Jack's Chemistry Revision Notes
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A level chemistry revision notes. From topic 1 to 6.
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 = -50OC, H2Se = -30OC, but H2O = 100OC)
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.
|
lone pair - lone pair repulsion > lone pair - bonding pair repulsion > bonding pair - bonding pair repulsion |
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]
|
Basic element |
Na2O |
Dissolves slowly to form strong alkali |
|---|---|---|
|
MgO |
Dissolves slowly to form weak alkali |
|
|
Amphoteric element |
Al2O3 |
Does not react |
|
Acidic element |
SiO2 |
Does not react |
|
P4O10 |
Forms phosphoric acid soln. |
|
|
SO2 |
Forms sulphuric acid soln. |
|
|
SO3 |
Forms sulphuric acid soln. |
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]
|
NaCl |
Dissolve easily, solution pretty neutral |
|---|---|
|
MgCl2 |
|
|
Al2Cl6 |
Reacts vigorously with water, produces HCl fumes Solution acidic, pH 2 or 3. |
|
SiCl4 |
|
|
PCl5 |
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]
|
Most electronegative: Fluorine Highest boiling point: Astatine Best oxidising agent: Fluorine Best reducing agent: Astatide ions (I/M forces are just Van der Waals) |
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).
|
reagent |
product |
|---|---|
|
NaF |
HF |
|
NaCl |
HCl |
|
NaBr |
HBr and a little Br2 and SO2 |
|
NaI |
I2 and a little HI and S, H2S, SO2 |
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:
state of division (i.e. The surface area, powdered form is faster)
temperature (more energy -> more collisions have required Ea)
(a small increase in T may lead to a large increase in rate)
concentration (increased P(collision))
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 (DHQ) refer to standard conditions: 1 atm, 298K, 1M
Types of Enthalpy [75]
+ve means that energy goes from surroundings to system (i.e. Endothermic)
|
Ionisation enthalpy |
One mole of e- from gaseous atoms |
DHQI 1 |
+ve |
|
Electron affinity |
One mole of e- added to gaseous atoms |
DHQe 1 |
-ve |
|
Lattice dissociation enthalpy |
Ionic solid dissociating to gaseous atoms |
DHQlatt |
+ve |
|
Solution enthalpy |
One mole of ionic solid dissolving in water to make aqueous ions |
DHQsoln |
+ve |
|
Enthalpy of hydration |
1 mole Gaseous ions forming aqueous ions |
DHQhyd |
-ve |
|
Enthalpy of combustion |
Buring 1 mole of reagents in standard states |
DHQc |
-ve |
|
Enthalpy of formation |
Forming 1 mole of something in standard state from elements in standard states |
DHQf |
|
|
Enthalpy of sublimation |
Making 1 mole of gaseous atoms |
DHQsub |
+ve |
|
Enthalpy of atomisation |
Often the same as sublimation. |
DHQdiss |
+ve |
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.
|
bond breaking - endothermic bond making - exothermic |
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]
|
|
Linear |
180 |
|---|---|---|
|
|
V-shaped |
105 |
|
|
Tetrahedral |
109.5 |
|
|
Pyramidal |
107 |
|
|
Planar |
120 |
|
|
Both tetrahedral |
109.5 |
|
|
Trigonal bipyramidal |
90 and 120 |
|
|
Octahedral |
90 |
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.
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A level chemistry revision notes source: http://www.jwhitham.org.uk/alevelchem/html/Chem2.html
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