<p class="MsoNormal"><span style="font-size: 1rem;">Hello, my beautiful readers. Today, I will
continue this article from where I stopped in my previous article on </span><a href="https://www.steemstem.io/#!/@empressteemah/molecules-in-interst-1561644120" target="_blank">MOLECULES IN INTERSTELLAR SPACE AND A CLOSE LOOK AT ELECTRONS</a><span style="font-size: 1rem;">&nbsp;as promised. So, I’ll be
starting my discussion on the atomic emission spectrum of hydrogen, followed by
the Niels Bohr’s model of the hydrogen atom, etc.</span></p><p class="MsoNormal"></p><div style="text-align: center;"><img src="https://res.cloudinary.com/drrz8xekm/image/upload/v1561844065/t3ssnrplojvyz775v1ai.jpg" data-filename="t3ssnrplojvyz775v1ai" style="font-size: 1rem; width: 325.5px;"><span style="font-size: 12px;"><a href="https://pixabay.com/illustrations/spectrum-psychedelic-green-gradient-553216/" target="_blank">Spectrum. Pixabay</a></span></div><p></p><h2><span lang="">THE ATOMIC EMISSION SPECTRUM OF HYDROGEN<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">When hydrogen is placed in a discharge tube at
low pressure and with a high voltage between the two plates, some of the bonds
in the hydrogen molecules (H<sub>2</sub>) are broken to give separate hydrogen
atoms. When the radiation emitted from the discharge tube is passed through a
spectroscope, a series of characteristic lines shows up in the visible
spectrum.</span></p><p class="MsoNormal" style="text-align: center; "><img src="https://res.cloudinary.com/drrz8xekm/image/upload/v1561844208/fc4izmfiikf49xjdpbuu.png" data-filename="fc4izmfiikf49xjdpbuu" style="width: 325.5px;"><span lang=""><o:p><br></o:p></span></p><p class="MsoNormal" style="text-align: center; "><a href="https://commons.wikimedia.org/wiki/File:Hydrogen_spectrum.svg" target="_blank"><span style="font-size: 12px;">The spectral series of hydrogen, on a logarithmic scale. OrangeDog, CC BY-SA 3.0</span></a><span lang=""><o:p><br></o:p></span></p><h2><span lang="">THE BALMER SERIES<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">You read in my&nbsp;</span><a href="https://www.steemstem.io/#!/@empressteemah/molecules-in-interst-1561644120" target="_blank">MOLECULES IN INTERSTELLAR SPACE AND A CLOSE LOOK AT ELECTRONS</a><span lang="">&nbsp;about
Bunsen and Kirchhoff’s work in identifying elements by their characteristic
spectral lines. These spectra are called line emission spectra for the
following reason. When an element’s atoms are given energy, they absorb it and
then emit radiation as discrete (separate) lines, at specific frequencies (and
hence energies) for that element. For example, hydrogen, which has just one
electron, has several prominent lines in the visible part of its emission
spectrum. These are known as the Balmer series, after a Swiss music teacher who
worked out a mathematical relationship between the lines.<o:p></o:p></span></p><h2><span lang="">OTHER SERIES<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">Other series of lines for hydrogen are found
in different parts of the electromagnetic spectrum. The Lyman series is found
in the ultraviolet section and the Paschen series in the infrared. Both these
series are named after their discoverers.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Like the photoelectric effect, the different
series of lines (characteristic for an element) was another 19th-century
mystery that could not be explained by the theories of that time. An
explanation had to wait for the ideas of Planck and Einstein to be developed.<o:p></o:p></span></p><h2><span lang="">ABSORPTION SPECTRA<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">Electrons absorb energy at specific
frequencies, producing a series of black lines on a coloured background. This is
called an absorption spectrum. The background is the visible electromagnetic
spectrum, while the black lines are frequencies that are missing because
electrons have absorbed these energies. The black lines occur in the same place
as the coloured lines in the line emission spectrum, showing that electrons
absorb radiation at specific frequencies and then release the energy at the
same frequency.</span></p><p class="MsoNormal" style="text-align: center; "><img src="https://res.cloudinary.com/drrz8xekm/image/upload/v1561844364/a6o8v3lw49neg0bclpof.png" data-filename="a6o8v3lw49neg0bclpof" style="width: 325.5px;"><span lang=""><o:p><br></o:p></span></p><p class="MsoNormal"><a href="https://commons.wikimedia.org/wiki/File:LymanSeries.svg" target="_blank">Lyman series of hydrogen atom spectral lines in the ultraviolet.  Adriferr , CC BY-SA 3.0</a><span lang=""><o:p><br></o:p></span></p><h2><span lang="">USING THE LIGHT OF ELECTRON TRANSITIONS<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">Neon is used everywhere in advertising signs.
An electric current is passed through the gas at low pressure. The fast-moving
electrons of the electric current excite electrons in the neon atoms into
higher energy levels and, when they return to lower energy levels, orange-red
light is emitted. The colour can be varied by adding other atoms such as argon
or mercury, or by colouring the glass tube the neon is in.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Street lights usually contain sodium or
mercury and they work on the same principle as neon lights. When excited
mercury atoms return to their ground state, the radiation they emit has
frequencies in the ultraviolet, yellow, green and blue parts of the spectrum.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Sodium lights have replaced mercury ones
across the UK road network because the radiation emitted by excited electrons
in sodium atoms when they return to their lower levels is centred upon yellow.
This has longer wavelengths than the light from mercury and is not as readily
scattered by fog, so it can illuminate further. Also, sodium atoms require less
energy to excite their electrons, and sodium is not as toxic as mercury.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Fluorescent lights and low-energy light bulbs
in the home or office contain low-pressure mercury vapour. The inside of the
lighting tube is coated with a phosphor, which absorbs the energy of
ultraviolet light when its electrons are excited. On returning to the ground
state, these produce many frequencies of light in the visible range, which
combine to give white light. A substance fluoresces when it takes in light of
one wavelength, usually outside the visible spectrum, and gives out light of
another, often in the visible spectrum.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">The advantage of fluorescent lights and low
energy light bulbs over filament lamps is that they use less energy, since
nearly all the energy is radiated in the visible spectrum. They feel cool to
touch, while ordinary light bulbs with tungsten filaments waste energy as they
become very hot. However, they do contain up to 50 mg of mercury, which is an
environmental hazard when they are thrown away.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">The Lyman series of lines is found in the
ultraviolet part of the spectrum. It is caused by the excited hydrogen electron
returning from higher levels to the n = 1 ground-state energy level. As this is
the lowest level nearest to the nucleus, far more energy is released when an
electron excited to a particular energy level returns n = 1 than when it
returns to n = 2, and so the lines show up in the more energetic ultraviolet
part of the spectrum.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Each line of hydrogen’s emission spectrum
represents one electron transition: a movement from a higher to a lower level.
As large numbers of hydrogen atoms are involved, all the possible transitions
are represented, which gives the full spectrum of lines.<o:p></o:p></span></p><h2><span lang="">NIELS BOHR’S MODEL OF THE HYDROGEN ATOM<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">In 1911, we were left with the Rutherford
model of the atom. However, there are problems with this model. Rutherford
proposed that electrons orbited the nucleus, rather like planets round the Sun.
Planetary motion was well understood by this time – the Sun’s gravity pulls
planets towards it, while their acceleration, caused by being in a circular
orbit, creates a balancing force. (For an object to travel in a circle, it is
constantly changing direction, so it has to accelerate constantly.)<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Negative electrons in circular motion are
attracted to the positive protons in the nucleus by electrostatic forces. If
they orbited the nucleus like planets, their acceleration would keep them from
falling into the nucleus. But electrons are charged particles, and accelerating
charged particles were known to emit electromagnetic radiation and lose energy.
If Rutherford’s model was correct, instead of a few separate lines, a
continuous spectrum should have been observed, with the atom emitting light all
the time and the electron losing its energy and falling into the nucleus,
causing the hydrogen atom to collapse. Clearly, this does not happen!<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Another model of the atom was needed to
explain the observations. Niels Bohr used the quantum ideas developed by Planck
and Einstein to propose his model for the hydrogen atom. Bohr’s model still had
the hydrogen electron orbiting the nucleus. But the orbits, or energy levels
that the hydrogen could occupy were quantized – that is, they had fixed energy
values. With this model, hydrogen’s line emission spectrum could be explained.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">In Bohr’s model, the electron normally
occupies the lowest possible energy level, called the ground state. This is the
energy level of the electron when it is not excited. Raising the electron to an
excited state (giving it energy), say by an electric discharge, causes it to
move up to a higher level by absorbing a quantum of energy. When it returns to
the lower, ground state energy level, it releases this quantum of energy as a
photon of light of a specific frequency, so giving a line in the emission
spectrum.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Imagine that the hydrogen electron is like a
ball on a staircase, the ball can rest on any step, but it cannot stop in
between. This is the case with the electron. The ball needs energy to go up the
step, and when it falls back down it releases this energy. The lines in an
emission spectrum are closer together at one end because the higher energy
levels that the electron can occupy are also closer together. This happens as
the electron moves away from the nucleus.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">When the electron is closest to the nucleus,
it is at its lowest energy level. Moving the electron away from the nucleus
requires energy, and the further away it is moved, the more energy it requires.
So the more energy an electron receives, the higher it can rise through the
energy levels, and the more energy it will release as it falls back down again.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">Each energy level is given a number, called
the principal quantum number, n. The term ‘principal quantum number’ is still
used to describe the main energy levels of electrons in an atom. When the
hydrogen’s electron is in the n = 1 level, it is not excited, so this is the
ground state.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">The Balmer series of lines is for the energy
transitions when the excited hydrogen electron falls back from higher energy
levels to n = 2. We see the lines because they are in the visible spectrum. The
higher the energy level the electron falls from, the higher the frequency of the
emitted photon. Note that the energy levels eventually merge at n = </span><span lang="">α </span><span lang="">(infinity). This is when the atom has become ionized and lost its
electron, which has escaped from the nucleus’s attraction.<o:p></o:p></span></p><h2><span lang="">IONIZATION ENERGY (IONIZATION ENTHALPY)<o:p></o:p></span></h2><p class="MsoNormal"><span lang="">From the Lyman series we can work out the
energy needed to remove an electron completely from a hydrogen atom. This is
the ionization energy for hydrogen. It is the energy required to take the
electron from the ground state, at n = 1, to where the energy levels converge
at n = </span><span lang="">α</span><span lang="">, when the electron is free of the attraction
of the nucleus. <o:p></o:p></span></p><h2><span lang="">SUMMARY OF THE BOHR MODEL OF THE ATOM<o:p></o:p></span></h2><ul><li><span lang="">Electrons exist only in certain permitted
energy levels and in these levels they do not emit energy.</span></li><li>Electrons move to higher energy levels by
absorbing quanta of energy. They return to lower energy levels by emitting
these quanta as photons of light, which show up as lines in different parts of
the electromagnetic spectrum.</li></ul><p class="MsoNormal"><span lang="">The Bohr model of the atom successfully
explained the lines in the emission spectrum of the hydrogen atom. It worked for
hydrogen, the simplest atom with just one electron, but it did not predict
accurately the spectral lines of atoms with several electrons.<o:p></o:p></span></p><p class="MsoNormal"><span lang="">The Bohr model is important because it used
the idea of quantized energy levels to explain atomic structure and provided a
foundation on which others could build. The Nobel Prize went to Bohr in 1922, one
year after Einstein had received the prize for his explanation of the
photoelectric effect.</span></p><p>







































































</p><h2><span lang="">REFERENCES</span></h2><p><a href="http://www.ece.utep.edu/courses/ee3329/ee3329/Studyguide/ToC/Fundamentals/Bohr/description.html" target="_blank">http://www.ece.utep.edu/courses/ee3329/ee3329/Studyguide/ToC/Fundamentals/Bohr/description.html</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://chem.libretexts.org/Courses/University_of_Missouri/MU%3A__1330H_(Keller)/06._Electronic_Structure_of_Atoms/6.3%3A_Line_Spectra_and_the_Bohr_Model" target="_blank">https://chem.libretexts.org/Courses/University_of_Missouri/MU%3A__1330H_(Keller)/06._Electronic_Structure_of_Atoms/6.3%3A_Line_Spectra_and_the_Bohr_Model</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://www.kentchemistry.com/links/AtomicStructure/waveenergy.htm" target="_blank">https://www.kentchemistry.com/links/AtomicStructure/waveenergy.htm</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://study.com/academy/lesson/electron-transitions-spectral-lines.html" target="_blank">https://study.com/academy/lesson/electron-transitions-spectral-lines.html</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://web.phys.ksu.edu/vqmorig/tutorials/online/hydrogen/emission.html" target="_blank">https://web.phys.ksu.edu/vqmorig/tutorials/online/hydrogen/emission.html</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://brilliant.org/wiki/energy-level-and-transition-of-electrons/" target="_blank">https://brilliant.org/wiki/energy-level-and-transition-of-electrons/</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://en.wikipedia.org/wiki/Emission_spectrum" target="_blank">https://en.wikipedia.org/wiki/Emission_spectrum</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://www.khanacademy.org/science/chemistry/electronic-structure-of-atoms/bohr-model-hydrogen/a/spectroscopy-interaction-of-light-and-matter" target="_blank">https://www.khanacademy.org/science/chemistry/electronic-structure-of-atoms/bohr-model-hydrogen/a/spectroscopy-interaction-of-light-and-matter</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://users.physics.ox.ac.uk/~smithb/website/coursenotes/qi/2016_QI_Lecture_4-6.pdf" target="_blank">https://users.physics.ox.ac.uk/~smithb/website/coursenotes/qi/2016_QI_Lecture_4-6.pdf</a><span lang=""><o:p><br></o:p></span></p><p><a href="https://www.toppr.com/guides/physics/atoms/atomic-spectra/">https://www.toppr.com/guides/physics/atoms/atomic-spectra/</a><a href="https://www.toppr.com/guides/physics/atoms/atomic-spectra/" target="_blank"></a></p><p><a href="https://www.chemguide.co.uk/atoms/properties/hspectrum.html" target="_blank">https://www.chemguide.co.uk/atoms/properties/hspectrum.html</a><br></p><p><a href="https://courses.lumenlearning.com/introchem/chapter/emission-spectrum-of-the-hydrogen-atom/" target="_blank">https://courses.lumenlearning.com/introchem/chapter/emission-spectrum-of-the-hydrogen-atom/</a><br></p><p><a href="http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/bohr.html" target="_blank">http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/bohr.html</a><br></p><p><a href="https://chem.libretexts.org/Courses/Solano_Community_College/Chem_160/Chapter_07%3A_Atomic_Structure_and_Periodicity/7.03_The_Atomic_Spectrum_of_Hydrogen" target="_blank">Atomic structure and periodicity</a><br></p><p><a href="https://en.wikipedia.org/wiki/Hydrogen_spectral_series" target="_blank">https://en.wikipedia.org/wiki/Hydrogen_spectral_series</a><br></p><p><br></p>