Light and Electromagnetic Radiation
Light is a form of electromagnetic radiation, which is energy traveling as waves. These waves are characterized by their frequency (f) and wavelength (λ). The relationship between frequency and wavelength is described by the equation c = fλ, where c is the speed of light (3.00 x 10⁸ m/s).
The energy (E) of electromagnetic radiation is related to its frequency by the equation E = hf, where h is Planck's constant (6.63 x 10⁻³⁴ J·s). This relationship shows that as frequency increases, energy increases, while wavelength decreases.
Highlight: The electromagnetic spectrum encompasses the entire range of electromagnetic radiation frequencies, from low-energy radio waves to high-energy gamma rays.
The visible spectrum, which is what we perceive as light, occupies a small portion of the electromagnetic spectrum. It ranges from about 400 nm (violet) to 700 nm (red) in wavelength.
Vocabulary: Wavelength is the distance over which a wave's shape repeats, while frequency is the number of times the wave repeats in one second.
Absorption Spectra
When white light (containing all visible wavelengths) passes through the vapor of an element, certain wavelengths are absorbed by the atoms. This process creates an absorption spectrum, characterized by dark lines in the otherwise continuous spectrum.
Definition: An absorption spectrum is a series of dark lines or bands in an otherwise continuous spectrum, resulting from the absorption of specific wavelengths of radiation by atoms or molecules.
These absorption lines correspond to the energy required to promote electrons from lower to higher energy levels within the atoms.
Emission Spectra
Emission spectra are produced when atoms are energized, typically by heat or an electrical field. This excitation causes electrons to move to higher energy levels. When the electrons return to lower energy states, they release energy in the form of photons with specific frequencies.
Example: The emission spectrum of hydrogen shows distinct colored lines corresponding to specific electron transitions, such as the Balmer series in the visible region.
The discrete nature of emission spectra provides evidence for the quantized energy levels in atoms. If electron energy levels were continuous rather than quantized, we would observe a continuous spectrum instead of distinct lines.
