Summary Detailed notes for Unit 1.1 Atomic Structure - A level Chemistry
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Unit 3.1.1 - Atomic structure
Institution
AQA
Book
A-Level Chemistry
This document provides detailed notes explaining all principles of the Atomic Structure Unit in AQA A-level chemistry that may appear in the exam. It is made with the specification in mind and attempts to go further than more simple explanations found in other books.
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Unit 3.1.1 - Atomic structure
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Chemistry A level
Topic 1 – Atomic Structure
Contents
- Subatomic particles
- Ions and isotopes
- Periodic table
- Atomic structure model
- Relative mass
- Mass spectrometry and using mass spectra.
- Electronic structure
- Ionisation energies.
The basics
Subatomic particles. There are three subatomic particles: protons, neutrons, and electrons.
Protons and neutrons are found in the nucleus and electrons are found in electron shells/
energy levels. Protons have a relative charge of +1 and a relative mass of 1, neutrons have a
relative charge of 0 and a relative mass of 1, and electrons have a relative mass of 1/2000
and a charge of -1. Since neutrons and protons are far heavier than electrons, most of the
atom’s weight is concentrated at the nucleus. However, the diameter of the nucleus is very
small compared to the diameter of the atom. This means that the electrons and the electron
shells make up the main volume of the atom. The atom itself is overall neutral/ has no
charge meaning that there is an equal number of the charged particles – protons and
electrons in the atom.
Ions are charged particles. They have lost or gained electrons meaning that they do not have
the same number of protons and electrons. Ions with more protons than electrons have a
positive charge and are known as cations. Ions with more electrons than protons have a
negative charge and are known as anions.
Isotopes are different versions of elements. One element will have many isotopes. Isotopes
of an element have the same number of protons but a different number of neutrons. This is
because the number of protons determines the element. As a result, all isotopes have the
same atomic (proton) number but a different mass number.
The chemical property of an element is determined by the number and the arrangement of
electrons. This means that all isotopes have the same chemical properties as they all have
the same electron number and configuration. However, the physical properties such as
density and rate of diffusion are determined by the mass number. Therefore, since isotopes
have different numbers of neutrons and thus different mass numbers, they may have
different physical properties.
Periodic table
The periodic table is a table used to organise and categorise elements by atomic number.
All elements are found in a box. The element will be represented by one of two letters.
There will be two numbers present in the box. The smaller number is the atomic number
, also known as the proton number whilst the larger number is the mass number – the
number of protons and neutrons found in the nucleus.
You can determine the number of protons, neutrons, and electrons in the element using the
mass number and the atomic number. The atomic number is the proton number as well as
the electron number – in neutral atoms. The mass number minus the atomic number is the
neutron number.
The periodic table is split into groups (vertical columns) and periods (rows.) This groups
shows how many electrons are in the outer shell of the element and the period shows the
number of electron shells/ energy levels.
Group 1 = 1 electron in the outer shell
Group 2 = 2 electrons in the outer shell
Group 3 = 3 electrons in the outer shell
Group 4 = 4 electrons in the outer shell
Group 5 = 5 electrons in the outer shell
Group 6 = 6 electrons in the outer shell
Group 7 = 7 electrons in the outer shell
Group 0/ Noble gases = full outer shell.
Row 1 = 1 electron shell
Row 2 = 2 electron shells
Row 3 = 3 electron shells
Row 4 = 4 electron shells.
Models of Atomic Structure
- John Dalton. At the start of the 19th century, John Dalton described atoms as solid
spheres. He claimed that different spheres made up different elements.
- At the end of the nineteenth century, in 1897, J.J Thomson discovered the ‘electron.’
He questioned Dalton’s solid sphere model and came up with the ‘Plum Pudding
model.’ The plum pudding model was a positively charged mass with negative
electrons embedded throughout it.
- In 1909, Ernest Rutherford and his students Geiger and Marsden conducted the
famous ‘Gold Foil’ experiment. They fired positively charged alpha particles at a thin
sheet of gold foil. It was expected that most of the alpha particles would be slightly
defected by the positive pudding. However, most of the particles passed straight
through the gold foil with only a small number of alpha particles being deflected
backwards. This clearly shows that the ‘plum pudding’ model isn’t accurate.
Therefore, Rutherford developed the nuclear model of the atom. This is a positively
charged nucleus surrounded by a cloud of negative electrons.
- However, the scientists soon realised that the nuclear model is not accurate. The
‘cloud’ of electrons would spiral down into the nucleus causing the atom to collapse.
Neil Bohr proposed a new model of the atom where the electrons exist in shells or
fixed energy levels. When electrons move between shells, electromagnetic radiation
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