CHEMISTRY
MODULE 2 – FOUNDATIONS IN CHEMISTRY
2.1 – ATOMS AND REACTIONS
Atomic structure and isotopes
● Describe the structure of an atom including the distribution of mass and
charge.
Protons Electrons Neutrons
Mass 1 1/2000th 1
Charge +1 -1 0
o Mass is in nucleus(protons and neutons) with electron cloud
surrounding it.
● Describe the difference between mass (nucleon) number and atomic
(proton) number
o Mass number – the total number of protons and neutrons in the
nucleus of an isotope
o Atomic number – the total number of protons in the nucleus of an
atom
● Explain what an isotope is.
o Atoms of an element (same number of protons) with different
numbers of neutrons and different masses.
● Define the term ‘first ionisation energy’
o The amount of energy needed to remove an electron from each
atom in 1 mole of gaseous atoms of an element.
Relative masses
● State what is used as the standard measurement for relative masses
o Carbon-12
● Define ‘relative atomic mass’ (Ar)
o The weighted average mass of an atom, relative to 1/12th the mass
of a carbon-12 atom.
THIS MEANS:
(The weighted average atomic mass of an element, relative to
carbon 12, on a scale where carbon 12 is 12.)
● Define ‘relative isotopic mass’
o The mass of an isotope compared to 1/12th the mass of a
carbon-12 atom.
THIS MEANS
(The mass of an isotope of an element relative to carbon 12 on a
scale where carbon 12 is 12.)
, ● How do you calculate the relative atomic mass of an element?
o Multipy each isotopic mass by the percentage for that isotope
o Add each of these valuve together
o Divide by 100.
● Define ‘relative molecular mass’ (Mr)
o The mass of a molecule of a compound relative to 1/12th the mass
of a carbon-12 atom.
o This should only be used for simple covalent substances
● Define ‘relative formula mass’ (Mr)
o The mass of one formula unit of a compound relative to 1/12th the
mass of a carbon-12 atom.
o This is better and can be used for ionic or covalent compounds
Compounds, formulae and equations
● Predict ionic charge from groups in the periodic table:
Group 1 2 3 4 5 6 7
Charge 1+ 2+ 3+ 3- 2- 1-
● Give the formulae of the following compound ions:
Ion Nitrate Carbonate Sulphate Ammonium
Formula NO3- CO32- SO42- NH4+
Ion Hydroxide Silver Zinc
Formula OH- Ag+ Zn2+
● Describe the procedure you should use for balancing equations
1. Write out the atoms involved in each compound
2. Check the the compound has the correct formula (balance
charges for ionic ones. Check using memory or dot and cross
diagrams for covalent ones)
3. Use big numbers to balance out the number of atoms/ions on each
side of the equation.
Amount of substance
MODULE 2 – FOUNDATIONS IN CHEMISTRY
2.1 – ATOMS AND REACTIONS
Atomic structure and isotopes
● Describe the structure of an atom including the distribution of mass and
charge.
Protons Electrons Neutrons
Mass 1 1/2000th 1
Charge +1 -1 0
o Mass is in nucleus(protons and neutons) with electron cloud
surrounding it.
● Describe the difference between mass (nucleon) number and atomic
(proton) number
o Mass number – the total number of protons and neutrons in the
nucleus of an isotope
o Atomic number – the total number of protons in the nucleus of an
atom
● Explain what an isotope is.
o Atoms of an element (same number of protons) with different
numbers of neutrons and different masses.
● Define the term ‘first ionisation energy’
o The amount of energy needed to remove an electron from each
atom in 1 mole of gaseous atoms of an element.
Relative masses
● State what is used as the standard measurement for relative masses
o Carbon-12
● Define ‘relative atomic mass’ (Ar)
o The weighted average mass of an atom, relative to 1/12th the mass
of a carbon-12 atom.
THIS MEANS:
(The weighted average atomic mass of an element, relative to
carbon 12, on a scale where carbon 12 is 12.)
● Define ‘relative isotopic mass’
o The mass of an isotope compared to 1/12th the mass of a
carbon-12 atom.
THIS MEANS
(The mass of an isotope of an element relative to carbon 12 on a
scale where carbon 12 is 12.)
, ● How do you calculate the relative atomic mass of an element?
o Multipy each isotopic mass by the percentage for that isotope
o Add each of these valuve together
o Divide by 100.
● Define ‘relative molecular mass’ (Mr)
o The mass of a molecule of a compound relative to 1/12th the mass
of a carbon-12 atom.
o This should only be used for simple covalent substances
● Define ‘relative formula mass’ (Mr)
o The mass of one formula unit of a compound relative to 1/12th the
mass of a carbon-12 atom.
o This is better and can be used for ionic or covalent compounds
Compounds, formulae and equations
● Predict ionic charge from groups in the periodic table:
Group 1 2 3 4 5 6 7
Charge 1+ 2+ 3+ 3- 2- 1-
● Give the formulae of the following compound ions:
Ion Nitrate Carbonate Sulphate Ammonium
Formula NO3- CO32- SO42- NH4+
Ion Hydroxide Silver Zinc
Formula OH- Ag+ Zn2+
● Describe the procedure you should use for balancing equations
1. Write out the atoms involved in each compound
2. Check the the compound has the correct formula (balance
charges for ionic ones. Check using memory or dot and cross
diagrams for covalent ones)
3. Use big numbers to balance out the number of atoms/ions on each
side of the equation.
Amount of substance
