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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