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Jack's Chemistry Revision Notes

or...

Why you'd have to be mad to study chemistry

VERSION 0.2.0 - 16 MAY 1999

Atomic Structure (1.1)

A - mass number

Z - number of protons

Mass spectrometer [1]

Vaporised sample is ionised, accelerated by an electric field, deflected by a magnet and made to hit a detector.

Order of electron shells [2]

1s 2s 2p 3s 3p 4s 3d 4p

[Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶

Ionisation energy across period 2 [3]

1. Shielding by existing electron orbitals causes this.

eg. Boron shielded by whole 2s orbital.

2. Pair effects

Nitrogen has 3 unpaired p electrons.

Oxygen has 2 unpaired p electrons, and a pair. It is easier to remove one of the p electrons from oxygen than from nitrogen, because there are less unpaired electrons.

• Ionisation energy decreases down a group. (Shielding effect of previous shells)

• Ionisation energy generally increases from left to right across a group, as nucleus becomes more positive but electrons are still at the same distance (greater attraction).

Amount of Substance (1.2)

Misc [4]

Avogadro Constant = 6.0 x 10²³ per mole

Ar = relative atomic mass

Mr = relative molecular mass

Mr is measured relative to Carbon 12.

Ideal Gas Equation [5]

pV = nRT

(R = 8.31JK^-1mol^-1)

Bonding (1.3)

Bond types [6]

Co-ordinate: A covalent bond where both electrons are from the same atom.

Metallic: lattice of +ve ions surrounded by delocalised electrons.

Pauling's scale of electronegativity [7]

Scale from 0.7 to 4.0.

electronegativity is the power of an atom to withdraw e^- density from a covalent bond

Covalent bonds are not always symmetrical: they may be polarised:

• NEGATIVE anions can be polarised by cations (+ve) of high charge density.

  eg. HCl is polar. The H is more positive (d+) than the Cl.

• Ionic bonds have covalent character unless the ions are perfectly spherical. The greater the charge to size ratio of the ions is, the more polarised the ionic bond is, and the more covalent character it will show.

eg. NaCl is truly ionic, MgCl2 has some covalent character, and AlCl3 and SiCl4 are covalent.

Intermolecular bonds: Hydrogen bonding [8]

(lone pair(s) and high electronegativity required for the O atom)

This is why water has an oddly high boiling pt.

(H2S = -50^OC, H2Se = -30^OC, but H2O = 100^OC)

H-bonds are very strong (less than covalent bonds though) and are called permanent bonds.

NH3 can make 1 hydrogen bond, HF can make 3.

How it works: electrons on the H are pulled away by the high electronegativity of the other atom, leaving an exposed side of the proton, which attracts negative areas on other molecules, eg. Lone pairs.

Intermolecular bonds: Dipole-dipole attraction [9]

weaker than H-bonding:-

dipole-dipole attraction between O and C where the dipole is.

Intermolecular bonds: Van der Waals forces [10]

Non-polar molecules do have I/M forces, they are the induced dipole-dipole type.

e^- movement induces a dipole in a molecule. These forces are attracting and breaking all the time.

Bond Rules

All things have Van der Waals I/M forces. If a molecule has a difference in electronegativity across it (i.e. A dipole) then there will also be dipole-dipole attraction. And if one of the atoms has very high electronegativity and lone pairs (N, O, F), and is bonded to a hydrogen atom, then hydrogen bonding will also exist.

Types of crystal [11]

• Ionic crystals (like NaCl) have a large cubic structure.

  (and a high melting point)

• Molecular (giant covalent) eg. Graphite, diamond

  Every atom shares electrons with neighbouring atoms.

• Giant Atomic - just atoms packed as tightly as possible - stacking like balls in pyramid structure. This is either cubic close packing or hexagonal close packing.

  (lower melt point than ionic crystals)

Repulsion rule [12]

A lone pair of electrons produces a greater repulsion effect than a bonding pair.

Periodicity (1.5)

Patterns in period 3 [17]

Period #3 is a bit like period #2.

• Atomic radius decreases from left to right (Na larger than Ar)

  (increased nuclear charge => increased nuclear size => pull on e^- is greater)

• melting pt. increases from Na to Si, then drops for P, increases slightly for S, then decreases again for Ar:

- For metals, melting point increases as the charge/size ratio increases - better metallic bonding.

- Silicon is macromolecular, so it has very strong bonds. (covalent bonds linking all atoms together)

- The final elements are simple molecular: with weak Van der Waals I/M forces.

• electronegativity increases from left to right.

(increasing nuclear charge => electrons attracted more strongly)

• first ionisation energy increases from left to right

  (same nuclear size, but greater nuclear charge, so electrons harder to remove)

Extra Solid Types

PCl5 actually exists as an ionic solid consisting of PCl4⁺ and PCl6 ^-. AlCl3 is a solid dimer at room temperature.

Patterns down group 1 [18]

• radius increases (more e^- shells)

• ionisation energy decreases (more e^- shells)

• electronegativity decreases (greater shielding effect)

• melting point decreases (metal-metal bond strength decreases)

Group I metals react with water to form metal hydroxides and hydrogen:

2Na + 2H2O -> H2 + 2NaOH

Metal oxides and chlorides [19]

MgO, SO2, MgCl2, and AlCl3 can be formed by direct combination (burning).

Reactions of oxides of period 3 elements with water [20]

So, the metals form alkalis, and the non-metals/metalloids either form acids or do not react.

Reactions of chlorides of period 3 elements with water [21]

Additional Notes on periodicity [22]

In each period the oxides of the metals and metalloids have giant structures, whereas the oxides of non-metals are composed of simple molecules.

• The simple molecules (non-metals) react much more readily than giant structures.

From left to right, the oxides of all these molecules change from being ionic, involatile, and metallic, through being giant molecular, involatile and amphoteric, to being simple molecular, volatile, and acidic.

Lithium (thermal stability of compounds) [23]

• Lithium Carbonate is much less stable than other group 1 carbonates.

• Lithium Hydroxide is also less stable.

• LiNO3 decomposes on heating to LiO2(s) but all other group 1 metals decompose to their corresponding nitrates, eg. NaNO2(s).

• Lithium compounds have more covalent character than other Group 1 compounds, eg. LiI is soluble in methanol and propanol.

The Halogens (2.4)

[40]

Products of reactions between NaX and Sulphuric Acid [41]

(where X is a halogen): NaX + H2SO4(conc)

Halide ions can reduce sulphur in H2SO4 by varying degrees. In H2SO4 the S has oxidation state +6. Iodide ions are able to reduce this to +4 (SO2), 0 (S) and -2 (H2S).

NaBr and NaI are oxidised in the reaction, but HF and HCl cannot be oxidised by the acid.

Br2 is a brown gas. I2 is a purple gas, and a black precipitate.

Reactions of Cl

[42] ..with water: chlorine + water -> hydrochloric acid + hypochlorous acid

Cl2 + H2O -> H⁺(aq) + Cl^-(aq) + HClO(aq)

Cl is simultaneously reduced and oxidised - a disproportionation.

The product is also known as chlorine water. It is used to make bleach and treat water, and kills germs.

[43] ..with NaOH

This reaction depends on concentration and temperature.

At room temperature and using dilute NaOH:

Cl2 + 2OH^- -> Cl^-(aq) + ClO^-(aq) + H2O(aq)

Forms Sodium Hypochlorite which is a bleach.

Testing for halides [44]

Silver nitrate (AgNO3) can be used to detect halide ions.

halide ion + silver nitrate -> nitrate ion + silver halide

or, for reaction with X^- (halide)

X^- + Ag⁺ + NO3^- -> NO3^- + AgX(s)

X can be identified from the colour of the silver halide AgX: [45]

AgI(s) is yellow AgBr(s) is creamy colour AgCl(s) is white

X can also be identified by dissolving the silver halide in ammonium hydroxide: [46]

AgI(s) doesn't dissolve AgBr(s) partially dissolves AgCl(s) dissolves fully

Kinetics (3.1)

Maxwell-Boltzmann Distribution [47]

increasing T shifts to right

Area under curve is total number of particles

Reaction rate [48]

Rate = Damount / Dtime or Dconcentration / Dtime

The rate is affected by:

1. state of division (i.e. The surface area, powdered form is faster)

2. temperature (more energy -> more collisions have required Ea)

   (a small increase in T may lead to a large increase in rate)

3. concentration (increased P(collision))

4. catalyst (provides alternate route with lower Ea)

Rate equation [50]

Rate = k [A]^m[B]ⁿ

m and n are the orders of the reaction with respect to reagents A and B.

k is the rate constant.

Thermodynamics & Energetics (4.1 1.4)

Enthalpy [74]

Definition: Enthalpy (H) is the energy content at constant pressure.

Standard enthalpy changes (DH^Q) refer to standard conditions: 1 atm, 298K, 1M

Types of Enthalpy [75]

+ve means that energy goes from surroundings to system (i.e. Endothermic)

Hess's Law [14]

Total Denthalpy is independent of the reaction route taken.

Bond Enthalpy [76]

Bond Enthalpy is the definite amount of energy associated with each chemical bond. Bond enthalpies can be used to predict whether or not, and how easily two substances will react.

example: all the bonds in CH4 are identical C-H bonds with the same bond enthalpy, E(C-H). To break all four, 4E(C-H) is required.

DH ( CH4(g) -> C(g) + 4H(g) ) = 4E(C-H)

Generally, the lower the bond enthalpy, the weaker the bond.

• But Ebond is not constant, it varies according to the environment of the bond. The same bond in different compounds will have different values of Ebond.

When heat or light can break bonds, it will usually break the weakest (lower bond enthalpy) first.

eg. Cracking works because C-C bonds are weaker than H-C bonds (C-C bonds split homolytically to form radicals).

eg. H2 + Cl2 works because Cl-Cl bonds split in the presence of light. (initiation, propagation, termination steps..)

Molecule shapes [16]

Spontaneous reactions [78]

DH, whilst important, is not sufficient to explain spontaneous change.

=> spontaneous exothermic change (eg. burning) makes sense in terms of DH but spontaneous endothermic change does not.

Entropy (S) [79]

UNITS: JK^-1mol^-1 (NOT kJ)

Positive DS for a reaction indicates that the reaction is feasible (ie. spontaneous).

S rules of thumb [80]

• S of gases > S of liquids > S of solids

• ions and molecules in solution have higher S than solids

• larger molecules breaking down into smaller molecules show an increase in entropy

• increasing T increases S

• A change of state has much more effect on S than a change of temperature

• The more complex a molecule is, the higher S is (at same temperature) because there are more bonds.

2^nd law of thermodynamics [81]

DStotal = DSsystem + DSsurroundings

DStotal must be positive because overall entropy is always increasing.

And because DSsystem = DSSproducts - DSSreactants

and DSsurroundings = - DH / T

the feasibility of a reaction can be determined by calculating the DSsystem and DSsurroundings from databook values. If DStotal is positive, the reaction is feasible at that temperature. Thus the feasibility of a reaction depends on the balance between entropy and enthalpy. [82]

Gibbs free energy change [83]

For a feasible reaction DG must be zero or negative, note this is the opposite way round to DS.

Equilibria (4.2 2.1)

Many reactions are reversible. At equilibrium, both forward and reverse reactions are proceeding. For a homogenous system in equilibrium, where: A + B <-> C + D,

Kc = Right / Left = [C]eq[D]eq / [A]eq[B]eq

[25] Le Chatelier's principle: The system resists change.

=> Increase conc. of something, it reduces it.

• Catalyst has no effect on Kc.

• Changes in overall pressure have no effect on Kc.

But Kc depends on temperature. The effect of T can be clearly defined:

DG = - RT ln K

• where R = gas constant (8.31)

Partial pressure and Kp [85]

=> partial pressure is analogous to concentration

The partial pressure of a gas in a gas mixture is the pressure that would be exerted if that gas alone filled the whole volume occupied by the mixture (Dalton's law)

partial pressure PA = Ptotal x mole fraction

(PA of oxygen in air is 0.2 atm, as this is 1/5 of 1 atm (4 parts nitrogen, 1 part oxygen))

Kp is the equilibrium constant for a gas system:

Kp = Right / Left = P(C)eqP(D)eq / P(A)eqP(B)eq

Note that changing pressure of the system does not affect KP, but changing temperature does affect KP.

Contact Process [26]

• Makes SO3 from SO2 by adding O2.

  Low temperature, high pressure, platinum/manganese catalyst

  (typically 450^OC, and a few atm - not really low T/high p, but obtains high yield at low expense)

Haber Process [27]

• makes ammonia from N2 and H2. Low T and high P.

  (typically 500^OC, and 100-300 a.t.m., with iron catalyst)

  The iron catalyst is supported on silica.

Acid-Base Equilibria (4.3 2.2)

Bronsted-Lowry: acid is a proton donor, base is a proton acceptor

In any acid-base reaction a proton is transferred.

pH [30]

pH = - log [H⁺] (H⁺ in moldm^-3)

Dissociation of water [31]

Water is very weakly dissociated: Kc = [H⁺][OH^-] / [H2O]

The ionic product of water, Kw = [H⁺][OH^-]

At 25^OC only, Kw = 10^-14 mol²dm^-3.

pH of a strong base found by finding [H⁺]: Kw = [H⁺] x [OH^-].

• weak acids dissociate only partially in aq. soln.

• strong acids/bases are fully dissociated: and fully ionised in solution

>>> Dissociations are always written with ions on the right.

Therefore, as Kc = right / left, Ka, Kc etc are always ions / molecules.

- Ions Right -

Monoprotic weak acids

A weak acid is only partially dissociated: HA <-> H⁺(aq) + A^-(aq)

The dissociation constant, Ka, is found by:

Ka = [H⁺(aq)][A^-(aq)] / [HA]

The higher the value of Ka, the stronger the acid.

Diprotic weak acids

These are like monoprotic weak acids, except they dissociate twice, producing 2H⁺.

One example is H2S(aq):

H2S <-> HS^-(aq) + H⁺(aq)

HS^-(aq) <-> S^2-(aq) + H⁺(aq)

There are two values of Ka, Ka1, and Ka2:-

Ka1 = [HS^-][H⁺] / [H2S]

Ka2 = [S^2-][H⁺] / [HS^-]

Ka2 is always much less than Ka1.

Finding Ka

Ka may be found by this approximation, as [H⁺] Z [HA^-], and [H⁺] Z 0. (as very few H⁺ come from the water that the acid is in)

Ka Z [H⁺]² / [HA]

Finding pH

and so, if [HA] and Ka are known for a weak acid, it's pH can be calculated:

[H⁺] Z †( [HA] x Ka )

=> pH Z -log10 †( [HA] x Ka )

pKa

pKa is a pH-like scale for Ka.

pKa = -log10 Ka

pH and pKa

You can combine the above two equations to make:

Weak bases

Just as Ka = [H⁺][A^-] / [HA] for weak acids,

Kb = [B⁺][OH^-] / [BOH] for weak bases.

Kb Z [OH^-]² / [BOH]

pH Curves

All acid/base titrations produce a pH graph like the one on the left. (a mirror image for an acid added to an alkali)

The endpoint is where the titration should finish: where equal amounts of acid and base have reacted.

Also known as the equivalence point where acid and base are equivalent.

Indicators change colour within a narrow pH range. The right indicator has to be used to show when the titration reaches the endpoint. This isn't always pH = 7 though.

For a strong acid and strong alkali, endpoint is about pH 7.

For a weak acid and strong alkali, endpoint is greater than pH 7.

For a strong acid and weak alkali, endpoint is less than pH 7.

For a weak acid and weak alkali, endpoint is about pH 7.

The indicator should show when the mixture is close to the endpoint.

Buffers

• an acid buffer consisting of a weak acid (eg. Ethanoic acid) and a salt of the acid (eg. Sodium ethanoate).

• An alkaline buffer consisting of a weak alkali (eg. ammonium hydroxide) and a salt of the alkali (eg. ammonium chloride).

Acid buffer: If H⁺ ions are added, they will react with the ethanoate ions from the salt (because there are more of them than ethanoate ions from the partially dissociated acid). Undissociated ethanoic acid will form. If OH^- ions are added, they will combine with the H⁺ ions to form water, and more of the acid will dissociate.

Alkaline buffer: same sort of mechanism

Uses of buffers

• calibration of pH meters

• biology eg. Maintaining correct conditions for enzymes

• medicines

• dyes

• electroplating

• injections.

Redox Titration Facts

MnO4 reduces to Mn²⁺

Cr2O7 ^2- reduces to 2Cr³⁺

Redox Equilibria (4.4)

The anode is where the oxidation reaction occurs.

Usually, electrodes are a sample of a metal in a solution of the same metal ions. If M = Zn and soln is ZnSO4(aq):

Zn(s) | ZnSO4(aq)

or

Zn(s) | Zn ²⁺(aq)

both describe this electrode.

Gas electrode

Sometimes a gas electrode must be used. Here the gas under test is bubbled over a platinum electrode in a solution of ions of the same element as the gas. eg. The standard hydrogen electrode is described as:

Pt (s) | H2(g) | H⁺(aq)

Standard

The standard for half cells is that everything happens under standard conditions (1 atm for gases, 298K, and 1M solutions).

Redox electrodes

This is where something in solution is either oxidised or reduced, but remains in solution. eg.

Pt (s) | Fe²⁺(aq) Fe³⁺(aq)

-> a solution containing 1M iron (II) and 1M iron (III) ions. e^- produced by oxidation of Fe²⁺ or lost by reduction of Fe³⁺ enter or leave the solution by the Pt electrode.

Cell reactions - standard layout

Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s)

• The | is a phase boundary.

• The || indicates a salt bridge.

The oxidation cell is put on the left. The overall reaction where oxidation of the left cell occurs is feasible if the left cell potential is less than the right cell potential.

Word Association

Oxidation occurs at the Anode on the Left which is Negative

POSITIVE POTENTIALS OXIDISE OTHERS

ie. the best oxidising agents have the highest E values.

Standard Cell Potential

Standard cell potential, E^Q is measured with the Standard Hydrogen Electrode (SHE) as the left hand electrode, and standard conditions. Cells with high E values are the best oxidising agents: they accept e^- easily.

Calculating E

Ecell = Eright - Eleft

Ecell is usually calculated by putting the most positive cell on the right.

Secondary standard

A calomel electrode with an E^Q value of +0.27V is often used in place of the SHE.

POSITIVE POTENTIALS OXIDISE OTHERS

Principles of Catalytic Action (4.5)

Heterogeneous catalytic action

A heterogeneous catalyst is in a different state to the reagents. One (or more) of the reagents is adsorbed onto the catalyst surface, making it more likely to react with the others. For example, it may become more accessible to collisions, or be held in a particularly active configuration, or be broken into more reactive fragments.

The strength of adsorption is very important. Tungsten adsorbs too well. Silver adsorbs poorly. So Ni and Pt are more commonly used.

A support medium minimises the amount of catalyst while maximising the catalyst surface area.

Homogenous catalytic action

The reaction goes through an intermediate species. If a transition metal catalyst is used, it's oxidation state changes.

Specificity

Catalysts are often highly specific. Enzymes will work for only one reaction. Acid/base catalysts (eg. Acid used to catalyse hydrolysis of an ester) are much more general.

Some Catalysts [35]

Transition metals often make good catalysts (change in oxidation state)

Iron -> Haber process

MnO2 -> decomposition of H2O2

Ni -> margarine (hydrogenation of vegetable oil)

Extraction of Metals (5.1)

Extraction of metals usually involves reduction of a metal oxide. There are several reduction methods: the one used depends on the cost of the reductant, the energy requirements, and the purity of the metal required.

Reduction of metal oxides with carbon

eg. production of iron.

1. Iron oxide is first heated with coke (mostly carbon) in a sintering plant.

2. Poured into a blast furnace along with more coke.

3. The carbon reduces the iron oxide to iron:

2Fe2O3 + 3C -> 4Fe + 3CO2

and also, CO produced by incomplete combustion of coke:

Fe2O3 + 3CO -> 2Fe + 3CO2

4. The impurities (slag) float on top of the liquid iron poured out of the blast furnace.

Waste gases are used to heat the furnace.

Iron can be further purified by:

• bubbling oxygen through it in a converter to remove carbon

• adding magnesium to remove sulphur

• adding lime (CaCO3) to remove other impurities (reacts with SiO2, the main impurity, producing slag, CaSiO3)

Carbon reduction is not perfect for all metals, as with some others carbides are formed.

Reduction of metal oxide with an active metal

Mainly used for chromium and in Thermite process. Advantages:

• no carbides formed

• lower temperatures than carbon reduction

• produces small quantities of very pure metal

Sort of substitution of one metal for another:

eg. Cr2O3(s) + 2Al(s) -> 2Cr(s) + Al2O3(s)

Reduction of metal oxide by electrolysis

Used for aluminium extraction because Al2O3 is too stable for carbon reduction: the temp would have to be very high. The bauxite is dissolved in molten cryolite (Na3^.AlF6) to lower it's melting point from 2000^OC to 950^OC. Carbon electrodes are used.

Cathode: Al reduced from Al³⁺ to Al

Anode: Oxide ions oxidised to oxygen

This is a continuous process using much electricity. Cryolite is damaging to the environment.

Reduction of a metal halide by another metal

Used for production of Ti, which is brittle unless very pure. This process produces very pure metal in quantity. Main Ti ore: TiO2.

TiO2 reacted with C and Cl at 900^OC to form TiCl4:

TiO2 + 2C + 2Cl2 -> TiCl4(l) + 2CO

TiCl4 is a dangerous product, it hydrolyses easily and forms HCl fumes in moist air.

It is reacted with Na: (at 550^OC)

TiCl4(l) + 4Na(s) -> 4NaCl + Ti(s)

This is also a dangerous reaction as it's highly exothermic. (temp rises to 1000^OC) The temperatures are so high that the reaction must happen in an inert atmosphere of argon so that titanium oxides are not formed again. It's all very expensive.

• dangerous process

• high temperatures needed

• sodium and chlorine needed

• inert atmosphere used

additional: H2 can be used as a reducing agent for metals: no carbides are formed.

Transition Metals (5.2 2.3)

General Properties [34]

1. coloured compounds

2. variable oxidation state

3. form complex ions

4. show catalytic activity

A transition metal has an incomplete d sub-shell in at least one of it's oxidation states. The properties arise from this.

The odd order of e.c.'s [37]

• Sc is [Ar] + 4s² + 3d¹

Most of them add 1 normally, eg.

• Ti is [Ar] + 4s² + 3d²

except Chromium and Copper

• Cr is [Ar] + 4s¹ + 3d⁵

• Cu is [Ar] + 4s¹ + 3d¹⁰

On ionisation, the 4s shell goes first, then the 3d shell.

Complex formation [39]

complex: central metal ion surrounded by ligands

co-ordination number: number of ligands

[Co(H2O)6]²⁺ and [Cu(H2O)6]²⁺ are octahedral.

[CoCl4]²⁺ and [CuCl4]²⁺ are tetrahedral.

The Colour Table

The colour of ions

is affected by a change in ligands, a change in co-ordination number, or change in oxidation state.

Ligands

A ligand is an electron pair donor.

Unidentate: one tooth: NH3, Cl^-

Bidentate: two teeth: ethanedioate ions

Polydentate: many teeth: EDTA, and haemoglobin

Reactions of inorganic compounds in aqueous solutions (5.3)

A Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor.

(A transition metal is one with a partially filled d sub-shell)

Metal ions form metal-aqua ions in water: [M(H2O)6]²⁺.

These ions can be present in the solid state: FeSO4.7H2O.

Acidity reaction (hydrolysis)

• M³⁺ ions are more acidic than M²⁺ ions (M³⁺ ion has greater polarising power)

• M³⁺ precipitates are amphoteric

• CO3^2- ions react with H3O⁺ to form carbonic acid. M²⁺ are less acidic than this, and so are unable to displace the CO2 from it, and that's why they form metal carbonates, whereas M³⁺ are more acidic, and form CO2 and metal hydroxides.

with water...

M²⁺ and M³⁺ ions can lose protons to water to

the extent that they have +1 charge:

[M(H2O)6]²⁺ + H2O <-> [M(H2O)5(OH)]⁺ + H3O⁺

with OH^- ions...

M²⁺ and M³⁺ ions lose protons as before, but

this time can go a step further

and form a metal hydroxide precipitate.

[M(H2O)4(OH)2]⁺ + OH^- <-> [M(H2O)3(OH)3] + H2O

M³⁺ pptes can redissolve in OH^- ions to form

[M(H2O)2(OH)4]^- ions.

M²⁺ pptes do not redissolve.

with CO3^2- ions...

M²⁺ ions form insoluble metal carbonates

(MCO3)

M³⁺ ions form metal hydroxides and CO2 gas

with NH3 ions...

M²⁺ and M³⁺ form hydroxides and then

substitution occurs.

Metal Hydroxide Table

Substitution Complex Ion Table

• Metal hydroxides are always initially formed. But with some bases in excess, ligand substitution will occur.

The Extra Bastard Hard Evil Table

• Vanadium may be reduced by zinc in dilute acid.

• These are all colours of -ate ions, I think.

• Manganese may be reduced by iron

Silver complexes

Silver ions produce linear complexes.

[Ag(NH3)2]⁺ - used in Tollen's reagent

[Ag(S2O3)2]^3- - used in photography

[Ag(CN)2]^- - forms when silver salts are dissolved in CN^- solutions:

used for electroplating

Transition metal complexes in the body

Haemoglobin is an iron (II) complex. It is surrounded by 4 porphyrin ligands, a globin ligands, and water or oxygen as the sixth ligand.

Cisplatin is an anticancer drug: a 4 co-ordinated complex of Pt²⁺, ligands are two Cl^- ions and two NH3 atoms.

Miscellaneous things

In alkaline solution, Co (II) and Cr (III) are oxidised by H2O2, which acts as an oxidiser or a reducer depending on the conditions.

Co (II) is oxidised by air in ammoniacal solution.

Organic Chemistry (3.2)

Cracking (thermal) [51]

• Free radical reaction

• usual catalyst is aluminium oxide at 450^OC

• involves homolytic fission of C-C bonds only

Main fractions of crude oil [52]

• petrol (C₅-C₁₀)

• kerosene (C₁₁-C₁₂)

• diesel (C₁₃-C₂₅)

• chemical feedstock - Naphtha

• refinery gas

• residue

Pollutants made by cars [53]

• NO_x, CO, unburnt hydrocarbons

• Catalytic converters exist.

Haloalkanes

• undergo nucleophilic substitution reactions (susceptible to nucleophilic attack) [61]

• with OH^- ions, and cyanides (NaCN, KCN) in aqueous solution:

aq. soln

warm

this mechanism also applies to OH^- ions.

The end result of the above (bromomethane + NaCN) is ethanenitrile, NOT methanenitrile, because the extra C counts as part of the carbon chain.

• with ammonia: where the lone pair attacks. Product: amines [62]

(haloalkane must be in excess, or secondary amines are formed (multiple substitution))- this makes it less likely that the amine produced will react with Br ^-, also warm

The NH3 reacts with the acid formed (HBr here): NH3 + HBr -> NH4Br

overall: RX + 2NH3 -> RNH2 + NH4X

(they often want this overall equation for the reaction)

Fluoroalkanes are very unreactive (C-F bond v. strong)

Alcohols [63]

• Made by hydration of ethene (phosphoric acid catalyst, 300^OC, 65 atm)

• Also by fermenting glucose - impure product

Alcohols: Suitable oxidising agents [64]

1. Acidified solution of potassium dichromate (VI) (K2Cr2O7)

- conditions: warm (except for making carboxylic acids from 1^O alcohols, in which case reflux)

- Cr2O7 ^2- ions change from orange to green as they are reduced in the reaction.

1. 3^O alcohols "are not easily oxidised".

Alcohols: Dehydration of alcohols to form alkenes [65]

an elimination reaction, reagents conc. sulphuric or conc. phosphoric acid @ 180^OC.

1. acid protonates the lone pair on the OH group

2. alcohol loses a molecule of water, left with carbonium ion

3. carbonium ion loses proton to produce an alkene.

The acid is actually a catalyst as it is regenerated.

Carbonyl compounds (having a =O but not -OH) [66]

aldehydes/ketones distinguished by

• Tollen's (silver mirror), reagent contains [Ag(NH3)2]⁺, this is reduced to silver ppte

  Silver mirror formed if an aldehyde is present

• Fehling's solution - normally deep blue, goes brick red with aldehyde. Contains copper (II) complex, that is reduced to red CuO2(s) by aldehyde.

[67] Carbonyl compounds may be reduced by (a) NaBH4 or (b) H2 gas + Nickel catalyst

symbol for reduction is [H]

• NaBH4 will not reduce alkenes, so it can be used for selective reduction of aldehydes/ketones.

• NaBH4 is sodium tetrahydridoborate (III)

• NaBH4 works in the presence of methanol

Carbonyl reaction with HCN (hydroxynitriles formation) [68]

works for aldehydes and ketones. Reagents: KCN, dilute sulphuric acid

The product is a hydroxynitrile, e.g. 2-hydroxypropanenitrile here.

The carbonyl compound actually reacts with HCN gas which is a product of KCN + acid.

Carboxylic acids + esters [69]

acid + alcohol <-> ester + water

in the presence of a strong acid catalyst (e.g. sulphuric)

ethylmethanoate

(methanoate group on left)

• Esters are used as solvents, artificial flavourings, and polymerised to make synthetic fibres. [70]

• Natural esters can be hydrolysed to other useful products eg. Soap, glycerol, higher fatty acids. [71]

Isomerism [72]

Structural isomer - same molecular formula, different structure

Optical isomer - same structure, not spatially identical

(aka enantiomers, stereoisomers). Rotate light in different dir.

Also geometric isomerism (cis/trans) which are also aka. stereoisomers

Additional [a]

CH4 + Cl2 does not produce any H^" radicals.

H2 + Cl2 does.

Aromatic Chemistry (6.1)

Benzene - an arene

• bond length is between single and double bond length [59]

• delocalised, therefore stable (better than Kekule benzene)

All reactions are electrophilic substitution:

A p complex ion (in the middle) is always formed.

Nitration of benzene [60]

Reacts at 50^OC with conc. H2SO4 and HNO3, sulphuric acid is catalyst.

HNO3 + 2H2SO4 <-> NO2⁺ + 2HSO4^- + H3O⁺

Friedel-Crafts Alkylation/acylation

-> where an alkyl/acyl group is added to a benzene ring.

A Friedel-Crafts carrier catalyst is needed, such as AlCl3. (Cl may be any halogen)

The acyl/alkyl group needs to be bonded to the same halogen, eg. Chloromethane.

The carrier catalyst is a Lewis acid, so it will accept a lone pair from the halogen in the alkyl/acyl group (shown as R here):

d+ d-

R - Cl: -> AlCl3

Result: R⁺AlCl4^-

That makes the R group much more d+, so that the benzene ring can attack it:

The AlCl3 is regenerated as the AlCl4^- is attacked by the H⁺ ion to form HCl.

-> Ethene may be used in place of RCl. This reaction is used to make ethylbenzene (styrene) used for making polystyrene. An additional HCl catalyst is needed, temp. 90^OC

Sulphonation

Very concentrated (fuming) sulphuric acid contains the electrophile SO3. When

• refluxed together for several hours

• mixed with concentrated HCl at 80^OC,

  benzene and concentrated sulphuric acid form benzenesulphonic acid.

Amines (6.2)

General formula: R-NH2

R can be an alkyl or an aryl (benzene) group. If the amine group is the main functional group, then compounds are named as -amine, eg. Phenylamine. If not, then they are named amino- eg. 2-aminopropane.

Named like alcohols. You can have primary, secondary and tertiary amines.

Amines have base properties. They are lone pair donors.

Methylamine is the best base: the availability of a lone pair is increased by the methyl group, because that group is an electron releasing group: it does not pull on the electrons.

Phenylamine is the worst base: the lone pair on the N is less available, as it is delocalised along with the e^- in the benzene ring.

Ammonia is somewhere between the two, in terms of basic strength.

Reactions of amines

The C-N bond in amines is weak compared to the C-O bond in alcohols, and it's also less polar. Amines tend to react by nucleophilic substitution.

eg. reaction with haloalkanes.

R-H2N: -> R'-Br -> R-H2N⁺-R' + Br^- -> R-HN-R' + HBr

(in this reaction it is often important to show what the HBr then does (ie. reacts with the amine) as an overall equation)

The reagent was a primary amine. The product is a secondary amine.

Reaction with haloalkanes adds to the order of the amine:

ammonia + haloalkane -> 1^O amine

1^O amine + haloalkane -> 2^O amine

4^O amines exist as salts: quaternary ammonium salts, used to make cationic detergents.

The problem with this reaction is that several products are formed, the product reacts again. This may be limited by putting the haloalkane in excess.

Cationic Detergents from quaternary ammonium salts

usually a long hydrocarbon chain, with a positively charged organic group.

Found in nappy cleaners, hair and fabric conditioners. Have poor cleaning properties but good for germicidal use.

Preparing aliphatic amines

Reduction of a nitrile is one method:-

R - CN -> RCH2NH2

Reducing agent: H2 with Ni catalyst (same as reduction (hydrogenation) of ethene to ethane)

Preparing aromatic amines

Usually by reduction of an appropriate nitro compound:

eg. nitrobenzene + [H] -> phenylamine

Same reducing agents will work, also tin in HCl.

Acid chlorides

These reaction by nucleophilic addition elimination, the mechanism's result is similar to substitution, but it's not the same. (Thing to remember: electrons move to the O on the double bond when the C atom is attacked)

In the above mechanism, water reacts with an acid chloride to form a carboxylic acid.

If one of the H's in water was replaced with an R' group (making it an alcohol, HOR') the result would be an ester RCOOR'.

The mechanism will also work for ammonia, making RCOONH2, which is called an amide (on left), and also with amines (general form NHR') to form RCOONHR': a substituted amide.

Reactions of acid anhydrides

This is what an acid anhydride looks like (right).

These are less reactive than acid chlorides. They exist as resonance hybrids: the p electrons can move around between the C-O and C=O bonds.

Acid anhydride + water -> carboxylic acid

Acid anhydride + alcohol -> carboxylic acid + ester

Acid anhydride + ammonia -> carboxylic acid + amide

Acid anhydride + amine -> carboxylic acid + substituted amide

Acid anhydrides may sometimes be used in place of acid chlorides (both are ethanoylating agents): this is better because they are cheaper and react less vigorously, and are therefore safer.

Alkene chemistry (6.3)

Note: Alkenes are unsaturated because of the C=C bond

• Alkenes form addition polymers

Inductive effect

Alkenes react by electrophilic addition. A d+ species (eg. H in HBr) adds to the alkene chain on either end of the double bond as that double bond breaks. A carbonium ion is formed at the other end as the bond has broken. (avoid ever saying C⁺ ion: carbonium is better)

Markovnikov: The carbonium ion is more likely to form on the end of the bond where it will be surrounded by the most alkyl groups. Alkyl groups exert an inductive effect; pushing e^- towards the carbonium ion, which makes it more stable, and likely to exist for longer, so that the d- species will be able to bond to it to complete the reaction.

So, in the reaction between propene and HBr, 2-bromopropane is more likely to be formed than 1-bromopropane.

Reaction of alkenes with hydrogen

(hydrogenation: adding hydrogen to ethene (margarine reaction))

Nickel catalyst at 150^OC

hydrogenation is NOT hydration. Hydrogenation = adding hydrogen

Reaction of alkenes with bromine

• alkenes decolourise bromine (it adds to them)

• adding Br2: (electrophilic addition)

Happens at room temperature.

• Bromine and hydrogen do not have dipoles, but in the above to reactions, one is induced by the presence of the double bond p electrons. nb. This doesn't work for Chlorine because chlorine's e^- cloud cannot be distorted as easily as Bromine's. So the reaction alkene + Cl2 will not happen

• HBr will also add to alkenes. H is the electrophile. Reaction happens in gas phase or in concentrated aqueous HBr soln. (I think the aq. bit is important)

Reaction of alkenes with steam (hydration)

hydration. Alkenes react with steam at 300^OC with Phosphoric acid catalyst,

to form alcohols. PRESSURE: 65 atm. (this is the industrial way to make alcohols)

Reaction of alkenes with sulphuric acid

Propene reacts with concentrated sulphuric acid to form propyl hydrogen sulphate. The electrophile is H⁺. Product drawn on right:-

Useful for making alcohols. The product hydrolyses (reacts with water) easily.

Propyl hydrogen sulphate + water -> propan-2-ol + sulphuric acid

The sulphuric acid is regenerated.

Making alkenes from haloalkanes: elimination reaction

In aqueous solution haloalkanes react with OH^- ions to form alcohols, as the OH^- ions act as nucleophiles. In ethanolic solution the OH^- ions act as a base, and alkenes are formed.

OH^-(eth) could be produced by dissolving KOH in ethanol. Other products would be KBr, and water.

In this reaction a little propanol will always be formed even though the OH^- ions are not in aqueous solution. The OH^- ions have some nucleophile properties. The reaction type that is favoured depends on:-

• the type of haloalkane (3^O favour elimination type)

• base strength of nucleophile: increased by ethanol, reduced by water

• reaction temperature (high temperatures favour elimination)

Epoxyethane

epoxyethane is manufactured by direct synthesis of ethene and oxygen at 250^OC and 1-2MPa with a silver-based catalyst. It is a colourless gas, flammable, explosive and very toxic.

It is a very reactive substance due to the strain on the 3 membered ring (-C-O-C-). The ring is an area of high electron density, so epoxyethane is susceptible to electrophilic attack. Breaking the ring releases much energy.

• It hydrolyses (not hydration reaction) to make ethane-1,2-diol (antifreeze), usually done at 60^OC.

• It is used to make polymers (terylene)

• reacts with alcohols to make alkoxyalcohols - like diols, but one of the OH groups is an OR group.

• Alkoxyalcohols used for solvents in paint industry and also plasticisers, brake fluid, and detergent manufacture (apparently that's only long chain alcohols though).

Structural Determination (6.4)

NMR: Nuclei with an odd mass number have a property known as spin.

Variable magnetic field used: nuclei align with the field (a state) or against it (b state). They may resonate from one state to the other when subjected to radio waves. A constant frequency radio transmitter is used, and the radio wave energy absorbed by resonance is detected by a detector coil. Samples rotates at about 25-35 rotations per second to average out any variations in the magnetic field.

The substance being tested is dissolved in tetrachloromethane, for examination to find ¹H nuclei (it contains no ¹H). TMS is mixed with this to provide a reference point for spectrometer calibration. TMS is ideal because:

• the ¹H nuclei are well shielded, so they resonate well above most organic substances

• it gives one sharp easily detected spectrum caused by the 12 equivalent ¹H.

• Non-toxic

• cheap

• non-volatile

Chemical shift

d = ( external field at resonance of TMS - external field at resonance of sample ) / external field at resonance of TMS

Chemical shift due to a ¹H nucleus depends mainly on how well it is shielded by e^- density around it. (e^- create a small opposing magnetic field) The H in O-H is less shielded than the H in C-H, because O is more electronegative than C and pulls the e^- away from the H.

...and then, there's singlets, doublets, triplets, quartets and so on. But that's not so hard.

Fingerprint region less than 1500cm^-1.

Mechanism notes

• Curly arrows good. Straight arrows bad.

• NMR questions often would like a structure

• Draw lone pairs on everything where -O: -> H etc.

• Don't just say "addition", say "nucleophilic addition" or whatever

• Show the structure when writing out chemical formulae: CH3CH(OH)CH3 (propan-2-ol)

• In equations, it's a good idea to put the catalyst over the ->

The Reactions and Conditions Table

** end of chemistry **