@Kalyani_Kumari see the spectra of various elements are their characteristic property..just as we have malleability, ductility, so in layman's terms, u can get it that every element has its own characteristic atomic spectrum, and the method of observing, detecting and extracting information from these spectra of atoms, is called spectroscopy. It is a very good method to identify the elements... various elements like Thallium, Indium, Rubidium, etc had been discovered from their atomic spectra only. Helium was discovered during a solar eclipse, by a French Astronomer. So, what basically an atomic spectrum is, is that, it is just like the photos that we click. Actually, in layman's terms, u can suppose of the atomic spectrum as the bands of wavelengths emitted or absolutely absorbed by that atom, such that they are obtained on a photographic plate, for further analysis...just as see , the old cameras were there, where I had those reels, and once clicked, u had to develop them... similar case is for atomic spectrum as well...
Atomic spectra are of broadly two types. One is Emission Spectrum and one is Absorption Spectrum.
In emission spectrum, what basically we do, is that we obtain the radiations that are emitted by the atom under observation, on a photographic plate, and then are further analysed. So what we basically do here, is that we take any solid element and we excite the atoms, and hence the electrons present in them, via either irradiation or through other methods of excitation..and hence, due to the energy gained by the electrons, they jump to higher levels and then as per Bohr's Hypothesis, they come down to the lower energy states, along with the emission of energy in the form of electromagnetic radiations, which are taken on a photographic plate. On such a plate, we find (in emission spectrum), bright lines, and rest all dark patches. These bright lines are actually the wavelengths of radiations that are emitted by that particular atom. Thus, every atom shows a characteristic different set of radiations on the plate and hence the spectrum of every atom is different. Thus on the photographic plate, we have only bright lines corresponding to the emitted wavelengths and rest all dark.
Absorption Spectrum can be through of exactly the photographic negative of the emission spectrum. What happens here is that we take a solution or molten form of the metal, and then we irradiate them with light or any means of excitation. Here, the best part is, the substance (solution) will allow all the radiations to pass through it ( and hence detected at the photographic plate), except those radiations, which are a characteristic is that element. Characteristic here means, that suppose \lambda_1 and \lambda_2 are two incident wavelengths..and let to go from the first excited state to the other, the atom needs \lambda_1 wavelength. So, this wavelength becomes the characteristic wavelength of that element, and hence will absorb it and the other one will be allowed to pass through its solution, by the element. So what happens at the photographic plate is that, we get all those wavelengths which are not the characteristic wavelengths of the element (and we're hence allowed to pass through its solution), except those which are absorbed by it, which causes the photographic plate to have only dark lines (corresponding to absorbed ones), and rest everywhere is bright. Thus, when compared emission and absorption spectra are photographic negative of each other.
So basically what u were asking about infrared, and UV was this only. When the electron used to de-excite from any higher level to lower level, the emitted wavelengths belonged to a region of the electromagnetic spectrum. When the electron, after absorbing energy, just only came to the first excited state (n=2), then the wavelengths were in visible region of the electromagnetic spectrum. Thus, the collection of all those emitted wavelengths, which were lying in the visible region, that is, which had their wavelengths ranging from around 3600 to 7200 angstroms nearly, were all grouped into a series, and named as the Balmer Series. This name was given after the scientist who discovered them, and the visible one was done by Balmer.
Now to identify which radiation belongs to which region, that is infrared or so, we use the Rydberg equation.
Here, as I said earlier, u need to see where the electron stops after de excitation. If it stops at n_1=1, then the wavelengths lie in the Lyman Series. When it depends excites, and reaches by maximum, to the first excited state (n1=2), then it corresponds to the visible region and thus corresponds to the Balmer Series. For all the other details excitations, that is n1>=3, the wavelengths all lie in the infrared region. n2 corresponds to the excited level, where the electron jumps after absorbing energy and n1 corresponds to the level where the he electron reaches after de excitation. When n1=3, the series is Paschen. What these series are, are just the bundles of wavelengths, where the least state after de excitation is the n1 state which demarcates the difference between them.
Do let me know if u have any doubts.