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Atomic Structure

AQA: AS Level Chemistry Unit 1

Kate Woodhead

on 1 January 2014

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Transcript of Atomic Structure

The Relative Atomic Mass is the
mass of one atom of an element compared with
one twelfth
of the mass of an atom of

The Relative Molecular Mass of a molecule is the sum of all the
relative atomic masses
of all the atoms in the molecule relative to
one twelfth
of the mass of
second ionisation
energy can be defined as the enthalpy change for:

third ionisation
energy can be defined as the enthalpy change for:

There are two anomalies:

Note that both of these elements have only 1 electron in the 4s orbital.
Atomic Structure @ AS Level
Fundamental Particles
Appreciate that there are various models to illustrate atomic structure.
Describe properties of protons, neutrons and electrons in terms of relative charge and relative mass.
Understand the importance of protons, neutrons and electrons in the structure of the atom.
Electron Configuration
Know that early models of atomic structure predicted that atoms and ions with noble gas electron-arrangements should be stable.
Know the electron configurations of atoms and ions up to Z = 36 in terms of levels and sub-levels (orbitals) s, p and d.
Recall the meaning of mass number (A) and atomic (proton) number (Z).
Explain the existence of isotopes.
Calculate the relative atomic mass from isotopic abundance.
Mass Spectrometry
Understand the principles of a simple mass spectrometer, limited to ionisation, acceleration, deflection and detection.
Know that the mass spectrometer gives accurate information about relative isotopic mass and also about relative abundance of isotopes.
Interpret simple mass spectra of elements.
Know that mass spectrometry can be used to identify elements.
Know that mass spectrometry can be used to determine relative molecular mass.
Fundamental Particles:
various models; protons, neutrons, electrons + properties; role within the atom
Electron Configuration:
shells, sub-shells and orbitals; s-block, p-block and d-block elements; d-block anomalies and ions
Ionisation Energies:
definitions; variation across period 3 (general and 3 features); variation down a group; successive I.E.s; evidence for sub-shells
definitions of A, Z, isotopes, RAM, RMM; properties; calculating RAM
Mass Spec:
principles; interpret spectra
Ionisation Energies
Know the meaning of the term ionisation energy.
Understand how ionisation energies in Period 3 (Na – Ar) and in Group 2 (Be – Ba) provide evidence for electron arrangement in sub levels and in levels.

Born in 460BC, Democritus was the first to suggest the idea of "atoms".
If matter were repeatedly divided in two, there would be a point where particles could no longer be divided.
These "atoms" were the smallest possible particles, and everything in the world was made of them arranged in different patterns.
Democritus - the Father of Atomic Theory
Dalton was born in 1766.
He suggested that atoms are small spheres that cannot be broken apart, and join together to make everything in the world around us.
He also proposed that one atom of oxygen joined to one atom of hydrogen to make a new compound - water!
John Dalton - the Atom Reaches the Modern World
J. J. Thompson (b. 1856) worked in the Cavendish Laboratory in Cambridge where he worked with cathode ray tubes.
He suggested that an atom was a sphere containing small, positively charged particles, and smaller, negatively charged particles, called electrons, arranged within the atom like fruit in a Christmas pudding.
J. J. Thompson - the Plum Pudding Model
New Zealander Ernest Rutherford (b. 1871) suggested that most of the atom was empty space.
He fired alpha particles at a thin piece of gold foil. Most went straight through, but a few came back.
Most of the atom's mass is contained in a tiny, central nucleus, with the electrons orbiting it.
Ernest Rutherford - the "Gold Foil" Experiment...
Neils Bohr, born in Copenhagen in 1885, took Thompson and Rutherford's ideas a little further by suggesting that the electrons moving around the nucleus fitted into shells, rather than moving about randomly.
Neils Bohr - the Model is Further Refined
James Chadwick discovered the neutron in 1920 when firing alpha particles at beryllium atoms.
He discovered that this new particle had a mass of one relative unit, but no charge.
His discovery explained the existence of isotopes.
James Chadwick - the Story is Complete (for now)
Sub-Atomic Particles
The atom consists of:
a small, massive, positively charged nucleus, containing protons and neutrons.
a cloud of electrons orbiting the nucleus.
the neutrons act as "nuclear glue", preventing the positive protons from repelling each other.
The Behaviour of Sub-Atomic Particles in an Electric Field
Electron Arrangement
Electrons are arranged in
energy levels
around the nucleus (1, 2, 3, 4...)
Each energy level is split into
(s, p, d, f...)
Each sub-level is split into a specific number of
(The number of orbitals depends on which sub-level it is).
Each orbital can hold a maximum of two electrons.
Electron Energy Levels and Sub-Levels
Electron Sub-Levels: the lowest level orbitals are filled first, in the order 1s 2s 2p 3s 3p 4s 3d 4p
Quick Review from IGCSE...
Draw the electron configuration diagrams for:
Quick Review from IGCSE...
Describe the structure of the atom in as much detail as you can, without using the following words:
Each orbital can hold a maximum of two electrons.
Electrons in an orbital have a property called "
Two electrons in the same orbital must be "
spin paired
", which means that they spin in opposite directions.
This reduces
between the electrons.
Spin Pairing
Orbitals of the same energy are filled singly before pairing takes place.
Filling p and d Orbitals
Writing Electron Configurations...
Electron configurations are written with the number of electrons in each sub-level indicated by a superscript number.
Just to Make Life Difficult...
The sub-levels are filled in the following order:
1s 2s 2p 3s 3p
4s 3d
The electron configuration is written in the following order:
1s 2s 2p 3s 3p
3d 4s
d-Block Anomalies
Electron Configurations of Ions
When elements form ions, electrons are lost from the highest occupied orbital.
When d-block elements form ions, the
electrons are lost BEFORE the
Classifying Elements
Elements with outer electrons in s orbitals are called
(i.e. elements in groups 1 and 2)
Elements with outer electrons in p orbitals are called
(i.e. elements in groups 3-8)
Elements with outer electrons in d orbitals are called
(i.e. the central block of elements)

The First Ionisation Energy
The First Ionisation Energy is the MINIMUM energy required to remove
one mole
of electrons from
one mole

It is an endothermic process.
The variation of the 1st I.E.s of the first 20 elements shows some interesting trends...
Variation in 1st I.E. Across Period 3
General trend: 1st I.E. increases across a period
Electrons are added to the same energy level, so there is no extra shielding.
Atomic radius decreases.
The force of attraction between the nucleus and the outer electrons increases.
The outer electrons are held more strongly.
The electrons are harder to lose, so more energy is required to remove the outermost electron.

Therefore, the first I.E. increases across a period.
Why is there a drop between Mg and Al?
Mg: outer electron is in
Al: outer electron is in
3p sub-shell sits at a higher energy than 3s sub-shell, and is
further from the nucleus
attractive force is less, so it is easier to remove the electron
Why is there a drop between P and S?
There are three
P: each of the
three 3p
electrons occupy separate orbitals
S: has
four 3p
electrons, so two electrons must
pair up
between the paired electrons make it easier for an electron to be removed
Why is there a drop between Ne and Na?
Na: the outer electron occupies the
sub-shell sits at a higher energy than the
sub-shell, and is further from the nucleus
It is
by the complete 2nd energy level
Despite the increase in nuclear charge, the attraction for the outer electron is less, so it is lost more easily
Successive Ionisation Energies
There is a large jump after the removal of the
2nd electron
The 3rd electron is removed from the 2p sub-shell, which is lower in energy than the 3s (and closer to the nucleus).
There is no longer a full 2p sub-shell
the outer electrons;
The attraction for the outer electron is much stronger, so it is much harder to remove.
Successive Ionisation Energies for Magnesium
Variation in 1st I.E.s Down Group 2
Down a group, there is a
in First Ionisation Energy.

nuclear charge is increasing
whole shells of electrons are being added
outer electrons are further from the nucleus
the outermost electrons are shielded by the complete inner shells of electrons
the attraction for the outer electrons is less
so they are easier to remove.
From the graph shown above, deduce which group Phosphorus belongs to. Explain your answer.
The Mass Number, A, of an atom is the combined number of the protons and neutrons in the nucleus of the atom.
Isotopes are atoms of the same element, which therefore have the same atomic number, but have a different mass number.
The Atomic Number, Z, of an atom is the number of protons in the nucleus of the atom.
Usual notation for an isotope.
Isotopes and Relative Atomic Mass
Many elements exist as more than one isotope.
Chemical properties are the same, as isotopes have the same
electron configuration
Physical properties are slightly different due to slightly
different masses
Relative Atomic Mass
The relative atomic mass (RAM) depends on the
relative abundance
of each isotope.
The RAM is the
weighted average mass
of each of the isotopes of an element, taking into account the relative abundancies of each isotope.
Mass Spectrometry
In order to work out the RAM of an element, the masses and relative abundancies of the individual isotopes must be known. These values are provided by a mass spectrometer.
1. Ionisation
sample, under low pressure, is introduced to the mass spectrometer.
The mass spectrometer must be under a
high vacuum
to make sure there are no other particles present.
electron gun
bombards the sample with
high energy electrons
These electrons
knock out an electron
from the gaseous particles, forming positive ions.
Usually, only one electron is knocked out, however it is possible for two or more to be lost.
2. Acceleration
The positively charged particles are accelerated by an
electric field
Negatively charged plates attract the positive ions to them.
The ions pass through slits in the plates.
The slits also focus the ions into a narrow beam.
3. Deflection
The beam of fast-moving, positive ions is deflected by a
magnetic field
The degree of deflection depends on the mass and charge of the ions.
the mass:charge (
) ratio, the
the ions are
Small, highly charged are deflected most.
Large, low charged ions are deflected least.
4. Detection
Ions with the same m/z ratio will be deflected by the same degree.
Ions will only reach the detector if the magnetic field is the correct strength.
When the ions reach the detector, a tiny
current is produced
The size of current for each stream of ions is dependent on the number of ions that strike the detector.
Therefore the
size of the current
is dependent on the
relative abundance
of those ions.
The deflection of the ions can be changed by altering either the magnetic field or the electric field.
From the RSC...
Interpreting Data from a Mass Spectrometer
Mass spectrum of chlorine ions
Mass spectrum of chlorine molecules
Relative Atomic Mass
and Relative Molecular Mass
Mass Spec. in Space
The word "atom" comes from the Greek word meaning "cannot be cut".
Over to You...
Write out the electronic configurations of:

Do you notice any patterns?
The orbital shapes are derived from quantum mechanics calculations.
These calculations are explored at degree level in both Physics and Chemistry (and related subjects).
The shapes of the s, p, and d orbitals.
How Much Do You Remember....
Without looking at your notes, write out the electronic configurations of:
Chromium (Z = 24)
Copper (Z = 29)
Iron (Z = 26)
Iron (II)
Iron (III)
Quick Review from IGCSE...
Write down everything you remember about isotopes in 5 bullet points.

Now condense these 5 bullet points into 5 words.

Now condense these 5 words into 1 word that epitomises the topic.
Ionisation energies provide evidence for the existence of energy shells and sub-shells.
How Much Do You Remember...
Without looking at your notes, write down three of the definitions introduced so far.

(Hint: if you are struggling to remember the whole definition, write down some of the key words...)
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