MOLECULES IN INTERSTELLAR SPACE AND A CLOSE LOOK AT ELECTRONS #3

in steemstem •  last year 

DISCOVERING SODIUM

Almost 50 years before Kirchhoff and Bunsen used emission spectra to identify elements, the German scientist Josef von Fraunhofer was experimenting with light. When he passed sunlight through a very fine slit and a high-quality prism, some wavelengths were missing from the sunlight’s absorption spectrum. They showed up on his spectrum as black lines in the otherwise continuous spectrum. These lines represented energies that were being absorbed by some unknown material.


A demonstration of the 589 nm D2 (left) and 590 nm D1 (right) emission sodium D lines using a wick with salt water in a flame. Cmglee - Own work, CC BY-SA 3.0

Fraunhofer recorded the wavelengths of several hundred black lines, which are now named after him. Two were close together in the yellow part of the spectrum at almost 600 nm. At exactly the same wavelengths, Bunsen and Kirchhoff found two characteristic yellow lines in the emission spectrum of sodium. They concluded that sodium must be in the Sun’s atmosphere. Using this hypothesis, they soon discovered other elements in the Sun.

To explain the Fraunhofer lines more fully, we need to look at the Bohr model of the atom again. When excited, electrons fall from higher energy levels to lower levels: they emit energy to give a line emission spectrum. In reverse, they reach the higher (excited) energy levels by absorbing quanta of energy (photons).

So, in an absorption spectrum, the quanta of energy the electrons absorb in going to these higher energy levels will be represented as black lines for the wavelengths missing from the spectrum.

For Fraunhofer’s spectrum for sunlight, the light has passed from the interior of the Sun through the Sun’s atmosphere. Assuming that there is sodium vapour in the Sun’s atmosphere and that its electrons are being excited to higher energy levels, then the corresponding points in the spectrum for the light reaching Earth will be missing – hence the dark lines for sodium.

DISCOVERING HELIUM

In 1868, the French astronomer Pierre Jules Janssen took his spectroscope to India to view a total eclipse of the Sun. In the spectrum of light from the Sun’s corona he saw a bright line very close to the two sodium lines. This line could not be accounted for by the line spectra of any known element. In the same year the English scientist, Norman Lockyer, also observed this line and suggested that it was caused by a new element. It was given the name helium, after the Greek word helios for the Sun.

At the time, Lockyer was ridiculed for his suggestion that he had identified a new element. Yet this discovery later helped helium to be discovered on Earth. In 1895, a quarter of a century after helium was found on the Sun, it was isolated by the Scottish chemist William Ramsey. He found the gas trapped in a uranium ore, where it had been formed as a product of radioactive decay.

Lasers – amplifying the energy from electron.

The letters in the word ‘laser’ stand for ‘light amplification by stimulated emission of radiation’. Lasers, invented in 1960, have revolutionized medicine, technology and science. The ruby laser was the first, and it works by using a burst of ultraviolet light to excite electrons in atoms of a ruby rod. In the same instant, a few electrons fall back to lower energy levels and produce photons. These photons are bounced back and forth by mirrors at the ends of the rod and stimulate other excited electrons to fall back to lower levels, emitting a simultaneous burst of photons of the same frequency. This laser pulse is very intense, and provided one of the mirrors is partially transparent, it can pass through it. The ruby laser produces a wavelength in the red part of the spectrum. Other lasers use other materials and produce wavelengths from infrared to X-ray.


Diagram of the first ruby laser. US gov, Public Domain

Electrons as waves producing images of smaller and smaller objects.

In 1923, Louis de Broglie proposed that, just as light could behave as a particle or a wave, so could matter. He suggested that electrons, regarded as particles, also had wave-like properties. Four years later, a beam of electrons was diffracted (bent and scattered) by a metal crystal to form a pattern. This diffraction could only be explained by assuming that the electrons were behaving as waves.

The electron microscope uses the wave-like behaviour of electrons. It allows us to see images of very small objects in much greater detail than we can see with ordinary microscopes, which rely on visible light. Light is visible to us between wavelengths of about 400 to 700 nm. To form an image of an object, the object cannot be smaller than half the shortest wavelength, that is, 200 nm (or a 20 thousandth of a millimetre).

Electrons travelling at high speed have very short wavelengths, and they allow images of objects measuring 2 × 10-3 nm to be made and magnified. For this reason, the electron microscope is a very helpful tool in biological and chemical research.

In 1986, the scanning tunnelling microscope (STM) took electron microscopy one stage further. With the aid of a computer to process the data and enhance the image, scientists used the charge on the electrons around atoms to make images of separate atoms and molecules. A newer technique involves probing the surface of a substance with fine conducting tip. The instrument that does this is the scanning probe microscope (SPM). SPMs have enabled the nanotechnology revolution to take off.


The first STM produced commercially, 1986.

Science Museum London / Science and Society, CC BY-SA 2.0

A MODERN MODEL OF HOW ELECTRONS ARE ARRANGED

Once it was realized that electrons could behave like waves, a new model of the atom was possible. This model is based on some very complicated mathematics to describe the wave properties of electrons. Erwin Schrödinger takes the credit for devising an equation that describes the energy levels for electrons in hydrogen and other atoms. His model is known as the quantum mechanical model.

ELECTRON SHELLS

Electron shells correspond to the energy levels that Bohr first identified in his model of the atom. The first shell has a principal quantum number n = 1, the second shell’s principal quantum number is n = 2, and so on. The first shell is closest to the nucleus and has the lowest energy. As the principal quantum number increases, so does its energy.


A Bohr diagram of lithium.

Pumbaa (original work by Greg Robson), CC BY-SA 2.0


Each shell can hold a maximum number of electrons. You will already have met or read about shells in some previous chemistry articles, and you may remember that the arrangement of electrons is known as its electron configuration, also referred to as electronic structure.

EXAMPLE

Question: Magnesium has an atomic (proton) number (Z) of 12. Work out the arrangement of electrons in the shells of a magnesium atom.

Answer: The atomic number tells you there are 12 electrons. (The atomic number is the number of protons in an atom, and as the atom has no net charge, it has the same number of electrons.

The shells of lowest energy are filled first, so:

  • Shell 1 will take 2 electrons, which is all it can hold;
  • Shell 2 will take the next 8 electrons, the maximum it can hold;
  • Shell 3 will take the remaining two.

So the electron configuration of magnesium is 2,8,2. You can use a dot-and-cross diagram to represent this also.

EVIDENCE FOR SHELLS FROM SUCCESSIVE IONIZATION ENERGIES

In my previous article MOLECULES IN INTERSTELLAR SPACE AND A CLOSE LOOK AT ELECTRONS, we met the term ionization energy, also referred to as ionization enthalpy, which is the energy that an electron must be given to remove it from an atom (ion). Now, let’s look at this term more closely. In order to measure ionization energy, the atoms have to be separated, and this means they must be in the gaseous state. For convenience, the value for ionization energies is given for one mole of atoms (or ions): So, the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of ions with a single positive charge is called the first ionization energy.

So the first ionization energy is the energy required to do this:

X(g) → X+(g) + e-                              ΔH = first ionization energy in kJ mol-1

where X is any element, (g) tells you that it is gaseous, and e- is the symbol for an electron.

The second ionization energy is the energy required to remove a second mole of electrons:

X+(g) → X2-(g) + e-

Let’s look again at magnesium with its electron configuration of 2,8,2.

First ionization energy:                  Mg(g) → Mg+(g) + e-                      ΔH = +738 kJ mol-1

Second ionization energy:            Mg+(g) → Mg2-(g)+ e-                      ΔH = +1451 kJ mol-1

Third ionization energy:                Mg2+(g) → Mg3+(g)+ e-                   ΔH = +7733 kJ mol-1

The ionization energy increases as each successive electron is removed. This is the reason: the positive nuclear charge stays the same, because there are still 12 protons in the nucleus, and each time an electron is removed, the remaining ones are attracted more strongly by the nucleus. Notice that there is a large jump in the energy required to remove the third electron. This is because we are breaking into the second shell, which is closer to the nucleus. This shell is also less shielded by inner electron shells from the positive charge of the nucleus, because there is only one full shell between it and the nucleus.

From these values, you can see that there is a large range in the ionization energies of magnesium.

ARRANGING ELECTRONS IN SUBSHELLS

Let us now return to the atomic emission spectra of elements. When looked in finer detail, the lines on these spectra are seen to be divided into more lines. Each finer line represents the energy level of a subshell.


Electron atomic and molecular orbitals. Patricia.fidi - own work, Public Domain

So electrons arranged in shells are also subdivided into subshells, or sub-levels. The fne lines result from electron transitions (movement of electrons) between the subshells. The subshells are known by letters:

  • s             subshell contains             2              electrons
  • p            subshell contains             6              electrons
  • d            subshell contains             10           electrons
  • f             subshell contains             14           electrons

The letters were given in the early 20th century and refer to spectral lines. Some of the lines were sharp, hence s, some were more spread out or difiuse, and some of the lines were brighter and called principal lines. The subshell (sub-level) takes the number of the principal quantum number or shell. Taking, for instance, the 3d subshell, it has a higher energy than the 4s subshell. This has important consequences for the chemistry of the transition elements.

Thanks for reading.

REFERENCES

https://en.wikipedia.org/wiki/Helium

http://www.rsc.org/periodic-table/element/2/helium

https://en.wikipedia.org/wiki/Sodium

https://en.wikipedia.org/wiki/Emission_spectrum

https://www.researchgate.net/figure/Sodium-discovery-spectrum-showing-D-1-and-D-2-sodium-resonance-emission-lines-within-the_fig1_225742092

https://en.wikipedia.org/wiki/Ruby_laser

https://www.physics-and-radio-electronics.com/physics/laser/rubylaserdefinitionconstructionworking.html

https://en.wikipedia.org/wiki/Scanning_tunneling_microscope

https://www.nanoscience.com/techniques/scanning-tunneling-microscopy/

https://en.wikipedia.org/wiki/Electron_configuration

https://www.chem.fsu.edu/chemlab/chm1045/e_config.html

https://opentextbc.ca/chemistry/chapter/6-4-electronic-structure-of-atoms-electron-configurations/

https://en.wikipedia.org/wiki/Ionization_energy

Physical and theoretical chemistry

https://en.wikipedia.org/wiki/Electron_shell

https://www.khanacademy.org/science/biology/chemistry--of-life/electron-shells-and-orbitals/a/the-periodic-table-electron-shells-and-orbitals-article




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Hello, thank you for the article.
In your next work, you can maybe tell the reader how to come from the electrons to the orbitals.
Thanks
Chapper

Thank you for coming, @chappertron.

I could have included an indepth discussion on electrons orbitals in this article but I was trying not to make it too lengthy and cumbersome.

On the other hand, I can't also write on the topic alone in another, because it can't be as length as standard post supposed to be.

So, for the benefit of my ever-supportive readers, here's a link to read up more on the electron orbitals.

Thanks.

But if you are in the mood you can write a little summary concerning this subject, @steemstem appreciates this.

Regards

Alright, @chappertron. I love this comment of yours and I promise to do as you have requested, albeit it may be a short article. But I'll start writing on it right away and I'll notify you as soon as it's posted.

Thanks for the love.
Manage the little kiss 😘

👍

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