Graduate Chemist (IChem.C)
PGDip in Food Science and Technology (University of SJP)
12+ years of teaching experience In Cambridge OL & AL
Experienced practical exam supervisor for Cambridge AS level
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Q: Describe the structure of an atom using a fully labelled diagram.
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Q:Identify and describe protons, neutrons and electrons in terms of their relative charges and
relative masses
Particle
Symbol
Relative charge
Relative mass
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Understanding the terms atomic and proton number; mass and nucleon number.
Q: Describe the distribution of mass and charge within an atom.
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Q: Describe the behaviour of beams of protons, neutrons and electrons moving at the same
velocity in an electric field
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2006/ON/02 Q:1
In the 19th and 20th centuries, scientists established the atomic theory and showed that three
sub-atomic particles, electron, neutron and proton, exist. The masses and charges of these
three particles were subsequently determined.
When separate beams of electrons, neutrons or protons are passed through an electric field in
the apparatus below, they behave differently.
(a) (i) Which of these three particles will be deflected the
most by the electric field? .........................................
(b) (ii) In which direction will this particle be deflected?
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(c) (iii) Explain your answer.
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(b) (i) Define the term proton number.
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(ii) Why is the proton number of an atom of an element usually different from the nucleon
number of an atom of the element?
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(c) Protons and neutrons have been used in nuclear reactions which result in the formation of
artificial elements. In such processes, protons or neutrons are accelerated to high speeds and
then fired like ‘bullets’ at the nucleus of an atom of an element. Suggest why neutrons are more
effective than protons as ‘nuclear bullets’.
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(d) In some cases, when neutrons are fired at atoms of an element, the neutrons become part
of the nucleus of those atoms. What effect does the presence of an extra neutron have on the
chemical properties of the new atoms formed? Explain your answer.
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Q: Determine the numbers of protons, neutrons and electrons present in both atoms and ions
given atomic or proton number, mass or nucleon number and charge
Particle
Atomic
number
Mass
number
Number of
protons
Number of
neutrons
Number of
electrons
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Q: Define the term isotope in terms of numbers of protons and neutrons.
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Element and isotopes Number of protons
Number of neutrons
Number of electrons
Hydrogen
Chlorine
Oxygen
Magnesium
Bromine
Carbon
Q: State that and explain why isotopes of the same element have the same chemical
properties.
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Q: State that and explain why isotopes of the same element have different physical
properties.
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Q: State and explain qualitatively the variations in atomic radius and ionic radius across a
period and down a group.
Types of radii
1. Metallic radius
2. Covalent radius
3. Vander waal’s radius
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Atomic radii show predictable patterns across the Periodic Table
o They generally decrease across each Period
o They generally increase down each Group
These trends can be explained by the electron shell theory
o Atomic radii decrease as you move across a Period as the atomic number
increases (increased positive nuclear charge) but at the same time extra
electrons are added to the same principal quantum shell
o The larger the nuclear charge, the greater the pull of the nuclei on the electrons
which results in smaller atoms
o Atomic radii increase moving down a Group as there is an increased number of
shells going down the Group
o The electrons in the inner shells repel the electrons in the outermost
shells, shielding them from the positive nuclear charge
o This weakens the pull of the nuclei on the electrons resulting in larger atoms
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The diagram shows that the atomic radius increases sharply between the noble gas at
the end of each period and the alkali metal at the beginning of the next period
This is because the alkali metals at the beginning of the next period have one extra
principal quantum shell
o This increases shielding of the outermost electrons and therefore increases the
atomic radius
Ionic radius
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The ionic radius of an element is a measure of the size of an ion
Ionic radii show predictable patterns
o Ionic radii increase with increasing negative charge
o Ionic radii decrease with increasing positive charge
These trends can also be explained by the electron shell theory
o Ions with negative charges are formed by atoms accepting extra electrons while
the nuclear charge remains the same
o The outermost electrons are further away from the positively charged nucleus
and are therefore held only weakly to the nucleus which increases the ionic
radius
o The greater the negative charge, the larger the ionic radius
o Positively charged ions are formed by atoms losing electrons
o The nuclear charge remains the same but there are now fewer electrons which
undergo a greater electrostatic force of attraction to the nucleus which
decreases the ionic radius
o The greater the positive charger, the smaller the ionic radius
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Electrons, energy levels and atomic orbitals
Shells
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The arrangement of electrons in an atom is called the electronic configuration
Electrons are arranged around the nucleus in principal energy levels or principal
quantum shells
Principal quantum numbers (n) are used to number the energy levels or quantum shells
o The lower the principal quantum number, the closer the shell is to the nucleus
o The higher the principal quantum number, the higher the energy of the shell
Each principal quantum number has a fixed number of electrons it can hold
o n = 1 : up to 2 electrons
o n = 2 : up to 8 electrons
o n = 3 : up to 18 electrons
o n = 4 : up to 32 electrons
Principal quantum shells
Sub-shells
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The principal quantum shells are split into sub-shells which are given the
letters s, p and d
Elements with more than 57 electrons also have an f shell
The energy of the electrons in the sub-shells increases in the order s < p < d
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The order of sub-shells appear to overlap for the higher principal quantum shells as seen
in the diagram below:
Orbitals
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Sub-shells contain one or more atomic orbitals
Orbitals exist at specific energy levels and electrons can only be found at these specific
levels, not in between them
o Each atomic orbital can be occupied by a maximum of two electrons
This means that the number of orbitals in each sub-shell is as follows:
o s : one orbital (1 x 2 = total of 2 electrons)
o p : three orbitals ( 3 x 2 = total of 6 electrons)
o d : five orbitals (5 x 2 = total of 10 electrons)
o f : seven orbitals (7 x 2 = total of 14 electrons)
The orbitals have specific 3-D shapes
Shapes of the electron orbitals
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Filling the shells and orbitals
The most stable electronic configuration of an atom is the one that has the lowest amount of
energy. The order in which the sub-shells are filled depends on their relative energy.
The sub-shell with the lowest energy, the 1s, is therefore filled first, followed by those that are
successively higher in energy
The Aufbau principle describes the pattern of electron filling in atomic orbitals, stating
that electrons first occupy the lowest energy levels before filling higher ones. This "building-up"
process follows a specific order determined by increasing energy levels of the orbitals.
Representing electronic configurations
Here is a detailed way of writing the electronic configuration of an atom of hydrogen, that
includes information about the number of electrons in each sub-shell.
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Element
Atomic
number
Electron configuration
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Orbitals and the Periodic Table
The arrangement of elements in the Periodic Table reflects the electronic structure of the
elements. The Periodic Table can be split into blocks of elements.
 Elements in Groups 1 and 2 have outer electrons in an s sub-shell. These are therefore
together called the s-block.
 Elements in Groups 13 to 18 (apart from He) have outer electrons in a p sub-shell. These
are therefore together called the p-block.
 Elements that add electrons to the d sub-shells are called the d-block elements. Most of
these are transition elements.
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Filling the orbitals - Electron Box Notation
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A useful way of representing electronic configurations is a diagram that places electrons
in boxes.
Each box represents an atomic orbital.
The boxes (orbitals) can be arranged in order of increasing energy from bottom to top.
An electron is represented by an arrow.
The direction of the arrow represents the ‘spin’ of the electron. (We imagine an electron
rotating around its own axis either in a clockwise or anticlockwise direction.) When there
are two electrons in an orbital, the ‘spins’ of the electrons are opposite, so the two
arrows in this box point in opposite directions.
Electrons in the same region of space repel each other because they have the same
charge. This is called spin-pair repulsion.
Wherever possible, electrons will occupy separate orbitals in the same sub-shell to
minimise this repulsion, so these electrons have their ‘spin’ in the same direction.
Electrons are only paired when there are no more empty orbitals available within a subshell. The spins are then opposite to minimise repulsion.
Examples:
Carbon
Nitrogen
Oxygen
Fluorine
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Sodium
Magnesium
Aluminium
Argon
Free radicals
 A free radical is a species with one or more unpaired electrons. An example of a free
radical is an isolated chlorine atom, which has the electronic configuration
1s22s22p63s23p5.
 In the 3p orbitals, two of the orbitals have paired electrons and the remaining orbital has
an unpaired electron. The unpaired electron in a free radical is shown as a dot ·
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Electronic configuration of ions
 Positive ions are formed when electrons are removed from atoms.
Examples:
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 Negative ions are formed when atoms gain electrons.
Examples:
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 Note that, in general, electrons in the outer sub-shell are removed when metal ions form
their positive ions. However, the d-block elements behave slightly differently. Reading
across the Periodic Table from potassium to zinc, the 4s sub-shell fills before the 3d subshell. But when atoms of a d-block element lose electrons to form ions, the 4s electrons
are lost first.
Examples :
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