What is Atomic Spectra? Explained for Class 12 Physics and Beyond

Atomic spectra are the collection of specific wavelengths of light. We observe these phenomena in real life and through microscopes when atoms emit or absorb light. This could happen when electrons of atoms move between different energy levels. 

The basic nature of atomic spectra is governed by anything that energises an atom's electrons. For instance, when heat, electric current, or light energises atoms, their electrons jump to higher energy states and then fall back down. Then they continue to release photons with wavelengths with different characteristics. 

As a result, this energy transfer creates a unique pattern of bright or dark lines. By learning about this pattern, anyone can easily identify which element is which.

Atomic spectra are produced whenever there is a difference in the energy states of electrons. 

To understand this energy change, we need to know the aftermath of what happens to an electron after we heat it, light it, or pass some electrical current to it. So, the treatment you choose, will tell or show you that an electron will absorb or emit energy. 

Let's say the electron jumps up to a higher energy level. That's an excited state for the electron. 

The excited state tells us that the electrons cannot ever remain at rest. So, the next thing after reaching an excited state, they would quickly fall back down to lower energy levels. 

And when they do, they release that excess energy as photons. These are particles of light. They will have specific wavelengths, each of them, in fact.

It’s the energy difference between levels that determines the exact colour (wavelength) of light that is emitted. 

Remember that each element will always have its own unique arrangement of electron energy levels. What we mean is that every one of them will have its own characteristic pattern of wavelengths. That is the atomic spectrum.

And this basic understanding would be required for your CBSE boards. 

We just need to cover the basics of three types of atomic spectra. 

Emission Spectra

Emission spectra look quite close to neon signs. It comes up when you heat a gas or pass electricity through it. That makes the electrons to get excited, that is, moving up to a higher energy level. Consequently, they emit light. 

These appear usually as bright, coloured lines with a dark background. Each line in the emission spectra tells a specific electron transition from a higher to a lower energy state. 

Absorption Spectra

In comparison with emission spectra, absorption spectra work in the opposite direction. You guessed it, in absorption spectra, there is this movement of electrons from a lower to a higher energy state. 

Consider here that you shine white light through a cool gas, the atoms absorb specific wavelengths. That is still exciting their electrons. 

What you see here is a continuous rainbow with dark lines. Only certain colours in this scenario have been stolen/taken by the atoms. In other words, absorbed. 

It's like someone with a pair of molecular scissors cutting out specific colours from a rainbow.

Continuous Spectra

Continuous spectra are visible when matter is so hot and dense that individual atomic transitions get blurred. It creates a smooth progression of colours. It’s how you see a rainbow in its perfect form, there are different colours, but merging. 

The core of the Sun produces a continuous spectrum, though the cooler gases in its atmosphere add absorption lines to create the spectrum we actually observe.

After the Rutherford Scattering Experiment, the global scientific community adopted the view that the atom is a tiny solar system. It was like the concentrated mass of the nucleus, centred like the sun, while electrons revolved around it, similar to how planets orbit around the sun. 

But this nuclear model of the atom couldn’t explain, or even puzzle many of Rutherford’s contemporaries, why electrons emit only certain energies of light and how an atom could be stable if the gravity of the nucleus would pull the electrons into the centre. 

That’s when we see the rise of quantum mechanics, when Niels Bohr in 1913 explained that electrons do not randomly orbit around the nucleus. 

Bohr, instead, found out that electrons exist in “discrete energy levels”. That would appear something like the rungs on a ladder. It was also found out that electrons cannot exist between these energy levels. And each spectral line corresponds to an electron transition between specific energy levels.
How did Bohr find this out? Go back to the basics to remember the developments leading to the Bohr Model of Atom. 

Line Spectra vs Continuous Spectra

A continuous spectrum looks like a smooth rainbow. It looks more like a blending of one colour into the next. 

Atomic spectra are different. Instead of a smooth flow, they show up as sharp lines.

These can be bright emission lines or dark absorption lines.

That lined pattern comes from the fact that electrons in atoms can only jump between fixed energy levels.

Element-Specific Patterns

Every element has its own unique fingerprint in light.

Hydrogen’s atomic spectrum doesn’t even look anything like helium’s.

Carbon has its own signature too.

This is why scientists can instantly tell which elements are present, just by looking at their atomic spectra.

Intensity Variations

Not all lines shine with the same brightness.

Some glow strongly, others are faint.

The strength of a line depends on things like temperature, pressure, and how likely a certain electron jump is.

These variations give extra clues about the source.

Wavelength Precision

Spectral lines are extremely precise.

They sit at very specific wavelengths, often measured to several decimal places.

This strong accuracy lets scientists detect tiny shifts in wavelength. For instance, when an object moves towards or away from us (the Doppler effect) or when magnetic fields are present.

Temperature Sensitivity

Atomic spectra also change with heat.

The lines that show up and their relative brightness can look very different at different temperatures.

A hot star can display one set of features, while a cooler star with the same elements shows another.

Hydrogen is the simplest atom with just one electron. It became the testing ground for Bohr's theory. Also, if you remember from your earlier Chemistry topic on the Bohr Model of the Atom, you know the shell designations starting from K (n=1) that’s closest to the nucleus, and it’s the ground state. 

This idea will help when you learn the different lines in the atomic spectrum of hydrogen. 

These series of lines are called the spectral series. Each of them were discovered by different scientists. 

• Lyman Series

The Lyman series occurs when electrons fall all the way down to the ground state (n=1). It’s closest to the nucleus. 

Now these transitions from one state to another end up releasing high-energy photons that are visible in the ultraviolet range. We can't see them, but are detectable by instruments.

• Balmer Series

The Balmer series is the only series visible to the human eye. In this, the electrons drop to the second energy level (n=2) from higher levels. They begin to emit the red, blue-green, and violet lines. These may be quite familiar in school. 

The bright red line (Hα at 656.3 nm) is particularly prominent in astronomical observations.

• Paschen, Brackett, Pfund, and Humphreys Series

If you have heard of infrared radiation, it’s the Paschen, Brackett, Pfund, and Humphreys series. Respectively, the electrons drop down to the third, fourth, fifth, and sixth energy levels.

Even these are invisible to the eyes, but they help us study interstellar gas clouds and similar. 

This spectral series is understood by the Rydberg formula. 

1/λ = R(1/n'² - 1/n²). 

R is the Rydberg constant

n' is the lower energy level, and 

n is the upper energy level. 

This simple equation accurately predicts the wavelength of every hydrogen spectral line.

For Astronomical Findings

The star’s spectrum can inform us about its chemical composition, temperature, and motion. Studying and analysing the atomic spectra in broad daylight, physicists have figured out that the Sun contains hydrogen, helium, and iron, along with other elements.

To Monitor the Environment

Atomic absorption spectroscopy enables the detection of any amount of heavy metals in water supplies. Normally, one would shine light through a sample to measure which wavelengths are absorbed. 

Quality Control in the Pharmaceutical Industry

Drug manufacturers use spectroscopic techniques to verify purity and detect contamination. They can find the characteristic spectrum of each compound. That’s how they assure quality of the products.