CHEMISTRY Part 2
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
Level Inorganic Chemistry Periodic Table Index *
Periodic Table history
* Part 2
Electron configurations, spectroscopy,
ionisation energies *
Period 1 survey H to He *
Period 2 survey Li to Ne * Part
5 Period 3 survey Na to Ar *
Period 4 survey K to Kr and important trends
down a group *
s-block Groups 1/2 Alkali Metals/Alkaline Earth Metals *
p-block Groups 3/13 to 0/18 *
Group 7/17 The Halogens *
3d block elements & Transition Metal Series
Group & Series data & periodicity plots * All
11 Parts have
their own sub-indexes near the top of the pages
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
- 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
- c = ,
= 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
= energy of a single photon (J), = h = Planck's
Constant (6.63 x 10-34 JHz-1),
- 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
- 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.
spectrum consists of several series of sharp spectral lines and
the 1st series is illustrated in Fig.1
series, the lines get closer and closer together and eventually
To understand the
origin of the series and their 'convergent' character you need study
- 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
- 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.
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
- 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
- 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
- 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
- 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
2.7 Evidence of quantum levels
from ionisation energies
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keywords: Na(g) ==> Na+(g)
+ e- * Na+(g) ==> Na2+(g) + e- * He(g) ==> He+(g) + e- *
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