Notes

Spectral Lines:

When electrons absorb energy, they jump from lower energy levels to higher energy levels. Recall that electrons in an atom are located around the nucleus at a specific quantum energy level. This energy level is called the ground state of the electron. As the distance of the electron from the nucleus increases, the energy level also increases.

If an atom of an element is exposed to a source of energy, such as light, the electrons in the atom absorb that energy. This enables an electron to jump to a higher energy level and places it in an "excited state."

Electrons cannot, however, remain in an excited state for long. They quickly return to a lower energy state. As they return, they release energy in the form of electromagnetic (EM) radiation. These abrupt movements of electrons between energy levels are known as electronic transitions or quantum jumps. The radiation emitted by atoms as a result of electronic transitions is in the form of photons. These photons are released when electrons return from an excited state to the ground state or any other lower energy state.
The range of EM radiations emitted by an element makes up the spectrum of the element. You might have seen the spectrum of sunlight when it passes through a prism.

The radiation spectra resulting from electronic transitions are different for each atom. You know that EM radiation has different frequencies. White light is made up of all frequencies between infrared and ultraviolet radiation.

During electronic transitions, electrons emit specific frequencies of radiation that depend on the atom. These frequencies are a unique characteristic of an atom, and they correspond to specific frequencies of radiation. Each element will emit only certain sharply defined lines of colors.

The line spectrum emitted by an element is called its atomic emission spectrum. The unique emission spectrum of each element is useful in identifying the element in unknown samples with the help of a spectrometer.
Apart from light, other sources of energy, such as electricity, can also excite electrons and cause electron transition between energy levels. The energy levels are expressed in terms of the principal quantum number, n. The ground state is expressed as n = 1. The subsequent higher energy states are expressed as n = 2, 3, and so on. This number is based on the distance of the electron from the nucleus of the atom. Therefore, larger distances mean greater quantum numbers and, as a result, higher energy levels.
The spectrum of each element contains a specific set of frequencies characteristic of that element.

Energy and Radiation:

According to the classical wave theory of light, light is a continuous electromagnetic wave that has no mass. This theory assumes that the energy in light waves is emitted continuously and can be absorbed and emitted by matter in any quantity. This theory was adopted and accepted until the end of the nineteenth century.

At the beginning of the twentieth century, scientists could not explain the results of certain experiments using classical wave theory. For example, when Max Planck was studying light emitted from heated solids, he observed that as energy, in the form of temperature, was increased, the color given off by the solid changed from red to orange and finally to white. However, there was no increase in the intensity or brightness of the light as would have been expected. Planck could not explain this radiation pattern using the classical wave theory.
Based on his experiments, Planck proposed a theory to explain what classical wave theory could not explain. His theory stated that an atom will not absorb or emit energy continuously, but only in discrete units or packets of energy called photons. The value of each energy packet depends on the frequency of the radiation absorbed or emitted by the object. The frequency is represented by the Greek letter ν (nu). Therefore, the energy of a photon is given by E=hν, where h is a constant called Planck's constant and its value is
6.626 × 10-34 joule seconds.
You know that c=νλ, where c is the speed of light and λ is the wavelength of the radiation.
So, we can also write E=hcλ.
Planck postulated that an atom could only gain or emit energy within integral multiples of hn. So, the energy absorbed or emitted by an electron during its transition to energy level "n" can be given by E=nhν, where n is the energy level the electron reaches
When a specific electron transition occurs in the hydrogen atom, a photon with a wavelength of 21.106 centimeters is emitted. This
"21–centimeter line" of hydrogen is in the radio spectrum and is extensively used in radio astronomy.

Let's calculate the energy change associated with this particular electron transition for hydrogen. Any change in the energy of the hydrogen atom will be equal to the energy of the photon it emits. We know that the energy of a photon is given by E=hν, where h is Planck's constant and v is the frequency.
We need to determine the frequency of the photon. Remember that ν=cλ.
Therefore, the change in energy of the hydrogen atom,
E=hcλ=6.626×10−34Js×2.998×108m/s0.211m
=19.8647×10−260.211joules

=9.4146×10−25joules.
Therefore, when an atom of hydrogen releases a photon having a wavelength of 21.106 cm, the change in energy is 9.4146 × 10−25 joules

Emission and Absorption Lines:

You now know that every element has a unique atomic spectrum. In fact, an element can emit three types of spectra: a continuous spectrum, an emission spectrum, and an absorption spectrum. The spectrum formed from white light consists of all frequencies, and is observed as a continuous band of all radiations. Such a spectrum is called the continuous spectrum. Solids and liquids heated to very high temperatures and gases under pressure produce continuous spectra.

The spectrum emitted by electrons returning to a lower energy level from a higher one is called an emission spectrum. It appears as a series of bright individual emission lines when the light is viewed with a spectroscope. Therefore, this spectrum is also called the line spectrum.

Alternately, when white light is passed through a cool gas, the gas atoms absorb specific frequencies of light and their electrons go the other way—from low energy to high. The resulting spectrum appears as dark lines against a continuous spectrum. These dark lines are the element's absorption lines or absorption spectrum. Absorption lines for an element appear at the exact same frequencies as its emission lines.
The hydrogen spectrum consists of different series of spectral lines in different ranges of EM radiation. The series of lines, or line spectrum, produced depends on the lower energy level involved in the electron transitions.

You have seen in the simulation of atomic spectra that electrons in the excited state may return to the ground state in more than one step. For example, an electron at n = 5 can first jump to n = 3 and then to n = 1. The different spectral series of hydrogen arise because of such transitions.

For example, when electrons in a hydrogen atom jump from higher energy levels to the ground state (n = 1), spectral lines are observed at the ultraviolet end of the spectrum. These lines were first observed by Theodore Lyman in 1906, and were named the Lyman series. The transition of electrons from higher energy levels to n = 2 generate spectral lines in the Balmer series, first observed by Johann Balmer in 1885. This series is in the visible range of the spectrum. When electrons move from higher energy levels to n = 3, the Paschen series is produced. First observed by Friedrich Paschen in 1908, these spectral lines are in the infrared range of the spectrum.
he spectral lines of hydrogen have applications in several fields. In astronomy, spectral lines in the radio spectrum of hydrogen are useful in observing objects that are not in the visible spectrum.

The 21–centimeter line of hydrogen is one example. Most objects in space emit radiation at radio frequencies. Detecting these frequencies with radio telescopes can help in imaging and analyzing celestial objects. Much of our knowledge about the universe is based on observing hydrogen spectral lines in the ranges of visible light, radio waves, and x–rays.

Hydrogen lines of certain frequencies are also used to measure the salinity of ocean surfaces and moisture content of soil, which are very useful for studying the water cycle and predicting the climate conditions on Earth.