방출 스펙트럼

When an atom absorbs energy, electrons become excited and move to a higher energy level. As these electrons relax back to a lower energy state or ground state, they release the excess energy as a photon. The wavelength of the emitted light depends on the energy difference between the higher and lower energy levels involved in the transition.

Pure elemental species produce line spectra because electrons can only occupy specific, discrete energy levels. When electrons transition between these quantized energy levels, they emit photons of specific wavelengths rather than a continuous range. This creates distinct spectral lines at characteristic wavelengths unique to each element, which is why line spectra are used to identify substances.

The Balmer series represents spectral lines in the visible light region produced when electrons transition from higher energy levels down to n = 2. For hydrogen, the Balmer series appears as four visible lines at wavelengths of 410, 434, 486, and 656 nanometers, corresponding to transitions from n = 3, 4, 5, and 6 respectively. These characteristic wavelengths make hydrogen's visible spectrum easily recognizable.

The Rydberg formula uses the Rydberg constant and principal quantum numbers to predict all of hydrogen's emission line wavelengths across ultraviolet, visible, and infrared regions. By inputting the lower energy level (n1) and higher energy level (n2) for any electron transition, the formula accurately calculates the resulting wavelength. This mathematical relationship was revolutionary because it explained discrete spectra using a simple empirical expression.

Absorption and emission spectra of hydrogen display lines at identical wavelengths but appear opposite in character. Emission spectrum lines are bright, showing wavelengths of light emitted when electrons relax to lower energy levels. Absorption spectrum lines are dark, representing wavelengths absorbed when electrons are excited to higher energy levels by continuous white light exposure.

Beyond the visible Balmer series, hydrogen exhibits additional spectral series in other regions of the electromagnetic spectrum. The Lyman series appears in the ultraviolet region with transitions to n = 1, while the Paschen series appears in the infrared region with transitions to n = 3. Each series represents electron transitions ending at a different principal quantum number, extending hydrogen's observable spectrum far beyond visible wavelengths.

Different atoms have distinct energy level arrangements, so electrons transitioning between their unique energy levels emit photons of different wavelengths. Since each element possesses its own characteristic set of energy levels, the spectral emission lines vary from element to element. This fundamental difference allows scientists to use emission spectra as a fingerprint to identify and distinguish between different substances.