Download Atomic Structure

ATOMIC STRUCTURE
THE FUNDAMENTAL PARTICLES
The atom consists of three principal sub-atomic particles: the ELECTRON, the PROTON and
the NEUTRON. Although other particles are known to exist, only these three particles are
considered in chemistry.
Particle
Mass
Charge
Symbol
Electron
1/1836
-1
e-
Proton
1
+1
p
Neutron
1
0
n
The History of the atom
The ancient Greeks had ideas about atoms thousands of years ago, however, it is only the last
100 years or so that our understanding of the atom has been supported by strong experimental
evidence. We begin this recent part of the story with an English scientist called John Dalton.
Dalton
From his experiments and observations, he suggested that atoms were like tiny, hard balls.
Each chemical element had its own atoms that differed from others in mass. Dalton believed
that atoms were the fundamental building blocks of nature and could not be split. In chemical
reactions, the atoms would rearrange themselves and combine with other atoms in new ways.
Thomson
At the end of the nineteenth century, a scientist called J.J. Thomson discovered the electron.
This is a tiny negatively charged particle that is much, much smaller than any atom. When he
discovered the electron, Thomson was experimenting by applying high voltages to gases at low
pressure. He noticed an interesting effect.
Thomson did experiments on the beams of particles in his tube. They were attracted to a
positive charge, so Thomson correctly concluded that they must be negatively charged
themselves. Other experiments showed that it would take about 2000 electrons to weigh the
same as the lightest atom, hydrogen. He called the tiny, negatively charged particles electrons.
TOPIC 12.1 ATOMIC STRUCTURE 1
But where had these tiny particles come from? Since they were so small, Thomson suggested
that they could only have come from inside atoms. So Dalton's idea of the indestructible atom
had to be revised.
Thomson proposed a different model for the atom. He said that the tiny negatively charged
electrons must be embedded in a cloud of positive charge (after all, atoms themselves carry no
overall charge, so the charges must balance out). Thomson imagined the electrons as the bits
of plum in a plum pudding (rather like currants spread through a Christmas pudding – but with
lots more space in between).
Rutherford
1) The next development came about 10 years later. Two of Ernest Rutherford's students,
Hans Geiger and Ernest Marsden, were doing an experiment at Manchester University with
radiation. They were using the dense, positively charged particles (called alpha particles) as
'bullets' to fire at a very thin piece of gold foil. They expected the particles to barge their way
straight through the gold atoms unimpeded by the diffuse positive charge spread throughout
the atom that Thomson's model described. The majority of the particles indeed passed
through the gold foil with little or no disturbance, but some were deflected backwards through
angles greater than 90o. On increasing the thickness of the foil, more and more particles
were deflected backwards. Rare though these deviations were, they occurred too frequently
to be explicable by the Thomson model.
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
parallel beam of -particles
+
+
+
+
+
+
+
+
+
gold atoms
In 1911, Ernest Rutherford interpreted these results and suggested a new model for the atom.
He said that Thomson's model could not be right. The positive charge must be concentrated in a
tiny volume at the centre of the atom, otherwise the heavy alpha particles fired at the foil could
never be repelled back towards their source. On this model, the electrons orbited around the
dense nucleus (centre of the atom).
TOPIC 12.1 ATOMIC STRUCTURE 2
Bohr
The next important development came in 1914 when Danish physicist Niels Bohr revised the
model again. It had been known for some time that the light given out when atoms were heated
always had specific amounts of energy, but no one had been able to explain this. Bohr
suggested that the electrons must be orbiting the nucleus in certain fixed energy levels (or
shells). The energy must be given out when 'excited' electrons fall from a high energy level to a
low one.
Bohr said that these shells could only hold a certain number of electrons and when the shells
were full the atom would be stable and would not react (it would be inert).
The Bohr model was subsequently modified to include subshells when scientists discovered that
not all electrons in the same shell have the same energy.
Chadwick
The existence of a neutral particle had first been predicted by Rutherford in about 1920, but it
proved to be elusive. The discovery of the neutron is credited to James Chadwick. When
beryllium was bombarded with alpha particles, a high energy “particle” with great penetrating
power was produced. The “particle” was unaffected by electric and magnetic fields and was
therefore neutral.
Chadwick measured the recoil produced when the particle was made to collide with nuclei of
different gases, such as nitrogen and hydrogen. The results showed that a neutral particle
existed and that it had a mass marginally greater than that of the proton.
The model of the atom therefore shows protons and neutrons tightly packed in the nucleus.
Since protons are positively charged, there must be a very strong force holding them together.
This is called the strong nuclear force and is much stronger than the electrostatic force of
attractions that holds the electrons in orbit around the positive nucleus.
TOPIC 12.1 ATOMIC STRUCTURE 3
MASS SPECTROMETRY
Mass spectrometry is a useful technique for determining the relative atomic mass of an element
and the relative molecular mass and structure of a compound.
Basic features
1. The apparatus is kept under vacuum to prevent ions produced from colliding with air
particles.
2. Ionisation: The sample is dissolved in a volatile solvent and a high voltage is applied.
Positively charged droplets are released as a fine spray and the solvent evaporates. This
leaves individual positively charged ions.
3. Acceleration: The ions are accelerated by attraction towards a negatively charged plate
which gives all of the ions the same kinetic energy. Since kinetic energy = ½ mv2, lighter ions
move faster than heavier ones.
4. Ion drift: The ions are passed through a hole in the negatively charged plate forming a
beam. The ions drift down the ‘flight tube’ towards a detector. They are separated based on
their different velocities. The faster (lighter) ions reach the detector first.
5. Detection: The separated streams of ions strike a negatively-charged detector plate,
which is contained in the detector and flight times are recorded. The arriving ions generate a
current which sends a signal to a computer or a chart recorder. A plot of the intensity of the
signal against m/z is drawn out.
TOPIC 12.1 ATOMIC STRUCTURE 4
The time of flight of the ions in a mass spectrometer can be calculated since all particles are
accelerated to the same kinetic energy. If a 25Mg+ ion is accelerated to a KE of 4.52x10-16J and
the drift region is 2.13m we can find the velocity of the particle using KE = 1/2mv 2 (remember
mass is in kg) and then the drift time using v = d/t (velocity in ms -1). You will also need the
Avogadro constant 6.022x1023.
The mass of 1 mole of 25Mg+ ions is 25g = 25 x 10-3 kg
so the mass of a single ion is 25 x 10-3 /6.022 x 1023
KE = 1/2mv2
so rearranging
v =  (2 x 4.52x10-16/4.15 x -26)
since v = d/t
4.15 x 10-26 kg
=
v =  (2 x KE/m)
=
147,591 m/s
t = d/v
t = 1.44 x10 -5 s
t = 2.13 / 147,591
(a little over 1/100,000 of a second)
The mass spectrum consists of a series of sharp peaks; the abundance of an ion is proportional
to its peak height. It is usual to convert the spectrum to a “stick diagram” in which only the major
peaks are drawn, the height of the line representing either absolute intensity or a percentage of
the height of the highest peak.
The mass spectrum of an element (magnesium).
Intensity
100
Ar (Mg) = (63 x 24) + (8.1 x 25) + (9.1 x 26)
(63 + 8.1 + 9.1)
80
= 24.3
63
60
40
20
8.1
5
10
15
m/z
TOPIC 12.1 ATOMIC STRUCTURE 5
20
9.1
24 25 26
30
Applications of mass spectrometry
1. Mass spectrometry can be used for the determination of the isotopic composition of a
sample of an element and hence for the measurement of relative atomic mass.
2. A mass spectrum can be used to determine relative molecular masses of compounds since
the line with the highest m/z value is usually the molecular ion M +. There is always a very
small peak at (M+1) for organic compounds which is due to the 13C isotope, present with a
natural abundance of 1.1%.
3. The structure of a molecule can be determined from the fragmentation pattern.
We will look at (2) and (3) in more detail next year.
TOPIC 12.1 ATOMIC STRUCTURE 6
Electronic structure
Electrons are arranged in ENERGY LEVELS or SHELLS. These are given a number from 1
upwards, with 1 being the energy level closest to the nucleus. As you move away from the
nucleus the energy levels get closer and closer together. The electrons furthest from the
nucleus are the highest in energy because the electrons in these energy levels are overcoming
the attraction of the nucleus more than those electrons in closer energy levels.
The electronic configurations for the first 20 elements were learned at GCSE in simple form,
noting down only the energy levels e.g. 2,8,1 for Na.
Energy levels are further divided into SUB-LEVELS (or SUB-SHELLS). These are labelled s, p,
d or f and these increase in energy moving from s  p  d  f. Each sub-level consists of one
or more ORBITALS, and each orbital can hold a maximum of two electrons.
Energy level
1
Distance from
nucleus
Sublevels
Number of
orbitals
closest to the
nucleus
(lowest in
energy)
s
1
2
3
getting further
away from
nucleus
(higher in
energy)
Number of eper sublevel
2e-
Total no. of e- in
the energy level
2
s
1
2e-
p
3
6e-
s
1
2e-
p
3
6e-
d
5
10e-
s
1
2e-
p
3
6e-
d
5
10e-
f
7
14e-
4
8
18
32
At A-Level, this additional detail is included in any electron configurations we use, since it gives
us a greater understanding of the behaviour of electrons. Complete the table below for the
‘simplified’ and full electron configurations of the first 20 elements. When asked for an electron
configuration from now onwards, give the FULL configuration.
A typical energy level diagram is shown below:
TOPIC 12.1 ATOMIC STRUCTURE 7
Energy
___ ___ ___ (4p)
___ ___ ___ ___ ___ (3d)
Energy Level 4
___ (4s)
Energy Level 3
___ ___ ___ (3p)
___ (3s)
Energy Level 2
___ ___ ___ (2p)
___ (2s)
Energy Level 1
___ (1s)
Each orbital (shown by ___ on the diagram) can hold a maximum of 2 electrons (this is PAULI’S
EXCLUSION PRINCIPAL).
All electrons possess a property known as ‘spin’. When the two electrons are placed in an
orbital they must have opposite spin ( or ).
The LOWEST energy orbitals (the ones closest to the nucleus) are filled first (this is the
AUFBAU PRINCIPAL). Orbitals possessing the SAME ENERGY, such as the three p orbitals,
must each be filled by single electrons first before the electrons begin to pair (this is HUND’S
RULE).
Try the following
N
___ 1s
___ 2s
___ ___ ___ 2p
O
___ 1s
___ 2s
___ ___ ___ 2p
Note that the 3d sub-level is HIGHER in energy than the 4s sub-level so the 4s sub-level is filled
FIRST, before the 3d. This explains the position of the 3d elements on the Periodic Table (they
appear after the 4s elements).
NB Diagrams of the different shapes of the orbitals can be seen below, however, knowledge of
why these shapes are the way they are, is not required for A Level.
TOPIC 12.1 ATOMIC STRUCTURE 8
Writing electron configurations
The sub-levels are always written in the following order:
1s
2s
2p
3s
3p
4s
3d
4p
etc.
The number of electrons held in each sub-level is indicated using superscript numbers:
Li
1s2 2s1
Be
1s2 2s2
F
1s2 2s2 2p5
Note: the number of OUTER SHELL ELECTRONS are all the electrons in the HIGHEST
ENERGY LEVELS. These may be contained in different sub-levels e.g.
Li = 1 outer shell electron Be = 2 outer shell electrons F = 7 outer shell electrons (NOT 5!)
NB the electronic configurations of the Cr and Cu atoms are not as expected and should be
learned.
To write electronic configurations for IONS, write out the configuration for the atom and then add
(for negative ions) or remove (for positive ions) the appropriate number of outer electrons
e.g.
Cl
=
1s22s22p63s23p5
Cl-
=
1s22s22p63s23p6
Na
=
1s22s22p63s1
Na+
=
1s22s22p6
NB when forming ions the d block elements lose their 4s electrons BEFORE they lose any 3d
electrons.
e.g. Fe2+ = 1s22s22p63s23p64s0 3d6
TOPIC 12.1 ATOMIC STRUCTURE 9
(4s0 need not be written)
Electron configurations and the Periodic Table
The Periodic Table is arranged according to electron configuration, and all you need to work out
the electron configuration of an atom is a copy of the Periodic Table.
 Elements in Groups 1 and 2 are collectively known as the S BLOCK ELEMENTS. This is
because their outer electrons are always in an s sub-level.
 Elements in Groups 3 to 8 are collectively known as the P BLOCK ELEMENTS. This is
because their outer electrons are always in a p sub-level.
 The transition metals are also collectively referred to as the D BLOCK ELEMENTS.
The emission spectrum of hydrogen (discussed shortly) provides evidence about the
arrangement of electrons about the nucleus. Further evidence about the arrangement of
electrons comes from a consideration of ionisation energies.
Ionisation Energies
Ionisation energies provide important evidence about the arrangement of electrons in an atom.
The magnitude of the ionisation energy measures the ease with which a particular electron may
be removed from a gaseous atom/ion. This in turn depends on the energy (i.e. the distance from
the nucleus) of the electron; the closer the electron is to the nucleus, the harder it is to remove.
The distance of the electron from the nucleus is determined by:
 the positive charge on the nucleus (the number of protons)
 the number of inner electron energy levels which screen the full effect of the nuclear
charge from the outer energy level
 the sub-level
1. First Ionisation Energies
The molar first ionisation energy is the energy required to remove one mole of
electrons from one mole of gaseous atoms. (unit: kJ.mol-1)
e.g. K(g)
K+(g) + e-
I1 = 418kJ.mol-1
N.B. State symbols must be written in these equations.
If first ionisation energy is plotted against atomic number, it is obvious that a definite pattern
exists. The plot provides evidence for the arrangement of electrons in energy levels:
 in H and He, the electron is being removed from the lowest energy level (n=1)
 from Li to Ne, the electron is being removed from a higher energy level (n=2)
 from Na to Ar, the electron is being removed from a still higher energy level (n=3)
 the irregularities in the plot occurring between, for example, Be and B and N and O
suggest that sub-divisions of the main energy levels exist.
TOPIC 12.1 ATOMIC STRUCTURE 10
First Ionisation energies across Period 2
1st Ionisatino energy (kJ/mol)
2500
2000
1500
1000
500
0
Li
Be
B
C
N
O
F
Ne
Na
With the benefit of current knowledge, the following conclusions about the shape of the curve
can be drawn:
1) The sudden drop in ionisation energy, which occurs between an inert gas and an alkali
metal, arises because the outermost electron in the alkali metal is in a new quantum level,
further from the nucleus, and is therefore more easily removed.
2) As a period is traversed, the value of I1 increases. The successive electrons enter the same
quantum level, but the nuclear charge increases. Thus the atom decreases in size across a
period, and the outermost electron, being closer to the nucleus, is harder to remove.
3) The kink in the plot between Be (1s22s2) and B (1s22s22p1) arises because the 2p electron in
B is in an energy sub-level further from the nucleus than the 2s electron in Be and is,
therefore, more easily removed.
4) There is a degree of stability associated with a half-filled p or d sub-shell. The kink in the plot
between N (1s22s22p3) and O (1s22s22p4) can be attributed to the increase in energy caused
by the pairing and hence mutual repulsion of electrons.
5) As a group is descended, e.g. from Be to Ba, the value of I 1 decreases. Although the nuclear
charge increases in successive members of a group, the increased shielding more than
compensates, and the size of the atom increases. Thus, the outermost electron, being
further from the nucleus, is more easily removed.
TOPIC 12.1 ATOMIC STRUCTURE 11
2. Successive Ionisation Energies
The molar second ionisation energy is the energy required to remove one mole of
electrons from one mole of gaseous unipositive ions. (unit: kJ.mol-1)
e.g. K+(g)
K 2+(g) + e-
I2 = 3070 kJ.mol-1
N.B. State symbols must be written in these equations.
Very convincing evidence for the existence of energy levels comes from the plotting of
successive ionisation energies for a particular atom. For example, with potassium, successive
ionisation becomes more difficult because, as each electron is removed and the ion decreases
in size, the influence of the nucleus on the outermost electron increases. However, sharp
increases in ionisation energy are observed when the 2nd,10th and 18th electrons are involved
suggesting that:
 the 2nd electron is in an energy level closer to the nucleus than the 1st.
 the 10th electron is in an energy level closer to the nucleus than the 9th.
 the 18th electron is in an energy level closer to the nucleus than the 17th.
The nucleus of the potassium atom is, therefore, surrounded by electrons grouped into a
number of energy levels.
n=1
2 electrons
n=2
8 electrons
n=3
8 electrons
n=4
1 electron
Determination of Ionisation Energies by emission spectroscopy
The energy levels in e.g. the hydrogen atom converge and therefore the lines in the different
spectral series converge as frequency increases. A SERIES LIMIT is reached for each series,
beyond which the spectrum is continuous.
Series limit
frequency
When the electron is at a distance from the nucleus corresponding to the series limit, it has been
excited into an energy level so high (n=infinity) that it has effectively escaped from the influence
of the nucleus: i.e. the atom has been ionised. The energy required to promote the electron from
the ground state to n = infinity is the energy required to ionise one atom. The molar first
ionisation energy is obtained by multiplying this energy by the Avogadro constant.
TOPIC 12.1 ATOMIC STRUCTURE 12