AS Chemistry Unit F321: Atoms, Bonds and
Groups
Module 1: Atoms and Reactions
Subatomic Relative Relative
Topic 1: Atoms Particle Charge Mass
Protons +1 1
Atomic Structure Neutrons / 1
Electrons -1 1/2000
● Majority of mass = nucleus
● 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
invisible particles) -> Dalton (start of 19th 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;
e.g. 76% Cl35, 24% Cl37
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.
● AVAGADRO CONSTANT (NA): no. particles p/ mole (6.02 x 1023 mol-1).
● MOLAR MASS: mass p/ mole of a substance (gmol-1).
● No. of moles = no. particles / 6.02 x 1023
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:
CO2 moles = mass/Mr = 4.4/44 = 0.1 moles
CO2 has 1 mole of carbon therefore there are 0.1 moles of carbon at the start.
H2O moles = 1.8/18 = 0.1 moles
H2O has 2 moles of hydrogen therefore there are 0.2 moles of hydrogen at the start.
2. Ratio:
C:H
1.1 : 0.2
3. Divide by smallest:
0.1/0.1 : 0.2/0.1
1:2
Therefore empirical will be CH2.
● Calculating molecular formula:
e.g Molecule has empirical formula of C4H3O2 and molecular mass of 166g. Work out
molecular formula.
1. Find empirical mass:
(4 x 12) + (3 x 1) + (16 x 2) = 83g
2. Divide molecular by empirical mass:
166/83 = 2 therefore 2 units of empirical in molecule.
3. Multiply empirical by units:
C4H3O2 x 2 = C8H6O4
Chemical Reactions
● Balance equations
Calculation of reacting masses, mole concentrations and volume of gases
● Mass calculation: n = m/Mr
● Gas calculation: n = vol/24 or vol/24,000
● Concentration calculation: n = conc x vol(/1000)
● CONCENTRATION: (of a solution) is the am of solute (mol) is dissolved p/ dm 3 of solution.
● Concentrated > 10dm3 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+ .
● Reactive Metal + Acid --> Salt + Hydrogen
Metal Oxide + Acid --> Salt + Water
, Metal Hydroxide + Acid --> Salt + Water
Metal Carbonate + Acd --> Salt + Water + Carbon Dioxide
Alkali + Acid --> Salt + Water
● 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, MgSO 4 . 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:
Mass of water lost: 3.210 – 1.567 = 1.643g
No. of moles of water lost: n = 1.643/18 = 0.0913 moles
2. Find no of moles of anhydrous salt:
n = 1.567/120 = 0.0131 moles
3. Ratio of water to anhydrous salt and divide by smallest no.:
0.0131 salt : 0.0913 water
0.0131/0.0131 : 0.0913/0.0131
1 mole : approx. 6.97 moles
Therefore X = 7 and formula is MgSO4 . 7H2O
● Titrations – find out exactly how much of an acid needed to neutralise a quantity of alkali.
1. Measure out alkali using pipette and put in a flask with some indicator.
2. Do rough titration first for rough idea of end point (which is the point where alkali is
neutralised and changes colour). Add the acid to the alkali using a burette, giving flask a
regular swirl.
3. Do accurate titration; put acid till 2cm3 to end point and then add acid drop wise carefully.
4. Record amount of acid used to neutralise. Repeat to ensure accuracy.
● 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
1st 1s 2=2
2nd 2s 2p 2+6=8
3rd 3s 3p 3d 2 + 6 + 10 = 18
4th 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
● Compound ions – Nitrate (NO3-), Sulfate (SO42-),
Carbonate (CO32-) 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
● (6 electron pairs w/ no lone pairs = octahedral, 90o)
● Treat double bonds as single bonds
Groups
Module 1: Atoms and Reactions
Subatomic Relative Relative
Topic 1: Atoms Particle Charge Mass
Protons +1 1
Atomic Structure Neutrons / 1
Electrons -1 1/2000
● Majority of mass = nucleus
● 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
invisible particles) -> Dalton (start of 19th 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;
e.g. 76% Cl35, 24% Cl37
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.
● AVAGADRO CONSTANT (NA): no. particles p/ mole (6.02 x 1023 mol-1).
● MOLAR MASS: mass p/ mole of a substance (gmol-1).
● No. of moles = no. particles / 6.02 x 1023
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:
CO2 moles = mass/Mr = 4.4/44 = 0.1 moles
CO2 has 1 mole of carbon therefore there are 0.1 moles of carbon at the start.
H2O moles = 1.8/18 = 0.1 moles
H2O has 2 moles of hydrogen therefore there are 0.2 moles of hydrogen at the start.
2. Ratio:
C:H
1.1 : 0.2
3. Divide by smallest:
0.1/0.1 : 0.2/0.1
1:2
Therefore empirical will be CH2.
● Calculating molecular formula:
e.g Molecule has empirical formula of C4H3O2 and molecular mass of 166g. Work out
molecular formula.
1. Find empirical mass:
(4 x 12) + (3 x 1) + (16 x 2) = 83g
2. Divide molecular by empirical mass:
166/83 = 2 therefore 2 units of empirical in molecule.
3. Multiply empirical by units:
C4H3O2 x 2 = C8H6O4
Chemical Reactions
● Balance equations
Calculation of reacting masses, mole concentrations and volume of gases
● Mass calculation: n = m/Mr
● Gas calculation: n = vol/24 or vol/24,000
● Concentration calculation: n = conc x vol(/1000)
● CONCENTRATION: (of a solution) is the am of solute (mol) is dissolved p/ dm 3 of solution.
● Concentrated > 10dm3 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+ .
● Reactive Metal + Acid --> Salt + Hydrogen
Metal Oxide + Acid --> Salt + Water
, Metal Hydroxide + Acid --> Salt + Water
Metal Carbonate + Acd --> Salt + Water + Carbon Dioxide
Alkali + Acid --> Salt + Water
● 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, MgSO 4 . 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:
Mass of water lost: 3.210 – 1.567 = 1.643g
No. of moles of water lost: n = 1.643/18 = 0.0913 moles
2. Find no of moles of anhydrous salt:
n = 1.567/120 = 0.0131 moles
3. Ratio of water to anhydrous salt and divide by smallest no.:
0.0131 salt : 0.0913 water
0.0131/0.0131 : 0.0913/0.0131
1 mole : approx. 6.97 moles
Therefore X = 7 and formula is MgSO4 . 7H2O
● Titrations – find out exactly how much of an acid needed to neutralise a quantity of alkali.
1. Measure out alkali using pipette and put in a flask with some indicator.
2. Do rough titration first for rough idea of end point (which is the point where alkali is
neutralised and changes colour). Add the acid to the alkali using a burette, giving flask a
regular swirl.
3. Do accurate titration; put acid till 2cm3 to end point and then add acid drop wise carefully.
4. Record amount of acid used to neutralise. Repeat to ensure accuracy.
● 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
1st 1s 2=2
2nd 2s 2p 2+6=8
3rd 3s 3p 3d 2 + 6 + 10 = 18
4th 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
● Compound ions – Nitrate (NO3-), Sulfate (SO42-),
Carbonate (CO32-) 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
● (6 electron pairs w/ no lone pairs = octahedral, 90o)
● Treat double bonds as single bonds
