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Edexcel AS Chemistry: (Unit 1) Atomic Structure & the periodic table

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J Amuah-Fuster

on 8 July 2016

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Transcript of Edexcel AS Chemistry: (Unit 1) Atomic Structure & the periodic table

All matter is made of atoms
All atoms of the same element are identical, but have different masses to atoms of other elements
A compound is made when atoms of two or more elements combines
Atoms cannot be created nor destroyed
In a chemical reaction, the atoms in the reactants are rearranged to give the products
Learning objectives
The Atom
The Story of the atom - key scientists
Electron Configuration
Electron Affinity
Ionisation Energy
It is important that you understand these two terms as you will have to use them in the exam.
The atom would have
(one or more)
. Therefore,

between the electrons,
reducing the attractive force
between the nucleus and the electrons. This results in the shells moving further from the nucleus.
460 B.C.
A brief history of the atom
J.J. Thompson
You need to be able to present electron configurations in both ways.
More examples
However, there are some rules that determine how the orbitals are filled?
You will also be required to show the electrons in their orbitals (electron-in-box notation). A single arrow represents 1 electron.
An example - 12Mg
With the exception of the first quantum shell, the other shells contain sub-shells of specific energy.
The electrons will first occupy the orbitals of the lowest energy before filling the orbitals of higher energy. [See aufbau principle]
Hence, electrons fill the 4s orbital before the 3d orbitals because the 4s orbital has a lower energy.
Sub-shells and energy levels
The s-orbitals have lower energy and therefore the electrons in them are more stable. Hence, the s-orbitals are filled first. See next slide
This diagram could be very easily drawn in the exam to help you write electron configurations.
Writing electron configurations
There are a number of ways of remembering how to write electron configurations.
Writing electron configurations
There are three p-orbitals, each holding two electrons. Hence, the maximum number of electrons in the p-orbitals is 6.
The 3 p-orbitals exist together.
The p-orbitals
The s-orbitals have the lowest energy of all the orbitals. In common with all orbitals, an s-orbital holds 2 electrons.
The s-orbital increases in size as the quantum shell number increases.
The s-orbital
Meet the orbitals
Electron density maps reveal that electrons orbiting a nucleus can be found in probability spaces called orbitals.
An electron within an orbital behaves like a wave where there are regions where you will not find an electron (regions called nodes).
Atomic orbitals
When hydrogen atoms are heated, the electrons move from their lowest energy sate (ground state) to a higher, excited state.

The excited state is not stable so the electron returns to a lower energy level. In doing so, it releases a energy in the form of light (electromagnetic radiation).

The frequencies of light emitted are specific to the element.
Emission spectra
In this lesson you learn:

about the arrangement of electrons around the nucleus of an atom
that electrons exist in orbitals that have specific energy levels
that the electron structure is central to the chemistry of an element
In this example the
of the peaks are given.

Therefore, the abundance of each isotope is a
fraction of the total height
, i.e. 500 in this example.
Another example
Atoms of an
will simply lose an electron, becoming ionised (positively charged):
Ag + energy Ag+ + e-

Compounds will also lose an electron, but this creates instability and may cause
C2H5OH C2H5OH+ C2H5+ + OH+

Ionisation of the sample
All the air is pumped out so there are no particles inside the machine.
All samples have to be in the
How does it work?
Measures the masses of atoms and fragments of molecules.

Identification of elements and compounds.
The mass spectrometer
In this lesson you learn how:

the mass spectrometer can be used to identify the isotopes of an element

to calculate the RAM

the MS can be used to identify unknown substances (‘CSI’ in action!)
Lesson 3 - The Mass Spectrometer
Moseley measured the positive charge of atoms. This positive charge was called the atomic number.
Proton then discovered and this corresponded to the atomic number.
To account for the differences in masses of atoms of the same element the neutron was suggested.
Problem solved – the discovery of the proton and neutron
Why do you think that this happens?
If you recall, the d-orbitals fill after the s-orbital has filled of the next quantum shell.

However, there are
two exceptions
to this rule as illustrated by Period 4 elements
Exceptions to the aufbau principle
Edexcel expect you to be able to write the electron configurations for the
first 36 elements (H to Kr)
using the s, p, d notation
‘electron-in-box’ notation.
The first 20 elements
Examples for you to try
Write the electron configuration for:
Try these
The quantum shells, n, are number 1, 2, 3, etc.
Within the shells there are sub-shells (orbitals)
Quantum shells and orbitals
There are 5 d-orbitals, each holding two electrons. Hence, the maximum number of electrons in the d-orbitals is 10.
The 5 d-orbitals exist together.
The d-orbitals
Explaining the emission spectra for hydrogen is easy – after all it only has one electron.

What about other elements? The complicated emission spectra produced by other elements suggested that within one quantum shell there were sub-shells. And in these sub-shells the electrons existed in orbitals.
Sounds too simple...
What would the emission spectra look like if the electron could take any energy value, i.e. if the energy of an electron was continuous?
The hydrogen electron is found in the first quantum shell (n=1) but can be excited to higher energy levels. The arrows represent the electron returning back to lower energy levels. The fact that the electron can only move between specific energy levels results in the emission spectra shown previously.
Energy transitions in hydrogen
Emission spectra for an element
Spectrum of light
What is the key difference between the emission spectra and the spectrum of light?
Max Karl Ernst Ludwig Planck
1858 - 1947
Max Planck proposed his Quantum theory in 1900.
He stated that energy existed in tiny packets called ‘quanta’.
Therefore, electrons also have specific levels of energy.
Quantum theory
trace amounts of chemicals can be detected in body fluids.
Uses of the mass spectrometer
Identify the species responsible for the peaks at m/e = 65, 63, 43 and 41
This parent peak is of 2-chloropropane containing the 35Cl isotope.
This parent peak is of 2-chloropropane containing the 37Cl isotope.
Identification of molecules
These species represent the parent molecules.
The mass spectrometer will

the molecules
and break them into the
. The fact that chlorine exists as two isotopes produces the three diatomic molecules detected.
Identification of molecules
Some books may use
. It is the same as
Calculating the RAM of Neon
– the sample to be tested is heated and turned into a gas, this allows particles to be ionised and their masses analysed

– the electron gun produces a beam of electrons which ionises the gas and breaks up molecules into charged fragments.

– an electric accelerates the positively charged ions.

– powerful electromagnets deflect the charged particles. Ions of smaller mass are deflected the most (track D in diagram). The magnetic field is increased gradually to deflect ions of greater mass.

– a detector calculates the m/e ratio (mass/charge) of each positive ion and its relative abundance is compared with the most abundant particle (the parent ion). A graph is generated.
The subatomic particles
Group 1 or 0? Both have a mass of 40. So which Group would you have put them in?
If elements had isotopes then how can atoms with different positive centres still be the same element!!!

Mendeleev had ordered elements according to their masses…..some elements had the same mass, e.g., K and Ar both have a mass of 40. Which Groups do you put them in?
A new problem
What do you think they discovered using the mass spectrometer?
Aston invented the mass spectrometer which accurately measured the mass of atoms.
Invention of the mass spectrometer
The modern atom
SUMMARY: the old and the new
What did the results mean?
“It was as if you fired a gun at a piece of paper and the bullet came back and hit you in the face!”
Alpha particles were fired at a piece of thin gold foil.
A detector was used to find out what happened to the alpha particles.
The gold-foil experiment
Dalton thought that the atom was a solid sphere, rather like a marble. After all the word atom comes from the Greek word ‘atmos’ or ‘uncuttable’
However, much of what Dalton said about chemical reactions paved the way for the huge advances in chemistry that were made during the 19th century.
Dalton’s atomic view




Niels Bohr Physicist, 1885 - 1962
Bohr extended the ideas of quantum theory to electrons.

He explained why light (electromagnetic radiation) is emitted by atoms when they are heated.

The light emitted by an atom is of specific wavelengths and not one continuous spectrum.

Bohr went on to suggest the existence of quantum shells in which electrons exist around the atom. The orbit closest to the nucleus has the lowest energy and the electrons in this orbit can only have the energy associated with this orbit.

You and I know these quantum shells as just shells or energy levels.

At GCSE you learned the 2,8,8 rule for electron configuration. As you will see this was a gross oversimplification!
Niels Bohr and quantum shells
Why is the peak at a m/e=78 also a parent ion?
You can check this by working out the molar mass of the molecule
This is the mass spectrum for 2-chloropropane
This last peak is the parent ion – in other words the ionised whole molecule
Identification of molecules
The Mass Spectrometer
Isotopes have the same atomic number but different mass numbers
The relative isotopic mass is the mass of an atom of that isotope divided by 1/12th the mass of carbon-12 atom.
The RAM of an element is the weighted average mass of an atom of that element divided by 1/12th the mass of a carbon-12 atom.
The relative isotopic mass
You must learn these as stated.
The relative atomic mass
…it was the work of Ernest Rutherford and others that established the model of the atom which we currently accept.
The plum pudding model was an early attempt. The atom is a sphere of evenly spread positive charges and electrons.
Discovered the electron using a cathode ray tube.
Came up with the ‘plum pudding’ model of the atom.
J.J. Thomson
John Dalton was born in 1776.
He used measurement of the masses in which elements combine to propose his hypothesis.
Dalton’s atomic theory
The noble gas chosen is one Period above the Period in which the element you want to write the electron configuration is found, e.g., Na and Mg are in Period 3 and Ne is in Period 2.
noble gas configuration
is often used to abbreviate the electron configuration of an element.
Short method
All identified species must be written as positively charged ions.
m/e = 43: CH3CHCH3+
m/e = 41: CH2CHCH2+
m/e = 65: 37Cl-CHCH3+
m/e = 63: 35Cl-CHCH3+
Identification of molecules - answers
If for example the most abundant species of Ne had a +2 charge then the m/e value would be 10. But the actual mass is still 20 because you would multiply the m/e by 2.
The particles identified by the mass spectrometer are referred to as ‘species’ as they are not atoms or molecules. They can also be called ions.
Interpreting a mass spectrum
Geiger and Marsden worked for Rutherford at the University of Manchester. In 1909, they did the now famous gold foil experiment which revealed the structure of the atom.
Ernest Rutherford (1871 – 1937)
1908, Nobel Prize in Chemistry
Ernest Marsden
(1889 – 1970)
Hans Geiger
(1882 – 1945)
The story of the atom - key scientists
More stable
More stable
Explaining the exceptions
There are two exceptions to this (see later).
Half arrows can also used.
An orbital holds a maximum of 2 electrons with opposite spins. This is shown by arrows in opposite directions.

The aufbau (‘building-up’) principle says that the electrons fill orbitals of lowest energy first.

Hund’s rule states than when more than one orbital is available in the same sub-shell then electrons will try and remain unpaired.

Pauli’s exclusion principle states that two electrons in the same orbital must be opposite spins.
The rules and principles
Dalton’s theory (1803)
e. recall that ideas about electronic structure developed from:
i. an understanding that successive ionization energies provide evidence for the existence of quantum shells and the group to which the element belongs
ii. an understanding that the first ionization energy of successive elements provides evidence for electron sub-shells
f. describe the shapes of electron density plots (or maps) for s and p orbitals
g. predict the electronic structure and configuration of atoms of the elements from hydrogen to krypton inclusive using 1s …notation and electron-in-boxes notation (recall electrons populate orbits singly before pairing up)
h. demonstrate an understanding that electronic structure determines the chemical properties of an element
i. recall that the periodic table is divided into blocks, such as s, p and d
j. represent data, in a graphical form, for elements 1 to 36 and use this to explain the meaning of the term ‘periodic property’
k. explain trends in the following properties of the element from periods 2 and 3 of the periodic table:
i. melting temperature of the elements based on given data using the structure and the bonding between the atoms or molecules of the element
ii. ionization energy based on given data or recall of the shape of the plots of ionization energy versus atomic number using ideas of electronic structure and the way that electron energy levels vary across the period.

Ernest Rutherford
Usually the ionised species has a single positive charge. Therefore, the mass/charge (m/e) ratio is equal to the mass of the species.
Pharmaceutical industry
Isotopic analysis of elements & radioactive dating
Drug testing in sports
Space exploration
Write the short hand electron configurations for:
This is because an atom is more stable if it has half-filled or filled set of 3d-orbitals and a single electron in the 4s-orbital.
The Periodic Table is divided into blocks which corresponds to the highest energy level orbital in which you will find an electron of that element.
Francis William Aston
1877 – 1945
Questions on the Periodic Table - 1
Boiling Temperature
A closer look at the non-metals
Electrical conductivity
Atomic radii, 1st I.E., and electronegativity
Structure and bonding
Electron configuration
Questions on the Periodic Table - 2
Melting temperature
In this lesson you will learn:

Explain, in terms of structure and bonding, trends in the following properties across Period 3.

Melting temperature and boiling temperature
Electrical conductivity
Ionisation energy
The Periodic Table
Can you explain the trends?
Trends in 1st EA
Second electron affinity: The heat energy change when one mole of gaseous single negative anions gains one mole of electrons to form one mole of gaseous double negative anions.

X-(g) + e- X (g)
e.g. O-(g) + e- O (g)

The 2nd EA is endothermic as the electron is being added to an ion which is already negative and there is repulsion between similar charges and energy has to be supplied in order to enable the electron to enter the anion
2nd EA
First electron affinity (EA): The heat energy change when one mole of gaseous atoms gains one mole of electrons to form one mole of gaseous single negative anions.

X(g) + e- X-(g)
e.g. F(g) + e- F-(g)

The 1st EA is an exothermic quantity as the nucleus is partially unshielded by the incompletely filled outer shell and therefore attracts the electron towards the atom.
1st electron affinity (EA)

There is a
in the value for Boron.
This is because the
extra electron
has gone into one of the
2p orbitals
This 2p electron has a
greater energy level
than a 2s-electron, because it is
partially shielded
by the 2s-electrons.
The "blips" explained
There are a few
in the general trend – these will be explained later.
There is a ‘
general increase
’ in first ionisation energy across a period before the value
drops dramatically
for the
start of another period
Trend in 1st I.E of the elements
The first
big jump
in I.E. indicates a
change in shells
and hence tells us the
Group number
which the element belongs to.
Interpreting data on successive ionisation energies
The large jumps in I.E. shows the existence of quantum shells.

The electronic configurations of elements:
- 2 electrons in 1st shell,
- 8 electrons in 2nd shell,
- etc…
What can we learn ...
For you to do
Plot a graph of the successive ionisation energies of K (potassium)
You first have to calculate the log values of I.E. This will allow you to plot the graph more easily.

You will need
: graph paper, pencil and ruler.

: the rules for plotting and presenting graphs.
Successive ionisation energies show the energy required to remove each electron from an atom one by one.
Successive ionisation energies
Recall that electrons can only have specific levels of energy. If an electron is excited (when the atom is heated) then it occupies a higher energy level.
If the electron is
enough then it could
from the atom completely. If this happens then the atom is left
. The
amount of energy
required for this to happen is called the
ionisation energy
Losing your electrons!
Effective nuclear charge
Electrons are attracted to the nucleus. The greater the number of protons in the nucleus, the greater the attraction.

However, electrons further away from the nucleus are attracted less and moreover, the inner electrons repel those on the outside.

Its this interaction that will be the focus of this section.
In this lesson you learn:

to explain the interactions between the electrons and the protons in the nucleus

to apply these explanations when discussing the properties and reactions of elements and compounds .
The forces within an atom
There is little change up the group because there is little change in the effective nuclear charge.
Across a Period, the nuclear charge increases which makes it easier for the nucleus to attract an additional electron.

Furthermore, the distance between the nucleus and the outer shell decreases (atomic radii decreases).
Explaining the trends
Electron affinity is the ability of an atom or ion to attract an electron into its valence (outer shell).
In this lesson you learn:

to define the 1st and 2nd electron affinity.
to explain why the 1st EA is exothermic but the 2nd EA is endothermic.
to use EA to explain the properties and reactivity of elements

There is a
in the value for Oxygen. The extra electron has
paired up
with one of the electrons already in one of the 2p orbitals.
repulsive force
between the
paired-up electrons
means that
less energy is needed
to remove one of them.
The "blips" explained
There is a different reason to explain why there is a drop in I.E. from Be to B and from N to O. Based on what you have learned can you explain the blips?
Explaining the blips
Lithium atoms have 3 protons so you would expect the pull on electrons to be greater. However, the
1st I.E. of lithium is lower than that of helium

of electron
reduces the effective nuclear charge
electron is
further away

from the nucleus =
weaker nuclear attraction
for an electron
2370 kJ/mol
519 kJ/mol
1310 kJ/mol
What affects I.E.?
Mg+ (g) Mg (g) + e-
M+ (g) M (g) + e-
Second ionisation energy (2nd I.E.):
The heat energy needed to

remove one mole of electrons

one mole


gaseous single positive ions


form one mole of
gaseous double positive ions
(or dipositive).
Second Ionisation Energy
In this lesson you learn:

to define the 1st, 2nd and successive ionisation energies of an element.
to interpret data on IE as evidence for the existence of quantum shells/levels.
to use IE to explain the properties and reactivity of elements.
Ionisation Energy
Size of negative ions
There are fewer electrons (and sometimes one less shell of electrons). Hence, electron-electron
repulsion is reduced
effective nuclear charge increases
. This
increases the attractive force
of the nucleus on of the electrons, pulling the electrons closer to the nucleus.
Size of positive ions
Electron shielding and the effective nuclear charge will be terms that you will be using a lot when explaining the properties and reactivity of elements.
Electron shielding is when the inner electrons shield the nucleus from the outer electrons.
The effective nuclear charge is the net charge on a nucleus, after allowing for the shielding effect of the inner electrons.
Electron shielding and the effective nuclear charge
This equation illustrates the information in the text.
The forces that act on electrons are
(because they are
charged particles

magnitude of the force
between two objects depends on the


and the
distance between
Forces that act on electrons in an atom
418 kJ/mol
494 kJ/mol
519 kJ/mol
1st I.E.
down the group
even though nuclear charge increases
. The
outer s
is easier to remove
because of
increased shielding
greater distance from the nucleus
. Therefore, outer electron is
held less strongly
1st I.E. down a Group
Mg (g) Mg (g) + e-
M (g) M (g) + e-
Specific example
General equation
Third ionisation energy (3rd I.E.):
The heat energy needed to

remove one mole of electrons
one mole


gaseous double positive ions


form one mole of

triple positive ions
(or tripositive).
Third Ionisation Energy (kJ/mol)
You will need to know this for the exam.


must show the key features such as the
big jumps

and the

steady increases
Explaining the pattern
High values; electrons in shell closest to nucleus and no shielding.
Large jump = shell change.
Large jump = shell change.
I.E. values increase as effective nuclear charge increases.
Large jump = shell change; more difficult to remove 2nd electron.
Mg (g) Mg (g) + e
M (g) M (g) + e
Learn this word for word!
First ionisation energy (1st I.E.):
The heat energy needed to

remove one mole of electrons

one mole


gaseous atoms


form one mole of
gaseous single positive ions
(or unipositive).
Ionisation energies always have a
value – denoting an

First Ionisation Energy (kJ/mol)
Are the electrons at this end the furthest or closest to the nucleus?
Describe and interpret what the parts labelled on the graph show.
Successive I.E. for K
Specific example
General equation
Specific example
General equation
Learn this word for word!
Learn this word for word!
Chapter 2:
Atomic structure & the periodic table

a. recall the definitions of relative atomic mass, relative isotopic mass and relative molecular mass and understand that they are measured relative to 1/12th the mass of a 12C atom
b. demonstrate an understanding of the basic principles of a mass spectrometer and interpret data from a mass spectrometer to:
i. deduce the isotopic composition of a sample of an element, e.g., polonium
ii. deduce the relative atomic mass of an element
iii. measure the relative molecular mass of a compound
c. describe some uses of mass spectrometers, e.g., in radioactive dating, in space research, in sport to detect use of anabolic steroids, in the pharmaceutical industry to provide an identifier for compounds synthesised for possible identification as drugs
d. recall and understand the definition of ionization energies of gaseous atoms and that they are endothermic processes

Never state the mass of an electron as zero.
Always quote it as negligible.
shells &
(electron beam)
Ionic radius vs. atomic radius
Mass Spectra for a compound
a mass spectrometer can be used on a probe or a space shuttle to analyse elements and compounds found out in space and on planets (probes only so far).
detection of small amounts of chemicals,
distinguishing isomers (molecules with the same molecular formula but different structural formulae) of chemicals,
linked with gas chromatography (GC) and high performance liquid chromatography (HPLC) to identify individual chemicals in a mixture.
the isotopic composition of an element can be used to
determine the source
of the element, e.g., the radioactive polonium used to poison the Russian Alexander Litvinenko was traced to its source
by analysing the isotopic composition
of the polonium found in his body.
proportion of radioactive C-14
in a dead sample of tissue is used to determine
how long ago the tissue died
Successive ionisation energies can show us...
The values
get smaller down a group
as the
electron is being removed from energy levels further from the nucleus
- there is
more shielding
effective nuclear charge
determines the I.E.
Forces that act on electrons in an atom
Positive & Negative
A positive ion is always smaller than its neutral atom.
A negative ion is always larger than its neutral atom.
Full transcript