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INORGANIC
CHEMISTRY Part 2
sub-index: 2.1
The electronic basis of the modern Periodic Table * 2.2
The electronic structure of atoms (including s p d f
subshells/orbitals/notation) * 2.3
Electron configurations of elements (Z = 1
to 56) * 2.4 Electron configuration and the
Periodic Table * 2.5 Electron configuration of
ions and oxidation states * 2.6 Spectroscopy and
the hydrogen spectrum * 2.7 Evidence of quantum
levels from ionisation energies
Advanced
Level Inorganic Chemistry Periodic Table Index *
Part 1
Periodic Table history
* Part 2
Electron configurations, spectroscopy,
hydrogen spectrum,
ionisation energies *
Part 3
Period 1 survey H to He *
Part 4
Period 2 survey Li to Ne * Part
5 Period 3 survey Na to Ar *
Part 6
Period 4 survey K to Kr and important trends
down a group *
Part 7
s-block Groups 1/2 Alkali Metals/Alkaline Earth Metals *
Part 8
p-block Groups 3/13 to 0/18 *
Part 9
Group 7/17 The Halogens *
Part 10
3d block elements & Transition Metal Series
*
Part 11
Group & Series data & periodicity plots * All
11 Parts have
their own sub-indexes near the top of the pages
2.6 Spectroscopy
and the hydrogen spectrum
- Spectroscopy is the study of
how electromagnetic radiation (e.g. light) interacts with matter.
- Studying the spectrum of hydrogen is good
example to start with in studying spectroscopy, which in most cases, is the
interaction of electromagnetic radiation with atoms or molecules at the
quantum level.
- Electromagnetic radiation forms
a wide ranging spectrum from radio - microwave - infrared -
visible light - uv - x-rays - gamma rays.
- Light can be considered as energy
packets called photons which have both the properties of a 'particle'
or a transverse 'wave'.
- The relationship between the speed of
light, wavelength of the radiation and the frequency of the photon is given
by ...
- c =
  ,
=
wavelength (m),
= frequency (Hz = s-1), c = speed of light 3 x 108 ms-1
- The relationship between the energy of the
photon and its wave frequency is given by Planck's Equation
- E
=
h
, E
= energy of a single photon (J), = h = Planck's
Constant (6.63 x 10-34 JHz-1),
= frequency
(Hz)
- E is for one
photon interacting with one atom, so you need to multiply by the Avogadro
Constant (6.02 x 1023 mol-1) to get Jmol-1,
and then divide by 1000 to get kJmol-1
- When atoms absorb energy e.g.
in hot flames, high voltage discharge etc., they can become
excited from their normal stable ground state (n=1 in the case of
hydrogen), up to a higher 'energy level' state.
- When the excited atoms lose energy and
return to the ground state, they emit electromagnetic radiation, usually
in the infrared, visible or ultraviolet regions.
- The emitted light can be split and
analysed into its constituent frequencies, using a prism or
grating in a spectrometer, to produce an atomic
emission spectrum of 'lines' of different colour.
- Its also possible for the reverse process
to happen, so if light is passed through the atoms in their ground state,
absorption of energy occurs at exactly the same frequencies as observed in
the emission spectrum. This shows up as black lines against the coloured
spectrum background and is known as the absorption spectrum.
- Both emission and absorption spectra
can be used to identify elements from
their 'finger print' pattern, and from the intensity of the 'signal' quantitative
measurements can be made.
- Neils Bohr was the first
scientist to successfully explain the spectrum of hydrogen using the theory of 'quantisation
of energy' i.e. quantum theory.
- Atomic spectra are caused by electrons
moving between energy levels (shells or sub-shells) and the
accompanying quanta of energy being emitted or absorbed.
- When atom is 'excited', the electron
'jumps' to a higher electronic quantum level e.g. on absorption of a photon.
- This gives rise to absorption
spectra.
- The atom 'relaxes' back to lower/ground
electronic state and loses energy - emission of photons.
- This gives rise to emission
spectra. (see Fig.1)
- The electron can only exist in certain
definite energy (quantum) levels.
- For each atom a photon of light is
absorbed or emitted, the electron changes from one level to another.
- The energy of the photon is the
difference between the energies of the two quantised levels involved in
the electronic change.
- e.g. E of photon = En=2 - En=1 for
the 1st line in the 1st series of the hydrogen spectrum, (see Fig.2)
- where En=2 and En=1
are the specific energy values of the electron in the 1st and 2nd
principal quantum levels.
- The frequency of the emitted or absorbed light is given by
Planck's Equation: E = hv (details above)
-
Spectra are very complex, even for the
simplest single electron system of the hydrogen atom discussed below.
-
The hydrogen
spectrum consists of several series of sharp spectral lines and
the 1st series is illustrated in Fig.1
-
Within each
series, the lines get closer and closer together and eventually
converge.
-
To understand the
origin of the series and their 'convergent' character you need study
Fig.2 below.
-
- The horizontal lines on the
diagram Fig.2 represent the
increasingly higher electronic energy levels, as you go from the
ground state (closest to the nucleus, shell 1, level 1, principal quantum
number n = 1), to the point where the electron is lost in ionisation
(n = infinity)
- Each series arises from the possible
electronic transitions between a particular level and all the levels
above it.
- e.g. The 1st or Lyman Series is between n = 1
(ground state of H) and n = 2, 3, 4 etc. This is in the ultra-violet.
- The 2nd or Balmer Series
arise from electronic transitions from
n = 2 and n = 3, 4, 5, etc.
- Fig.3
-
- Particular
changes are represented on
electronic energy level diagram Fig.3. For hydrogen, arrow ..
- represents the 4th line in
the 3rd series of the emission spectrum (n=7 to n=3),
- represents the 4th line in
the 2nd series of the absorption spectrum (n=2 to n=6),
- represents the 6th line in
the 1st series of the absorption spectrum (n=1 to n=7), and
- represents the 4th line of
the 1st series of the emission spectrum (n=5 to n=1)
- If the absorbed photon has enough energy, it can
remove the most loosely bound electron in a process called ionisation
...
- The 1st ionisation energy
(or enthalpy) is defined as the energy required to completely
remove the most weakly held electron from 1 mole of the gaseous atoms.
- e.g. for the process: Na(g)
==> Na+(g) + e-
- this is the equation for the first
ionisation energy of sodium atoms
- ionisation is always
endothermic, heat absorbed ΔH = +493 kJ mol-1
- For hydrogen, this energy can be
calculated from the frequency of the light emitted or absorbed at the
conversion point in the first series because it corresponds to the
quantum level change from n =1 to n = infinity or vice versa. (see Fig.1)
- Note that the lines in any series, for
any atom, tend to converge in the increasing frequency direction
because the energy levels converge in quantum level value the further
they are from the influence of the positive nucleus.
- The spectra of multi-electron
systems, from He onwards, are much more complex, but from spectroscopy a great
deal can be learned about their electronic structure, which aids our
understanding of an elements chemical behaviour.
- The emission or absorption spectra
of elements can be used to identify and quantify elements from distant stars
to the analysis of steel samples.
- Every element has its 'fingerprint'
pattern, though usually, a few selected and unique frequencies are used in
practice.
- The astronomer Hubble provided some of
the first evidence of the 'Big Bang' or 'expanding universe' theory by
recognising the spectral pattern of the hydrogen series of lines in stars of
very distant galaxies. However all the frequencies were displaced to lower
values because the immense receding of these distance galaxies causes a
Doppler shift, known as the 'red shift'. In the visible spectrum,
VIBGYOR (left to right decreasing frequency, longer wavelength), you can
imagine the 'intergalactic' electromagnetic waves being 'stretched' producing
a longer wavelength i.e. lower frequency, that is a shift in the 'blue' to
'red' frequency direction. The 'red shift' is observed in every direction from
Earth.
- If the 'Big Bang' reverses, then the
'Big Crunch' would be preceded by observing a 'blue shift' as the waves get
'crunched up' by the Doppler effect.
- Incidentally a good sound Doppler
analogy is the increasing pitch of a car engine as it approaches you (a 'blue
shift') at high speed and the decrease in pitch as it moves away from you (a
'red shift').
- The element helium was identified by
its absorption spectrum in our Sun and also by its emission spectrum, when the
products of alpha particle decay were collected in a tiny glass container and
subjected to spectroscopic study i.e. high voltage discharge to create an
emission spectrum.
- -
2.7 Evidence of quantum levels
from ionisation energies
-
Evidence for
electronic 'shell structure' is obtained from spectroscopy and
ionisation energy measurements
-
Interpretations of graphs of the first and successive ionization
energies of the elements provides evidence for the
existence of the main quantum levels and the energy sub-levels too e.g.
-
(1) Ionisation energies
steadily decrease down a group
-
 e.g. for the process: Na(g)
==> Na+(g) + e- etc.
-
Generally speaking the 1st energy
decreases down a group of the periodic table (see graph on the right
of the 1st ionisation energies of the group 1 and metals).
-
As you go down the group, each
element has an extra shell of occupied electronic energy levels
which shield the outermost and most loosely bound electron.
-
Therefore the most outer electron
is becoming further and further from the nucleus as the atom gets
bigger.
-
The further the electron (in the
s sub-shell) is from the nucleus, the less strongly it is held, so
less energy is required to remove it in the ionisation process.
-
This appears to outweigh the
effect of the increasing nuclear charge (Z) because the volume of
the atom is also expanding, allowing for the space required by the
orbitals of the extra shell of electrons.
-
Although this simple pattern
shows some feature of the group 1/2 atoms is steadily changing, I
wouldn't say it was the greatest evidence for the existence of
principal quantum levels ('shells').
-
(2) The patterns of the 1st
ionisation energies when plotted against atomic number (Z)
-
This graph does provide
substantial evidence the principal quantum levels and directly
relates to the structure of the periodic table which is based on the
chemical properties of the elements.
-
The 1st ionisation energy,
and is the energy required to remove the most loosely bound electron from
one mole of the neutral
gaseous atom (it is always endothermic) e.g.
-
1st
IE of helium, He(g) ==> He+(g) + e- (ΔH
= +2370 kJ mol-1 )
- this is the equation for the first
ionization energy of helium atoms.
- You write a similar equation for ANY
element of the periodic table.
-
The energy
required to remove the 2nd most loosely bound electron is called
the 2nd ionisation energy (first possible with helium),
which is therefore defined as the energy required to remove an
electron from one mole of the monopositive ions e.g. Na+,
but here we are just concerned with the first ionization energy.
-
-
Graph of periodic
ionization data for elements 1 to 38
above.
-
Generally speaking the 1st
ionisation energy increase from left to right across a period of the
periodic table. As you go across the
period from one element to the next, the positive nuclear charge is
increasing by one unit as the atomic number increases by one unit and
the positive charge is acting on electrons in the same principal
quantum level. The effective nuclear charge can be considered
to be equal to the number of outer electrons (this is very
approximate and NOT a rule) and this is increasing from left to
right as no new quantum shell is added
i.e. no extra shielding. Therefore the outer electron is increasingly more strongly held by the
increasing positive charge of the nucleus and so,
increasingly, more energy is needed remove it.
-
BUT if the next 1st electron to
be removed from the next element is from a principal 'shell' of
electrons, it is much less strongly held, hence the minimum value
(as argued above). This occurs at atomic numbers 3, 11, 19, 37, 55
and 87 (alkali metals). The highest values indicate the most stable
electron arrangements and these occur at atomic numbers 2, 10, 18,
36, 54 and 86 (noble gases). These numbers themselves indicate a
numerical pattern.
-
The 1st
ionisation energy (1st IE) pattern shows evidence ...
-
From the broad
periodic patterns of 1st IE, electrons are distributed in fixed
patterns of principal quantum levels,
-
The minimum ionisation
energies correspond to the first element in a period (group
1 alkali metal) and the peaks correspond to the last element
in a period (group 0/18 noble gas).
-
These two sets on
minimums and maximums correspond to 'new' sets of ionisation
energies in a 'new' principal quantum level.
-
BUT, that's not all the
graphs show ...
-
From the 'kinks' evidence of
sub-shells of electronic energy levels even within principal
quantum levels (see section
(3) below
on the two 'unexpected' decreases in ionisation energy in period
3.
-
You can also see
evidence of the d block of elements (3d shell) if you look at the pattern
of first ionisation energies of elements 1-38.
-
 (3)
Evidence from sub-levels - the 'kinks' in the 1st
ionization energy graph
-
In the 1st ionisation energy
graph you 'kinks' or abrupt decreases (e.g. Be to B, N to O,
Mg to Al and P to S) which provides evidence of sub-shells of
principal quantum levels. On the right, period 3 ionisation energies
are shown in more detail and you can clearly see this effect and its
similar for period 4. To fully these two 'drops' in ionization
energy, counter to the period trend, you need to bring in your
hopefully gained electron configuration knowledge!
-
(i) A decrease from Mg
[1s22s22p63s2]
to Al [1s22s22p63s23p1]
Box spin diagram of 3s3p orbitals  
==> 
The anomalously low
value for aluminium is considered to be due to the first time a
3p electron is shielded by the full 3s sub–shell and, more
importantly, the 3p electron is a bit
further away (higher in energy) on average from the nucleus than the 3s electrons
(so less strongly bound), so less energy needed to remove it. The effect
to some extent overrides the effect of increasing proton number i.e. increase
in positive
nuclear charge from Mg to Al. However, after the kink, the
continued increase in nuclear charge ensures the Period 3 trend
for the 1st ionisation energy continues as expected until
sulfur, the 2nd anomaly.
-
(ii) A decrease from P
[1s22s22p63s23p3]
to S [1s22s22p63s23p4]
Box spin diagram of 3s3p orbitals 
==> 
Prior to the 4th 3p
electron, the other three p electrons occupy separate p
sub–orbitals (Hund's Rule of maximum multiplicity) to minimise
repulsion between adjacent orbitals. The
anomalously low values for sulphur is
considered to be due to the effect of the first pairing of
electrons in the 3p orbitals producing a repulsion
effect that to some extent overrides the effect of increasing proton
number (increase in positive nuclear charge), so less
energy needed to remove the 4th p electron. From the
'kink', the Period 3 trend for the 1st ionisation energy
continues as expected from sulfur to argon with increase in
nuclear charge.
-
(4) The consecutive ionisation enthalpies for the same element:
- e.g. for the process: Na(g)
==> Na+(g) + e-
- this is the equation for the first
ionisation energy of sodium
- ionisation is always
endothermic, heat absorbed ΔH = +493 kJ mol-1
-
2nd IE of sodium,
Na+(g) ==> Na2+(g) + e- (ΔH
= +4562 kJ mol-1)
- this is the equation for the 2nd
ionisation energy of sodium and dramatically more endothermic.
- 3rd: Na2+(g) ==> Na3+(g) + e- (ΔH
= +6940 kJ mol-1)
- 4th: Na3+(g) ==> Na4+(g) + e- (ΔH
= +9540 kJ mol-1)
- etc. etc. and you can do the same for ANY
element of the periodic table until you run out of electrons to
be removed!
-
The graphs of ionisation versus
ionisation number (1st, 2nd, 3rd etc.) also provide evidence
'shells' of electrons, by looking for sudden extra large leaps in
the progressively increasing ionisation energy as each electron is
removed.
-
Successive ionisation energies
for a given element will always increase because less electrons are
being increasingly more strongly held nearer the nucleus by the
constant positive nuclear charge Z.
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IONISATION
ENERGY PATTERNS
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