Cambridge International AS and A
Level Chemistry (9701)
1
,Contents INDEX Page number
-C1. Atoms, Molecules and Stoichiometry 4
-C2. Atomic Structure 6
-C3. Chemical Bonding 11
-C4. States of Matter 17
-C5. Chemical Energetics 22
-C6. Electrochemistry 26
-C7. Equilibria 28
-C8. Reaction Kinetics 33
-C9. The Periodic Table: Chemical Periodicity 34
-C10. Group 2 41
-C11. Group 17 43
-C12. Nitrogen and Sulfur 47
-C13. An Introduction to Organic Chemistry 49
-C14. Hydrocarbons 55
-C15. Halogen Derivatives 61
-C16. Hydroxy compounds 64
-C17. Carbonyl Compounds 69
-C18. Carboxylic Acids and Derivatives 64
2
,Contents INDEX Page number
-C1. Chemical Energetics (Lattice Energy and Entropy) 73 and 116
-C2. Electrochemistry 81
-C3. Equilibria 93
-C4. Reaction Kinetics 108
-C5. An introduction to the Chemistry of Transition Elements 121
-C6. An introduction to Organic Chemistry, Hydrocarbons 134
and Hydroxy Compounds
-C7. Carboxylic Acids and Derivatives 144
-C8. Nitrogen Compounds 148
-C9. Polymerisation 160
-C10. Analytical Techniques 175
-C11. Organic Synthesis 188
3
, AS
Atoms, Molecules and Stoichiometry
Relative atomic mass: The weighted average mass of the atoms of an
element on a scale on which the carbon 12 isotope has a mass of exactly
12 units.
Relative isotopic mass: The mass of a particular isotope of an element on
a scale on which the carbon 12 isotope has a mass of exactly 12 units.
Relative molecular mass: The weighted average mass of a molecule on a
scale on which the carbon 12 isotope has a mass of exactly 12 units. The
sum of the relative atomic masses.
Relative formula mass: The weighted average mass of one formula unit of
a compound on a scale on which the carbon 12 isotope has a mass of
exactly 12 units.
Mole: Amount of substance that has the same number of particles as
there are atoms in exactly 12g of the carbon-12 isotope.
Avogadro Constant: The number of defined particles in a mole of
substance, its 6.02 × 1023 .
Mass spectra:
Abundance of isotopes can be represented on a mass spectra diagram.
𝑃𝑒𝑎𝑐𝑘 𝐻𝑒𝑖𝑔ℎ𝑡
𝑅𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝐴𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 = × 100
𝑇𝑜𝑡𝑎𝑙 𝐻𝑒𝑖𝑔ℎ𝑡
Only works for mass spectra diagrams.
𝑚𝑎𝑠𝑠 × 𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒
𝐴𝑟 =
100
Empirical formula: The simplest ratio of the different atoms present in a
molecule.
Molecular formula: The actual number of each type of atoms in a
molecule.
4
,Percentage of Mass:
𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 × 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠
𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑜𝑓 𝑀𝑎𝑠𝑠 =
𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑
Molar mass: The mass of a mole of substance in grams.
To calculate the number of moles you use this formula:
𝑚𝑎𝑠𝑠 𝑖𝑛 𝑔𝑟𝑎𝑚𝑠
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑀𝑜𝑙𝑒𝑠 =
𝑀𝑟
Example: How many molecules are there in 17 grams of CaCO3:
So first we calculate the number of moles by,
Ca: Mr is 40, C: Mr is 12, O: Mr is 16 * 3 = 48
Total MR is 100. So, we use the formula, 17/100= 0.17 moles per grams of
CaCO3.
So, then we have using the Avogadro’s constant, 1 mole = 6.02 * 1023
molecules. And, 0.17 moles= x molecules.
Finally, we cross multiply and divide. So,
0.17 × 6.02 × 1023
= 1.02 × 1023
1
Molar gas volume: 1mole of any ideal gas in RTP occupies 24dm3.
RTP means room temperature.
Example, How much does 7 grams of hydrogen occupy:
We use the formula to calculate the number of moles.
So, 7g/ 2= 3.5. Then we multiply it by 24 as 1 mole of any ideal gas at room
temperature occupies 24dm3.
3.5 * 24= 84dm3.
Formula For Concentration:
𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒
𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 =
𝑉𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝑑𝑚3
5
,Example, Calculate the concentration of a solution that contains 27g of
calcium chloride in 300cm2 of solution:
First, we use the formula to calculate the number of moles.
Ca is plus and Cl is minus, so the formula is CaCl.
So, Ca has a Mr of 40 and Cl an Mr of 70. The total Mr is 110.
So, 27g/110= 0.245 moles of calcium chloride.
Then we convert the cm3 to dm3.
1dm3= 1000cm3
So 0.27g/0.3dm3= 0.83M.
Atomic Structure
Structure of an Atom:
The smallest amount of an element is an atom. Atoms are made of sub-
atomic particles. There are 3 main sub-atomic particles which are protons,
neutrons and electrons.
Sub-atomic particle Relative mass Relative charge
Proton 1 +1
Neutron 1 0
Electron About 1 -1
1836
Distribution of Mass and Charge in an Atom:
Most of the mass is concentrated in the centre of an atom, in the nucleus.
The nucleus is made up of protons and neutrons. While the electrons
move around in regions of space called orbitals. The number of protons
and electrons are the same as the number of positive charges must be
equal to the negative charges.
6
,Behaviour of a Beam of Subatomic Particles in Electric Fields:
Protons: positively charged, therefore they are deflected to the negative
pole.
Neutrons: no charge, therefore they aren’t deflected.
Electrons: negatively charged, therefore they are deflected to the positive
pole.
Electrons are lighter than protons, that’s why they are deflected at a
greater angle.
Proton number and nucleon number:
To calculate the number of protons, neutrons and electrons we used the
proton number and nucleon number. Nucleon Number or
mass number:
The number of protons +
A electrons.
Symbol of the element.
X Proton Number or atomic
Z number:
The number of protons which
equals the number of electrons.
Isotopes: Are atoms of the same element with the same number of
protons and electrons but different number of neutrons. The chemical
properties of an element depend on the number of electrons in the outer
electron shell. As isotopes of the same element have the same number of
electrons, they have the same chemical properties.
7
,Electronic Configuration:
Electrons are arranged in energy levels called shells. Each shell is described
by a principal quantum number.
The main quantum numbers are 1,2,3 and 4.
As the quantum number increases the energy of the shell increases. This is
because the lowest energy level, quantum number 1, is the closest to the
nucleus. Therefore, the shells further away from the nucleus have more
energy.
Inside the shell there are subshells: s, p, d and f.
Orbital: region in space where there is a maximum probability of finding
an electron.
Each orbital can hold 2 electrons in opposite directions
When electrons are placed in a set of orbitals of equal energy, they occupy
them singly and then pairing takes place.
Electrons placed in opposite direction, both negatively charged, create a
spin to reduce repulsion.
Completely filled or half-filled are more stable (reduced repulsion).
Shapes of Orbitals:
s – orbital:
Has spherical shape, it increases in size as quantum
number increases.
p – orbital:
Dumbbell shape
8
,Allocating Electrons:
Electrons are allocated in order of increasing energy. This is the order:
Chromium and Copper are exceptions, one of the electrons is promoted.
Ionisation energy: The first ionisation energy of an element, is the energy
needed to remove 1 electron from each atom in one mole of the element
in gaseous state to form one mole of gaseous 1+ ions.
Factors affecting Ionization Energy:
Nuclear Charge: As the atomic number increases, the positive nuclear
charge increases. The bigger the positive charge, the stronger the attraction
between the protons and the remaining electrons. More energy is needed
to overcome this.
Shielding Effect: The ability of inner shell electrons to reduce the effect of
the nuclear charge on outer electron shells. The more shielding, the less
energy needed.
Distance of outer electrons from the nucleus:
The shorter the distance between the electrons and the nucleus the greater
the forces of attraction and the more energy needed to overcome them.
9
, First Ionization Energy Trends in the Periodic Table:
Down a Group Across a Period
DECREASES INCREASES
The distance between the nucleus and The shielding remains reasonably
the outer shell increases. constant.
Shielding effect increases. The effective nuclear charge increases.
The nuclear charge increases; The distance between the nucleus and
however, the other 2 factors outweigh the outer shell remains reasonably
it. constant.
Interpreting Ionisation energy from position of element in periodic table:
Elements in group 1 and 2 have outer electrons in an s-subshell.
Elements which are transition metals have outer electrons in d-subshell.
Elements in group 3 to 18 have outer electrons in p-subshell.
Mass Spectrometer:
A mass spectrometer consists of an ion source, an analyser and a detector.
Ion source: Gaseous atoms are bombarded by electrons from an electron
gun and are ionized. Enough energy is supplied to form 1+ charged ions.
These ions are then accelerated out of the ion source by an electric field.
Analyser: The charged particles are deflected by a magnetic field. They are
deflected more or less depending on their mass to charge ratio (m/z). Ions
of heavier isotopes have a larger m/z ratio so they are deflected more. As
most ions normally have the same charge, +1, the deflection is affected only
by their mass.
Detector: Ions are detected by a detector electrically, by their mass to
charge ratio. The abundance is measured by the number of ions that hit the
detector.
10
Level Chemistry (9701)
1
,Contents INDEX Page number
-C1. Atoms, Molecules and Stoichiometry 4
-C2. Atomic Structure 6
-C3. Chemical Bonding 11
-C4. States of Matter 17
-C5. Chemical Energetics 22
-C6. Electrochemistry 26
-C7. Equilibria 28
-C8. Reaction Kinetics 33
-C9. The Periodic Table: Chemical Periodicity 34
-C10. Group 2 41
-C11. Group 17 43
-C12. Nitrogen and Sulfur 47
-C13. An Introduction to Organic Chemistry 49
-C14. Hydrocarbons 55
-C15. Halogen Derivatives 61
-C16. Hydroxy compounds 64
-C17. Carbonyl Compounds 69
-C18. Carboxylic Acids and Derivatives 64
2
,Contents INDEX Page number
-C1. Chemical Energetics (Lattice Energy and Entropy) 73 and 116
-C2. Electrochemistry 81
-C3. Equilibria 93
-C4. Reaction Kinetics 108
-C5. An introduction to the Chemistry of Transition Elements 121
-C6. An introduction to Organic Chemistry, Hydrocarbons 134
and Hydroxy Compounds
-C7. Carboxylic Acids and Derivatives 144
-C8. Nitrogen Compounds 148
-C9. Polymerisation 160
-C10. Analytical Techniques 175
-C11. Organic Synthesis 188
3
, AS
Atoms, Molecules and Stoichiometry
Relative atomic mass: The weighted average mass of the atoms of an
element on a scale on which the carbon 12 isotope has a mass of exactly
12 units.
Relative isotopic mass: The mass of a particular isotope of an element on
a scale on which the carbon 12 isotope has a mass of exactly 12 units.
Relative molecular mass: The weighted average mass of a molecule on a
scale on which the carbon 12 isotope has a mass of exactly 12 units. The
sum of the relative atomic masses.
Relative formula mass: The weighted average mass of one formula unit of
a compound on a scale on which the carbon 12 isotope has a mass of
exactly 12 units.
Mole: Amount of substance that has the same number of particles as
there are atoms in exactly 12g of the carbon-12 isotope.
Avogadro Constant: The number of defined particles in a mole of
substance, its 6.02 × 1023 .
Mass spectra:
Abundance of isotopes can be represented on a mass spectra diagram.
𝑃𝑒𝑎𝑐𝑘 𝐻𝑒𝑖𝑔ℎ𝑡
𝑅𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝐴𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 = × 100
𝑇𝑜𝑡𝑎𝑙 𝐻𝑒𝑖𝑔ℎ𝑡
Only works for mass spectra diagrams.
𝑚𝑎𝑠𝑠 × 𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒
𝐴𝑟 =
100
Empirical formula: The simplest ratio of the different atoms present in a
molecule.
Molecular formula: The actual number of each type of atoms in a
molecule.
4
,Percentage of Mass:
𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 × 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠
𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑜𝑓 𝑀𝑎𝑠𝑠 =
𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑
Molar mass: The mass of a mole of substance in grams.
To calculate the number of moles you use this formula:
𝑚𝑎𝑠𝑠 𝑖𝑛 𝑔𝑟𝑎𝑚𝑠
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑀𝑜𝑙𝑒𝑠 =
𝑀𝑟
Example: How many molecules are there in 17 grams of CaCO3:
So first we calculate the number of moles by,
Ca: Mr is 40, C: Mr is 12, O: Mr is 16 * 3 = 48
Total MR is 100. So, we use the formula, 17/100= 0.17 moles per grams of
CaCO3.
So, then we have using the Avogadro’s constant, 1 mole = 6.02 * 1023
molecules. And, 0.17 moles= x molecules.
Finally, we cross multiply and divide. So,
0.17 × 6.02 × 1023
= 1.02 × 1023
1
Molar gas volume: 1mole of any ideal gas in RTP occupies 24dm3.
RTP means room temperature.
Example, How much does 7 grams of hydrogen occupy:
We use the formula to calculate the number of moles.
So, 7g/ 2= 3.5. Then we multiply it by 24 as 1 mole of any ideal gas at room
temperature occupies 24dm3.
3.5 * 24= 84dm3.
Formula For Concentration:
𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒
𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 =
𝑉𝑜𝑙𝑢𝑚𝑒 𝑖𝑛 𝑑𝑚3
5
,Example, Calculate the concentration of a solution that contains 27g of
calcium chloride in 300cm2 of solution:
First, we use the formula to calculate the number of moles.
Ca is plus and Cl is minus, so the formula is CaCl.
So, Ca has a Mr of 40 and Cl an Mr of 70. The total Mr is 110.
So, 27g/110= 0.245 moles of calcium chloride.
Then we convert the cm3 to dm3.
1dm3= 1000cm3
So 0.27g/0.3dm3= 0.83M.
Atomic Structure
Structure of an Atom:
The smallest amount of an element is an atom. Atoms are made of sub-
atomic particles. There are 3 main sub-atomic particles which are protons,
neutrons and electrons.
Sub-atomic particle Relative mass Relative charge
Proton 1 +1
Neutron 1 0
Electron About 1 -1
1836
Distribution of Mass and Charge in an Atom:
Most of the mass is concentrated in the centre of an atom, in the nucleus.
The nucleus is made up of protons and neutrons. While the electrons
move around in regions of space called orbitals. The number of protons
and electrons are the same as the number of positive charges must be
equal to the negative charges.
6
,Behaviour of a Beam of Subatomic Particles in Electric Fields:
Protons: positively charged, therefore they are deflected to the negative
pole.
Neutrons: no charge, therefore they aren’t deflected.
Electrons: negatively charged, therefore they are deflected to the positive
pole.
Electrons are lighter than protons, that’s why they are deflected at a
greater angle.
Proton number and nucleon number:
To calculate the number of protons, neutrons and electrons we used the
proton number and nucleon number. Nucleon Number or
mass number:
The number of protons +
A electrons.
Symbol of the element.
X Proton Number or atomic
Z number:
The number of protons which
equals the number of electrons.
Isotopes: Are atoms of the same element with the same number of
protons and electrons but different number of neutrons. The chemical
properties of an element depend on the number of electrons in the outer
electron shell. As isotopes of the same element have the same number of
electrons, they have the same chemical properties.
7
,Electronic Configuration:
Electrons are arranged in energy levels called shells. Each shell is described
by a principal quantum number.
The main quantum numbers are 1,2,3 and 4.
As the quantum number increases the energy of the shell increases. This is
because the lowest energy level, quantum number 1, is the closest to the
nucleus. Therefore, the shells further away from the nucleus have more
energy.
Inside the shell there are subshells: s, p, d and f.
Orbital: region in space where there is a maximum probability of finding
an electron.
Each orbital can hold 2 electrons in opposite directions
When electrons are placed in a set of orbitals of equal energy, they occupy
them singly and then pairing takes place.
Electrons placed in opposite direction, both negatively charged, create a
spin to reduce repulsion.
Completely filled or half-filled are more stable (reduced repulsion).
Shapes of Orbitals:
s – orbital:
Has spherical shape, it increases in size as quantum
number increases.
p – orbital:
Dumbbell shape
8
,Allocating Electrons:
Electrons are allocated in order of increasing energy. This is the order:
Chromium and Copper are exceptions, one of the electrons is promoted.
Ionisation energy: The first ionisation energy of an element, is the energy
needed to remove 1 electron from each atom in one mole of the element
in gaseous state to form one mole of gaseous 1+ ions.
Factors affecting Ionization Energy:
Nuclear Charge: As the atomic number increases, the positive nuclear
charge increases. The bigger the positive charge, the stronger the attraction
between the protons and the remaining electrons. More energy is needed
to overcome this.
Shielding Effect: The ability of inner shell electrons to reduce the effect of
the nuclear charge on outer electron shells. The more shielding, the less
energy needed.
Distance of outer electrons from the nucleus:
The shorter the distance between the electrons and the nucleus the greater
the forces of attraction and the more energy needed to overcome them.
9
, First Ionization Energy Trends in the Periodic Table:
Down a Group Across a Period
DECREASES INCREASES
The distance between the nucleus and The shielding remains reasonably
the outer shell increases. constant.
Shielding effect increases. The effective nuclear charge increases.
The nuclear charge increases; The distance between the nucleus and
however, the other 2 factors outweigh the outer shell remains reasonably
it. constant.
Interpreting Ionisation energy from position of element in periodic table:
Elements in group 1 and 2 have outer electrons in an s-subshell.
Elements which are transition metals have outer electrons in d-subshell.
Elements in group 3 to 18 have outer electrons in p-subshell.
Mass Spectrometer:
A mass spectrometer consists of an ion source, an analyser and a detector.
Ion source: Gaseous atoms are bombarded by electrons from an electron
gun and are ionized. Enough energy is supplied to form 1+ charged ions.
These ions are then accelerated out of the ion source by an electric field.
Analyser: The charged particles are deflected by a magnetic field. They are
deflected more or less depending on their mass to charge ratio (m/z). Ions
of heavier isotopes have a larger m/z ratio so they are deflected more. As
most ions normally have the same charge, +1, the deflection is affected only
by their mass.
Detector: Ions are detected by a detector electrically, by their mass to
charge ratio. The abundance is measured by the number of ions that hit the
detector.
10
