|Blacklight is another term for an
ultra-violet (uv) light.
To the eye, the light they emit seems to be an intense electric purple that doesn't really brighten anything normal unless it's white.
When used in low light on anything uv-responsive, the effect is astounding. Colours become suddenly intensely rich and vibrant and hover above the surface, sometimes appearing to hang in mid-air.
Ultra Violet (UV) light represents a section of the overall electromagnetic spectrum of light, extending from the blue end of the visible (400nm) to the x-ray region (100nm).
It is subdivided into three distinct wavelength regions described as either UV-A, UV-B or UV-C in increasing order of photon energy.
UV-A 400nm-315nm: Often referred to as 'blacklight', this is the longest wavelength region and lowest energy, it represents the largest portion of natural UV light.
UV-B 315nm-280nm: Partially blocked by the ozone layer this is the most aggressive component of natural UV light and largely responsible for sunburn (erythema).
UV-C 280nm-100nm: Only generally encountered from artificial light sources since it is totally absorbed by the earth's atmosphere.
Fluorescence and phosphorescence - the theory
The excitation energy provided by UVA photons is much higher than the energy of the thermal motions of the molecules at physiological temperatures. Thus the absorbing molecules temporarily assume energy levels that otherwise they would never attain and thus acquire properties differing considerably from those effective in ordinary chemistry.
The lifetime of a molecule in its usual excited state (10-10 to 10-8 sec), which is still long compared with the time required for the energy absorption itself (approximately 10-15 sec), can be greatly extended if the excited electron is trapped in an (energetically somewhat lower) triplet excited state. In contrast to the usual singlet state, the triplet state is characterised by two electrons with unpaired spin. Because the return from the triplet state to the ground state is "forbidden" (i.e. occurs at a low probability), the triplet may last 10-3 sec or even longer and is, therefore, called metastable.
As an excited electron returns to a lower energetic state, its excess energy can be emitted as a photon, resulting in fluorescence. Fluorescent light is recognised by its usually longer wavelength, compared with the exciting radiation. Emission from molecules in the metastable excited state occurs over a longer period of time and is called phosphorescence.
The theory behind molecular luminescence (fluorescence and phosphorescence) is fairly well understood. When light hits a substance, the incoming energy may pass on through or it may be temporarily absorbed. This event takes around 10 -15 seconds (Guilbault 1973). The energy that is absorbed by a molecule becomes stored as increased electron vibrational or even rotational motion, and, if there is sufficient energy, as an elevation in the molecule energy states (molecular excitation). The whole process is best explained with quantum mechanics theory as follows. Every molecule has a series of energy levels, both main electron energy states and subsidiary vibrational levels. When a molecule is bombarded with energy, the molecule graduates from what is called the ground energy state, the lowest energy level, to one of a number of higher excited states. These are electronic states; in other words, they relate to the relative energy level demonstrated by the molecule's electrons. In addition, every state has a series of vibrational levels which the electron can occupy. After absorbing energy, an electron is typically elevated to a higher level, either a higher vibrational level for smaller amounts of energy, or a higher, excited, energy state for larger amounts. Photons of visible light and especially those of ultraviolet light wavelengths typically have sufficient energy to cause a transition into one of the excited states. If the energy is too little, only changes in the amount of vibrational or rotational energy level occur; if too much, photodecomposition can occur (see below). Further, the energy must be of a type that is appropriate to the molecular structure; excitation energy that is at less ideal wavelengths may still produce fluorescence, but at a lower intensity (Guilbault 1973). (Becker 1968; Dake and De Ment 1941; Guilbault 1973; Hurtubise 1986; McGown 1986; Wain 1965)
Over the next 10-4 seconds, the molecule relaxes, losing its excess vibrational energy in small amounts (vibrational relaxation) until the lowest energy level of the present state is attained (Guilbault 1973). Since excited states are unstable, an excited molecule must lose still more energy. For this, several paths are possible (Becker 1968: 76-8; Dake and De Ment 1941:51-2; Hurtubise 1986; Wain 1965:12-15). The loss may involve a direct jump to the ground state. This route can release a large quantum of energy in the form of a photon, which we see as fluorescence. Since the jump may terminate at any of the ground state's vibrational levels, the resulting spectrum for an entire sample will typically exhibit several wavelengths, the intensity of each corresponding to the likelihood of a particular jump (Schulman 1985). While excitation depends on the overall energy of the bombarding light, the resulting emission spectrum is independent of its wavelength (McGown 1986). This final decay process takes roughly 10-9 to 10-7 seconds. Since some energy is typically lost due to vibrational relaxation prior to fluorescence, the energy of the emitted light is less than that of the absorbed light, and the wavelength is corresponding longer. In other words, while the substance was bombarded with ultraviolet light, the fluorescence will be either higher in teh ultraviolet spectrum, or more typically up into the visible spectrum, or even in the infrared.