The atomic emission spectrum refers to the characteristic set of wavelengths of light emitted by atoms when they transition from excited states to lower energy states. Here are some examples of atomic emission spectra:
Hydrogen Spectrum: The hydrogen emission spectrum is one of the most well-known examples. It consists of a series of discrete lines in the visible, ultraviolet, and infrared regions of the electromagnetic spectrum. The Balmer series (visible region) includes the famous H-alpha (656.3 nm), H-beta (486.1 nm), H-gamma (434.0 nm), and H-delta (410.2 nm) lines.
Helium Spectrum: Helium also exhibits an emission spectrum with prominent lines in the visible region. The most notable lines are the D3 line (587.6 nm, yellow) and the D1 (587.5 nm) and D2 (588.9 nm) lines (both in the orange range).
Sodium Spectrum: The emission spectrum of sodium contains a pair of prominent lines known as the sodium D lines. These lines are in the yellow region of the spectrum, with wavelengths of 589.0 nm (D2) and 589.6 nm (D1).
Mercury Spectrum: Mercury vapor produces a complex emission spectrum with multiple lines in the ultraviolet, visible, and infrared regions. Some notable lines include the strong mercury blue line at 435.8 nm (visible) and the mercury green line at 546.1 nm (visible).
Neon Spectrum: Neon emission spectrum features several bright lines in the red, orange, and green regions of the visible spectrum. Notable lines include the neon red line at 640.2 nm, the orange line at 612.8 nm, and the green line at 540.1 nm.
Argon Spectrum: The argon emission spectrum displays a series of lines in the blue and green regions of the visible spectrum. The most intense line is the argon blue line at 488.0 nm.
These are just a few examples of atomic emission spectra. Different elements have their unique spectra due to the specific energy levels and transitions characteristic of their electron configurations. These spectra have been crucial in identifying elements and understanding atomic structure.