Summary AQA AS Level Physical Chemistry - Unit 3.1.1 - Atomic Structure - Full Notes
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Course
Unit 3.1.1 - Atomic structure (74041)
Institution
AQA
Detailed notes for AQA AS level chemistry unit 3.1.1 (atomic structure). Covers: history of the atom, relative atomic mass, time of flight (TOF) mass spectrometry and calculations, electronic structure and ionisation energies. Includes diagrams and example calculations for RAM and TOF. Also explain...
History of the Atom
Start of the 19th Century (John Dalton):
• John Dalton described atoms as solid spheres.
• He believed different spheres made up different elements.
1897 (JJ Thompson):
• JJ Thompson’s experiments of charge and mass showed that an
atom must contain smaller, negatively charged particles, i.e.
electrons.
• This showed that atoms weren’t solid and indivisible.
• The new model was known as the plum pudding model.
• The plum pudding model presents the atom as a ball of positive
charge with electrons stuck in it.
1909 (Ernest Rutherford):
• Ernest Rutherford and his students Hans Geiger and Ernest Marsden conducted the gold foil experiment.
• Positively charged alpha particles were fired at an extremely thin sheet of gold.
• Based on the plum pudding model, they expected most
particles to pass straight through the sheet (completely
missing the electrons) with only a few particles getting
slightly deflected, as the positive charge was very spread
out.
Results of the Gold Foil Experiment
• Most particles passed straight through the sheet.
• This means the atom is mostly empty space.
• Some particles were slightly deflected.
• This means the atom has a tiny but strong positively
charged region.
• Rarely, some particles were deflected backwards.
• This meant that most of the atom is empty space with the
positive charge being concentrated together.
The Nuclear Model
• In the nuclear model of the atom, there is a tiny positively charged nucleus at the centre, where most of the
mass is concentrated.
• A ‘cloud’ of negatively charged electrons surrounds this nucleus (meaning most of the atom is empty space).
The Bohr Model:
• Scientists realised that electrons in a ‘cloud’ around the nucleus would
quickly spiral down into the nucleus, causing the atom to collapse.
• Niels Bohr proposed a new model of an atom where electrons exist in
shells or orbits of fixed energy.
• When electrons move between shells, electromagnetic radiation (with
fixed energy or frequency) is either emitted or absorbed.
• The Bohr Model fitted experimental observations of radiation emitted
and absorbed by atoms.
The Current Model:
• Scientists later discovered that not all electrons in a shell have the
same energy.
• The model was refined to include sub-shells.
, Atomic Structure
Sub-Atomic Particle Relative Mass Relative Charge
Proton 1 1+
Neutron 1 0
Electron 1/1836 1-
Isotopes
Isotopes: Atoms with the same number of protons but a different number of neutrons. This means they are atoms of
the same element with the same atomic number but a different mass number.
Relative Atomic Mass
Relative Atomic Mass (Ar): The average mass of an atom of an element (taking into account all of its isotopes)
relative to 1/12 of the mass of a 12C atom.
Ar = Σ (Isotopic Abundance x Isotopic Mass Number)
Σ Isotopic Abundance
Example 1: Calculate the RAM of chromium given the data below:
Ar (Cr) = (4.3% x 50) + (83.8% x 52) + (9.5% x 53) + (2.4% x 54)
= 52.1 (1dp)
Example 2: The relative atomic mass of gallium is 69.72. It consists of two isotopes; 69Ga and 71Ga. Find the
percentage composition by mass of these two isotopes in gallium.
3. 69x + 71y = 6972 4. x + y = 100
– 69x + 69y = 6900 x + 36 = 100
2y = 72 x = 74
y = 36
69 71
Ga = 74% Ga = 36%
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