AS chemistry notes
Atomic structure
1. know the structure of an atom in terms of electrons, protons and neutrons
- same number of electrons and protons: the atomic number (normally the whole number)
- protons + neutrons: mass number
2. know the relative mass and relative charge of protons, neutrons and electrons
Mass Charge
Proton 1 +1
Neutron 1 0
electron 1/1837 -1
3. know what is meant by the terms ‘atomic (proton) number’ and ‘mass number’
atomic number: number of protons
mass number: number of protons and neutrons
4. be able to determine the number of each type of sub-atomic particle in an atom, molecule
or ion from the atomic (proton) number and mass number
number of protons and electrons: atomic number (or factor in if its an ion)
number of neutrons: mass number - atomic number
5. understand the term ‘isotopes’
“atoms of the same element with the same atomic number, but a different mass number, due to the
presence of a different number of neutrons”
6. be able to define the terms ‘relative isotopic mass’ and ‘relative atomic mass’, based on
the 12C scale
relative isotopic mass: “the mass of an isotope relative to 1/12 of the mass of C-12”
relative atomic mass: “the average of the masses of an atom of an element relative to 1/12 of the
mass of C-12”
7. understand the terms ‘relative molecular mass’ and ‘relative formula mass’, including
calculating these values from relative atomic masses
relative molecular mass: on simple molecules: total mass numbers- note: 2Mg don’t factor in
the “2”, but Cl -> do factor in the 2
,AS chemistry notes
relative formula mass: same, but for giant structures
8. be able to analyse and interpret data from mass spectrometry to calculate relative atomic
mass from relative abundance of isotopes and vice versa
if given the mass spectra:
9. be able to predict the mass spectra, including relative peak heights, for diatomic
molecules, including chlorine
10. understand how mass spectrometry can be used to determine the relative molecular mass
of a molecule
- the peak with the largest m/z= RFM
11. be able to define the terms ‘first ionisation energy’ and ‘successive ionisation energies’
first I.E: “amount of energy required to remove 1 mole of electrons from one mole of gaseous
elements”
successive I.E: “amount of energy to remove a second electrons from 1 mole of gaseous 1+ ions”
,AS chemistry notes
12. understand how ionisation energies are influenced by the number of protons, the electron
shielding and the electron sub-shell from which the electron is removed
Number of protons Higher effective nuclear charge
Electron shielding Repulsion from electrons on inner shells= less
attractive forces from the protons
Electron sub-shell Some shielding from P and S shells
Note: a p4 electron is easier to remove than a p3 electron because there is respulsion between the
pair
13. understand reasons for the general increase in first ionisation energy across a period
“increase in effective nuclear charge without a large increase in shielding”
14. understand reasons for the decrease in first ionisation energy down a group”
“distance from the nucleus increases, as does shielding”
15. understand how ideas about electronic configuration developed from:
a. the fact that atomic emission spectra provide evidence for the existence
of quantum shells
b. the fact that successive ionisation energies provide evidence for the existence of
quantum shells and the group to which the element belongs
c. the fact that the first ionisation energy of successive elements provides evidence
for electron sub-shells
The answers to this spec point are in the point
16. know the number of electrons that can fill the first four quantum shells
s= 2
p= 6
d= 10
17. know that an orbital is a region within an atom that can hold up to two electrons with
opposite spins
orbital: “ a region with a high density of electrons that can hold up to 2 electrons with opposite spins”
18. know the shape of an s-orbital and a p-orbital
20. know that electrons fill subshells singly, before pairing up, and that two electrons in the
same orbital must have opposite spins
s= 1 orbital
p= 3 orbitals
, AS chemistry notes
d= 5 orbitals
each orbital must have one electron before gaining another
21. know that elements can be classified as s, p and d-block elements
22. understand that electronic configuration determines the chemical properties of an
element
- Chemical reactions involve addition of removal of electrons
23. understand periodicity in terms of a repeating pattern across different periods
periodicity: “when a particular chemical and physical trend is repeated across different periods
25. understand reasons for the trends in the following properties of the elements
from periods 2 and 3 of the Periodic Table:
- the melting and boiling temperatures of the elements, based on given data, in terms of
structure and bonding
- ionisation energy based on given data or recall of the plots of ionisation energy versus
atomic number
Along a period Down a group
Melting and boiling point Increases: the atomic radii Decreases: atomic radii
decreases, which means a increases, because increase in
greater force of attraction shells- smaller forces of
between atoms attraction
Ionisation energy Increases: atomic radii Decreases: increase in
decreases, increase in effective shielding and distance from
nuclear charge w/o sig nucleus
increase in shielding
Topic 2: bonding and structure
2A: Bonding
1. know that ionic bonding is the strong electrostatic attraction between oppositely charged
ions
self-explaniatory
Atomic structure
1. know the structure of an atom in terms of electrons, protons and neutrons
- same number of electrons and protons: the atomic number (normally the whole number)
- protons + neutrons: mass number
2. know the relative mass and relative charge of protons, neutrons and electrons
Mass Charge
Proton 1 +1
Neutron 1 0
electron 1/1837 -1
3. know what is meant by the terms ‘atomic (proton) number’ and ‘mass number’
atomic number: number of protons
mass number: number of protons and neutrons
4. be able to determine the number of each type of sub-atomic particle in an atom, molecule
or ion from the atomic (proton) number and mass number
number of protons and electrons: atomic number (or factor in if its an ion)
number of neutrons: mass number - atomic number
5. understand the term ‘isotopes’
“atoms of the same element with the same atomic number, but a different mass number, due to the
presence of a different number of neutrons”
6. be able to define the terms ‘relative isotopic mass’ and ‘relative atomic mass’, based on
the 12C scale
relative isotopic mass: “the mass of an isotope relative to 1/12 of the mass of C-12”
relative atomic mass: “the average of the masses of an atom of an element relative to 1/12 of the
mass of C-12”
7. understand the terms ‘relative molecular mass’ and ‘relative formula mass’, including
calculating these values from relative atomic masses
relative molecular mass: on simple molecules: total mass numbers- note: 2Mg don’t factor in
the “2”, but Cl -> do factor in the 2
,AS chemistry notes
relative formula mass: same, but for giant structures
8. be able to analyse and interpret data from mass spectrometry to calculate relative atomic
mass from relative abundance of isotopes and vice versa
if given the mass spectra:
9. be able to predict the mass spectra, including relative peak heights, for diatomic
molecules, including chlorine
10. understand how mass spectrometry can be used to determine the relative molecular mass
of a molecule
- the peak with the largest m/z= RFM
11. be able to define the terms ‘first ionisation energy’ and ‘successive ionisation energies’
first I.E: “amount of energy required to remove 1 mole of electrons from one mole of gaseous
elements”
successive I.E: “amount of energy to remove a second electrons from 1 mole of gaseous 1+ ions”
,AS chemistry notes
12. understand how ionisation energies are influenced by the number of protons, the electron
shielding and the electron sub-shell from which the electron is removed
Number of protons Higher effective nuclear charge
Electron shielding Repulsion from electrons on inner shells= less
attractive forces from the protons
Electron sub-shell Some shielding from P and S shells
Note: a p4 electron is easier to remove than a p3 electron because there is respulsion between the
pair
13. understand reasons for the general increase in first ionisation energy across a period
“increase in effective nuclear charge without a large increase in shielding”
14. understand reasons for the decrease in first ionisation energy down a group”
“distance from the nucleus increases, as does shielding”
15. understand how ideas about electronic configuration developed from:
a. the fact that atomic emission spectra provide evidence for the existence
of quantum shells
b. the fact that successive ionisation energies provide evidence for the existence of
quantum shells and the group to which the element belongs
c. the fact that the first ionisation energy of successive elements provides evidence
for electron sub-shells
The answers to this spec point are in the point
16. know the number of electrons that can fill the first four quantum shells
s= 2
p= 6
d= 10
17. know that an orbital is a region within an atom that can hold up to two electrons with
opposite spins
orbital: “ a region with a high density of electrons that can hold up to 2 electrons with opposite spins”
18. know the shape of an s-orbital and a p-orbital
20. know that electrons fill subshells singly, before pairing up, and that two electrons in the
same orbital must have opposite spins
s= 1 orbital
p= 3 orbitals
, AS chemistry notes
d= 5 orbitals
each orbital must have one electron before gaining another
21. know that elements can be classified as s, p and d-block elements
22. understand that electronic configuration determines the chemical properties of an
element
- Chemical reactions involve addition of removal of electrons
23. understand periodicity in terms of a repeating pattern across different periods
periodicity: “when a particular chemical and physical trend is repeated across different periods
25. understand reasons for the trends in the following properties of the elements
from periods 2 and 3 of the Periodic Table:
- the melting and boiling temperatures of the elements, based on given data, in terms of
structure and bonding
- ionisation energy based on given data or recall of the plots of ionisation energy versus
atomic number
Along a period Down a group
Melting and boiling point Increases: the atomic radii Decreases: atomic radii
decreases, which means a increases, because increase in
greater force of attraction shells- smaller forces of
between atoms attraction
Ionisation energy Increases: atomic radii Decreases: increase in
decreases, increase in effective shielding and distance from
nuclear charge w/o sig nucleus
increase in shielding
Topic 2: bonding and structure
2A: Bonding
1. know that ionic bonding is the strong electrostatic attraction between oppositely charged
ions
self-explaniatory
