Notes made based on:
CGP AS and A2 OCR Chemistry textbook
OCR AS Chemistry Student Book
OCR Specification
Basically a condensed version of all of these cutting out the BS that's not needed; learn these by heart and do some past papers and that grade A/B is yours.
I used these for the July 2015 C...
AS Chemistry Unit F321: Atoms, Bonds and
Groups
Module 1: Atoms and Reactions
Topic 1: Atoms Subatomic Relative Relative
Particle Charge Mass
Atomic Structure Protons +1 1
Neutrons / 1
● Majority of mass = nucleus Electrons -1 1/2000
● Orbitals = most volume
● Diameter of nucleus v small compared to whole atom
● ISOTOPES: (of an element) atoms w/ same no. of protons but dif no. of neutrons
● Atomic structure model changed throughout the years; Greeks (thought all matter made of
th
invisible particles) -> Dalton (start of 19 century, described atoms as ‘solid spheres’ dif
spheres = dif elements) -> Thompson (1897, not solid, proved the subatomic particles
existence, ‘plum pudding model’) -> Rutherford (gold foil experiment, fire alpha particles at
thin sheet of gold, passed straight through few deflected back, made the nuclear model w/
‘cloud’ of electrons)-> Bohr Model (electrons can’t be in clouds etc.) now used today
Relative Masses
● Are masses of atoms compared to Carbon-12
● RELATIVE ATOMIC MASS (Ar): weighted mean mass of an atom of an element in
comparison to one twelfth the mass of a C-12 atom
● RELATIVE MOLECULAR MASS (Mr): weighted mean mass of a molecule in comparison to
one twelfth the mass of a C-12 atom
● RELATIVE FORMULA MASS: weighted mean mass of a formula unit in comparison to one
twelfth the mass of a C-12 atom
● RELATIVE ISOTOPE MASS: mass of an atom of an isotope of an element in comparison to
one twelfth the mass of a C-12 atom
● Ar and isotopic abundances;
35 37
e.g. 76% Cl , 24% Cl
1. Multiply each relative isotopic mass by percentage and add results up:
(76 x 35) + (24 x 37)
2. Divide by 100:
(76 x 35) + (24 x 37) / 100
=35.5 (1dp)
Topic 2: Moles and Calculations
The Mole
● AM OF SUBSTANCE: quantity whose unit is the mole. Used as a means of counting atoms.
● MOLE: am of any substance containing as many particles as there are carbon atoms in
exactly 12g of a C-12 isotope.
●
23 -1
AVAGADRO CONSTANT (NA): no. particles p/ mole (6.02 x 10 mol ).
●
-1
MOLAR MASS: mass p/ mole of a substance (gmol ).
●
23
No. of moles = no. particles / 6.02 x 10
Empirical and Molecular Formulae
● EMPIRICAL FORMULA: simplest whole no. ratio of atoms of each element in a
compound.
● MOLECULAR FORMULA: actual no. of atoms of each element in a compound.
● Calculating empirical formula:
, e.g. Hydrocarbon burnt in excess oxygen, 4.4g of carbon dioxide and 1.8g of water is made.
What is empirical formula of hydrocarbon?
(only care about the hydrogen and carbon as they are elements present in hydrocarbon)
1. Find out no. of moles of products present:
2. Ratio:
3. Divide by smallest:
● Calculating molecular formula:
e.g Molecule has empirical formula of C4H3O2 and molecular mass of 166g. Work out
molecular formula.
1. Find empirical mass:
2. Divide molecular by empirical mass:
3. Multiply empirical by units:
Chemical Reactions
● Balance equations
Calculation of reacting masses, mole concentrations and volume of gases
Mass calculation:
● Gas calculation:
● Concentration calculation:
● 3
CONCENTRATION: (of a solution) is the am of solute (mol) is dissolved p/ dm of solution.
●
3
Concentrated > 10dm dissolved
● Dilut < 10dm3 dissolved
Topic 3: Acids
Acids and Bases
●
+
Acid = releases H ions in aq solution, proton-donor
● Common acids: HCl, H2SO4, HNO3
●
+
Bases = remove H ions from aq solution, proton-acceptors
● Common bases: metal oxides, metal hydroxides, ammonia
●
-
Alkali = soluble base, releases OH ions in aq solution
● Common alkalis: NaOH, KOH, NH4
Salts
●
+
SALT: any chemical compound formed from an acid when a H ion from the acid has been
+
replaced by a metal ion or another positive ion such as the ammonia ion NH 4 .
, ●
+ -
Base readily accepts H ions from acid e.g. OH forming H2O
● ANHYDROUS: substance w/ no water molecules
● HYDROUS: substance w/ water molecules
● WATER OF CRYSTALISATION: water molecules that form an essential part of the crystalline
structure of a compound.
● Calculating formula of hydrated salt:
e.g heating 3.210g of hydrated magnesium sulfate, MgSO4 . XH2O, forms 1.567g of
anhydrous magnesium sulfate. Find value of X and write the formula of the hydrated salt.
1. Find no of moles of water lost:
2. Find no of moles of anhydrous salt:
3. Ratio of water to anhydrous salt and divide by smallest no.:
● Titrations – find out exactly how much of an acid needed to neutralise a quantity of alkali.
● Methyl orange – turns yellow to red when adding acid to alkali
● Phenolphthalein – turns red to colourless when adding acid to alkali
● Universal indicator no good as colour change too gradual
Topic 4: Redox
Oxidation Number
● Oxidation Is Loss Reduction Is Gain (in terms of electrons)
● When happen together = redox reaction
● Oxidising agent = accepts electrons and is reduced
● Reducing agent = donates electrons and is oxidised
● Oxidised (oxidation number increase) reduced (oxidation number decreased)
● Roman numerals = oxidation number e.g copper has 2+ in copper(II) sulfate
Redox Reactions
● Metals – generally form positive ions by losing electrons w/ increase in oxidation no.
● Non-metals - generally form negative ions by gaining electrons w/ decrease in oxidation no.
● Redox reactions of metals w/ dilute acids (hydrochloric and sulfuric) – metal ions oxidised,
lose electrons and form soluble metal ions, hydrogen ions are reduced, gaining electrons and
forming hydrogen molecules.
Module 2: Electrons, Bonding and Structure
Topic 1: Electron Structure
Ionisation Energies
● FIRST IONISATION ENERGIES: energy required to remove one electron from each atom in
one mole of gaseous atoms to form one mole of gaseous 1+ ions.
, ● SUCCESIVE IONISATION ENERGIES: energy required to remove one electron from each
atom in one mole of gaseous atoms to form one mole of gaseous 2+ ions.
● Ionisation energy influenced by:
1. Nuclear Charge – more protons
= more positively charged =
more nuclear attraction
2. Atomic radius – more distance =
less attraction
3. Electron shielding – as no. of
electrons between nucleus and
outer electron increases, outer
electron has less attraction
towards the nucleus charge.
● High ionisation energy = high attraction
between electron and nucleus
● From graph of ionisation energy can tell
which group element is in – no. of
electrons before first big jump in energy
Electrons: electronic energy levels, shells, sub-shells, atomic
orbitals, electron configuration
● Electrons -> shells, given numbers called principal quantum
numbers
● Further from nucleus = greater energy level
● Each shell = dif sub-shells = dif no. of orbitals
● No. of electrons in each type of sub-shell:
Sub-shell No. of orbitals Max electrons
s 1 1x2=2
p 3 3x2=6
d 5 5 x 2 = 10
f 7 7 x 2 = 14
● Sub-shells and electrons in first four energy levels:
Shell Sub-shell Total no. of electrons
st
1 1s 2=2
nd
2 2s 2p 2+6=8
rd
3 3s 3p 3d 2 + 6 + 10 = 18
th
4 4s 4p 4d 4f 2 + 6 + 10 + 14 = 32
● Orbital = 2 electrons, bit space electrons occupy,
opposite spins, orbitals in same sub-shell have
same energy
● s-orbitals = spherical, p-orbitals = dumbbell, 3 at
right-angles to each other Dot-and-cross diagram
Topic 2: Bonding and Structure
Ionic Bonding
● IONIC BONDING: electrostatic attraction between
two oppositely-charged ions
●
- 2-
Compound ions – Nitrate (NO3 ), Sulfate (SO4 ),
2- +
Carbonate (CO3 ) and NH4
,Covalent Bonding
● COVALENT BONDING: a shared pair of electrons
● Dative covalent bonding = when both electrons come from the same atom
● Shown in diagram w/ an arrow
Shapes of simple molecules and ions
● Valence-Shell Electron-Pair Repulsion theory: shape of molecule/ion determined by the no. of
electron pairs in the outer shell surrounding the central atom
● All electrons = negative charge so each pair repels another and push each other away
● Lone pairs = slightly more electron-dense than bonded pairs so repels more than bonded
●
o
(6 electron pairs w/ no lone pairs = octahedral, 90 )
● Treat double bonds as single bonds
Electronegativity and bond polarity
● ELECTRONEGATIVITY: ability of an atom to attract
the bonding electrons in a covalent bond
● In covalent bond both nucleuses are attracted to the shared pair of electrons; few compounds
are 100% covalent/ionic apart from diatomic gases (e.g. H2, O2 etc.)
● In molecule of H2, both atoms are identical = equal share of electrons in bond = 100%
covalent bond = electrons evenly distributed in bond = non-polar bond (charges evenly
distributed) another example is Cl2
, ● If bonding atoms are dif, one of them will be (most likely) more attracted to the bonding
electrons = more electronegative e.g. in HCl, Cl more electronegative than H, Cl has greater
attraction for bonding pair = bonding pair closer to Cl than H
● This results in a slight charge difference across the H – Cl bond
● Charge difference = permanent dipole = polar covalent bond (all the same thing)
● Polar molecules = overall dipole (non-symmetrical)
● Some molecules are symmetrical and dipoles of the molecule bonds may cancel
each other out e.g. CCl4, in each C – Cl bond it is polar however it’s symmetrical
and all the dipoles act in dif directions and cancel each other out so CCl4 is non-
polar w/ polar bonds
● Electronegativity increases across the group (F is the most electronegative) and up a period
● Reactive non-metallic elements (O, F and Cl) form compounds w/ the most electronegative
atoms
● Reactive metals (Na and K) form compounds w/ the least electronegative atoms
● The greater the difference in electronegativity = the greater the ionic properties
● The greater the similarity in electronegativity = the greater the covalent properties
Intermolecular Forces
● Ionic and covalent bonds IN molecules, intermolecular forces are BETWEEN molecules
● Intermolecular forces = much weaker than ionic/covalent, caused by weak attractive forces
between v small dipoles in dif molecules
● Permanent dipole-dipole interactions:
-polar molecules have permanent dipoles, the permanent dipole of one molecule attracts the
permanent dipole of another causing weak permanent dipole-dipole force
● Van der Waals’ forces (induced dipole-dipole interaction):
-exist between all molecules
-between v small and temporary dipoles in neighbouring molecules
-caused by the movement of electrons in the shells which cause a temporary difference in
charge (a temporary dipole)
-this induces a dipole in neighbouring molecule which then does so to another molecule etc. --
-these small induced dipoles attract one another causing weak intermolecular forces between
them = VDW forces
-greater the no. of electrons = larger the induced dipoles = greater the attractive forces
between molecules = stronger the VDW force
-so the larger the molecule = larger no. of electrons = stronger VDW force
-greater S.A = stronger VDW because have a bigger exposed electron cloud
-stronger VDW forces = higher boiling point as more energy needed to break the bonds
-e.g. as you go down group 4 the VDW forces and boiling points increase due to a) the
atomic/molecular size increase b) no. of shells of electrons increases
● Hydrogen bonding:
-happens with a covalently bonded H to O, N, or F
-usually occurs between O – H or N – H (e.g. water and ammonia)
-O, N and F are v electronegative and H has a high charge density (because it’s so small)
-the extreme difference in charge (v polarised/ permanent dipole-dipole interaction is so
strong) causes a bond to form between the H of one molecule and the O/N/F of another
molecule
-molecules w/ H-bonding are usually organic
● H-bonding strong enough to cause unexpected properties of water:
-Ice less dense than water: ice has an open lattice w/ H-bonds holding the water molecules
apart, when melted the molecules move closer together
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