
Used to research the development of drugs originally made from natural plant
substances from reagents that are more readily available.
– Paclitaxel is a cancer drug that was discovered in the bark of the Pacific Yew tree.
– Pacific Yew trees are a rare, protected species, that are often inhabited by protected birds.
– Chemists developed a new method of synthesising Paclitaxel from the needles and twigs of
other species of yew trees that grow in Europe.
•
Used to change a natural product to improve its qualities or reduce its side effects.
– First form of aspirin was salicylic acid and was derived from the barks of willow trees.
•
Caused mouth and stomach irritation.
– Chemists modified the functional groups to make acetylsalicylic acid.
•
Same painkilling properties but without the side effects.