What Can Happen to an Electron When Sunlight Hits It

When an electron is hit by a photon of light, information technology absorbs the quanta of energy the photon was carrying and moves to a higher energy land.

One way of thinking about this college free energy state is to imagine that the electron is now moving faster, (it has but been “striking” by a chop-chop moving photon). But if the velocity of the electron is now greater, it’s wavelength must also have changed, so it tin no long stay in the original orbital where the original wavelength was perfect for that orbital-shape.

And so the electron moves to a different orbital where once more its own wavelength is in phase with its self.

Electrons therefore accept to leap effectually within the cantlet as they either gain or lose energy. This property of electrons, and the energy they absorb or give off, tin can be put to an every day use.

Well-nigh any electronic device you purchase these days comes with 1 or more
Light Emitting Diodes
(commonly chosen “LEDs“). These are tiny bubbles of epoxy or plastic with two wire connectors. When electricity is passed through the diode it glows with a feature color telling you that the device is working, switched on and ready to do it’s work.

Deep in the semiconductor materials of the LED are “impurities”, materials such as aluminum, gallium, indium and phosphide. When properly stimulated, electrons in these materials motility from a lower level of energy upward to a higher level of energy and occupy a different orbital.

Then, at some point, these college energy electrons give up their “extra” energy in the course of a photon of light, and fall back downwardly to their original energy level. The light that has suddenly been produced rushes abroad from the electron, atom and the LED to colour our earth.

Typically, the low-cal produced by a LED is only one color (red or green being strong favorites). Although they are cheap, easy to make, don’t toll a lot to run, LEDs are
not
usually used to lite a room, because they cannot normally produce the wide range of dissimilar colors needed in “white” calorie-free.

This is considering of the quantum nature of the atoms being used in the LED and the quantum energies of the electrons within them.

When an excited electron within a LED gives up energy it must do and then in those lumps called
quanta. These are fixed packets of energy that cannot exist changed or used in fractions; they must e’er be transferred in whole amounts.

Thus, an excited electron has no selection but to give off either one quanta or ii quanta of energy, information technology cannot give up 1.five quanta, or 2.3 quanta. As well, the electron can only move to very express orbitals within the atom; it must end up in an orbital where the wavelength is now uses is “in phase” with itself. These two restrictions limit the quality of the quanta of energy being released past the electron, and thus the nature of the photon of light that rushes abroad from the LED.

Since the energy given off is strongly restricted to quanta, and quanta that allow the electron to move to a suitable place inside the atom, the photons of light are similarly restricted to a tiny range of values of wavelength and frequency (a property we meet equally “colour”).

Many LEDs have electrons that can only requite upwardly quanta of energy that, when converted into photons, produce low-cal with a wavelength of about 700 nm – which we then come across as

red

light. These electrons are so restricted in the quanta they can emit that they never shine bluish light, or green low-cal, or xanthous low-cal, only red light.

Long, long before their were LEDs in our lives, scientists trying to empathise electrons in atoms noted a similar miracle when lite was either shone on certain materials or given off by certain materials.

In 1859 the German physicist Gustav Robert Kirchoff, and his older friend Robert Wilhelm Bunsen came up with a clever idea. They used Bunsen’s burner to strongly rut tiny pieces of various materials and minerals until they were and then hot that they glowed and gave off low-cal.

Sodium, for example, when heated to incandescence, produced a strong xanthous calorie-free, but no blueish, green or red. Potassium glowed with a dim sort of violet low-cal, and mercury with a horrible light-green light but no cerise or xanthous.

When Kirchoff passed the emitted light through a prism it separated out into its diverse wavelengths (the same way a rainbow upshot is produced when white light is used), and he got a daze. He could just see a few thin lines of light in very specific places and often spread far apart.

Clearly glowing sodium was not producing anywhere near all the dissimilar wavelengths of white lite, in fact it was only producing a very characteristic band of lite in the yellow region of the spectrum – just similar a LED!

Kirchoff and Bunsen carefully measured the number and position of all the spectral lines they saw given off by a whole range of materials. These were called


emission spectra

, and when they had collected enough of them information technology was articulate that each substance produced a very feature line spectrum that was unique. No two substances produced exactly the aforementioned serial of lines, and if two unlike materials were combined they collectively gave off
all
the lines produced past both substances.

This, thought Kirchoff and Bunsen, would exist a good way of identifying substances in mixtures or in materials that needed to be analyzed. So they did. In 1859 they found a spectrum of lines that they had never seen before, and which did not represent to whatever known substance, and then, quite rightly, they deduced that they had institute a new element, which they called

cesium

from the Latin word meaning “sky blue”. (Judge in what part of the spectrum they found the lines!).

All the research on atomic construction and the hideously difficult-to-understand properties of electrons come together in the topic of “electron energy”.

An cantlet such as
lithium
has iii electrons in diverse orbitals surrounding the atomic center. These electrons can be bombarded with energy and if they absorb enough of the quanta of energy being transferred they jump near and in the almost extreme example, get out the lithium atom completely. This is called

ionization
.

The corporeality of energy needed to remove the first electron from a lithium is 124 kilocalories/mole, an amount of energy that is not difficult to supply, so lithium atoms ionize hands.

Notwithstanding, it takes most 1740 kilocalories/mole of free energy to dislodge the 2nd electron from effectually the lithium ion (it is now an “ion” because it has already lost one electron). It takes a massive 2820 kilocalories/mole to dislodge the third and final electron from around the lithium ion.

Partly this difference in the corporeality of free energy needed to dislodge different electrons away from the lithium diminutive centre is due to the fact that the center of the lithium cantlet is carrying the positive charges of three protons. Moving a negatively charged electron away from a positively charged atomic center needs more and more energy as the amount of united nations-neutralized charge increases, thus;

Yet, the amount of energy needed to remove the first electron is a good measure of what it takes to stimulate an electron to go out its atom, and how tightly it is held there in the first identify.

Within the atom, as Bohr pointed out, there are dissimilar possible positions for electrons to be constitute as divers by the


primary quantum number

, usually written every bit “
n
“.

Bohr divers the energy of electrons located at these different locations of quantum state by the formula:

In this formula

Due east_o


is a whole drove of concrete constants, which for an cantlet such equally hydrogen has a value of 313 kilocalories/mole. Using this formula it is possible to calculate how much energy an electron has at each of the other, different, quantum states (n = 2, n = three, n = 4, etc.). This is usually presented in the course of a diagram (see left).

For an electron at the ground state (n = one) to be moved up to the next level (n = ii) information technology must absorb a quantum of energy that is the perfect amount to make this motion. If the quantum is too small the electron could not reach the next level, then it doesn’t endeavor. If the quantum is too big the electrons would overshoot the side by side level, so over again, it does not endeavour. Only quanta of exactly the correct size will be absorbed and used.

Similarly, if an electron is already at the second level (n = 2), and there is a space for the electron at the lower level (due north = 1), it can release a quantum of free energy and driblet down to the lower level. Merely the amount of energy given off will be a whole number quantum. If this energy is given off as calorie-free (such as happens with emission spectra) then the photons rushing away from the falling electron volition be of merely ane size and quality (colour). Hence glowing sodium, or LEDs, only requite off very detached bands of low-cal with distinct colors or bands within their spectrum.

All this implies that if white light (with all the possible wavelengths, colors and possible quanta of energy) is shone on certain materials or substances only certain wavelengths (and their quanta of free energy) volition exist absorbed by the electrons in that substance. Only a narrow ring of calorie-free will take just the right quanta to move an electron to the next level, or the level above that, and so on.

That wavelength will be taken out of the spectrum of lite and leave a nighttime band of no-calorie-free behind. Absorption spectroscopy, therefore, is the equal and opposite of emission spectroscopy. However, in both kinds, it is the absorption of quanta to move electrons, or the emission of quanta to move electrons around in the atom that is the reason why only certain wavelengths of light are affected.

Although Bohr’s original flick of a quantum atom has been modified in the years since he start proposed the concept, never the less, the chief principles however stand:

1. Electrons are to exist found occupying certain volumes of space around an diminutive center (“nucleus”) – these volumes of space are called
   orbitals

2. An electron in an orbital has a defined wavelength. The actual wavelength tin be determined using the de Broglie formula “wavelength = Plank abiding / momentum.

3. The shape and location of the orbital is determined past the fact that the only stable shapes and locations are those where the electrons (acting as waves) can have a number of waves that are whole numbers (technically these are called “
   standing waves
   “). Standing-wave orbitals are the only ones in which the occupying electrons exercise not either radiate energy, or collapse.

4. The energy carried past electrons has to exist a whole number of quanta of energy as given past the formula
   
   Due east_n
   = – Eastward_o/n^two
   
   

   where “n” is the principal quantum number. The energy of an electron, and the atom that carries it, is therefore restricted, or
   
   quantized,
   
   to a limited number of values.