Luminescence
Rearranged from Wikipedia, the free encyclopedia; S. Lang
Luminol and hemoglobin, an example of chemoluminescence
Luminescence is light that usually occurs at low temperatures, and is thus a form of cold body
radiation. It can be caused by chemical reactions, electrical energy, subatomic motions, or stress on
a crystal. This distinguishes luminescence from incandescence, which is light generated by high
temperatures. Historically, radioactivity was thought of as a form of "radioluminescence", although
it is today considered to be separate since it involves more than electromagnetic radiation.
The dials, hands, scales and signs of aviation and navigational instruments and markings are often
coated with luminescent materials, in a process known as luminising.
The following are types of luminescence:
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Chemoluminescence
o Bioluminescence
Crystalloluminescence
Electroluminescence
o Cathodoluminescence
Mechanoluminescence
o Triboluminescence
o Fractoluminescence
o Piezoluminescence
Photoluminescence
o Phosphorescence
o Fluorescence
Radioluminescence
Sonoluminescence
Thermoluminescence
On the following pages you can read about Chemoluminescence, Bioluminescence,
Phosphorescence and Fluorescence
On the following pages we focus on Chemoluminescence, Phosphorus, Phosphorescence and
Fluorescence.
Chemiluminescence
From Wikipedia, the free encyclopedia
A chemoluminescent reaction carried out in an erlenmeyer flask producing a large amount of light.
Chemiluminescence (sometimes "chemoluminescence") is the emission of light with limited
emission of heat (luminescence), as the result of a chemical reaction. Given reactants A and B, with
an excited intermediate ◊,
[A] + [B] → [◊] → [Products] + light
For example, if [A] is luminol and [B] is hydrogen peroxide in the presence of a suitable catalyst we
have:
luminol + H2O2 → 3-APA[◊] → 3-APA + light
where:
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where 3-APA is 3-aminophthalate
3-APA[◊] is the excited state fluorescing as it decays to a lower energy level.
The decay of the excited state[◊] to a lower energy level is responsible for the emission of light. In
theory, one photon of light should be given off for each molecule of reactant, or Avogadro's number
of photons per mole. In actual practice, non-enzymatic reactions seldom exceed 1% QC, quantum
efficiency.
A standard example of chemiluminescence in the laboratory setting is found in the luminol test,
where evidence of blood is taken when the sample glows upon contact with iron. When
chemiluminescence takes place in living organisms, the phenomenon is called bioluminescence. A
lightstick emits a form of light by chemiluminescence.
Chemiluminescence of luminol (voluntary)
To exhibit its luminescence, the luminol must first be activated with an oxidant. Usually, a solution
of hydrogen peroxide (H2O2) and a hydroxide salt in water is used as the activator. In the presence
of a catalyst such as an iron compound, the hydrogen peroxide is decomposed to form oxygen and
water:
2 H2O2 → O2 + 2 H2O
In a laboratory setting, the catalyst used is often potassium ferricyanide. In the forensic detection of
blood, the catalyst is the iron present in hemoglobin. Enzymes in a variety of biological systems
may also catalyze the decomposition of hydrogen peroxide.
When luminol reacts with the hydroxide salt, a dianion is formed. The oxygen produced from the
hydrogen peroxide then reacts with the luminol dianion. The product of this reaction, an organic
peroxide, is very unstable and immediately decomposes with the loss of nitrogen to produce 3aminophthalic acid with electrons in an excited state. As the excited state relaxes to the ground
state, the excess energy is liberated as a photon, visible as blue light.
Use by crime scene investigators
Theory
Luminol is used by crime scene investigators to locate traces of blood, even if it has been cleaned or
removed. The investigator prepares a solution of luminol and the activator and sprays it throughout
the area under investigation. The iron present in any blood in the area catalyzes the chemical
reaction that leads to the luminescence revealing the location of the blood. The amount of catalyst
necessary for the reaction to occur is very small relative to the amount of luminol, allowing the
detection of even trace amounts of blood. The glow lasts for about 30 seconds and is blue.
Detecting the glow requires a fairly dark room. Any glow detected may be documented by a long
exposure photograph.
Drawbacks
Luminol has some drawbacks that may limit its use in a crime scene investigation:
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Luminol chemiluminescence can also be triggered by copper or copper-containing alloys,
horseradish, and certain bleaches; and, as a result, if a crime scene is thoroughly cleaned with a
bleach solution, residual bleach will cause the entire crime scene to produce the typical blue glow,
effectively camouflaging any organic evidence, such as blood.
Luminol will also detect the small amounts of blood present in urine and it can be distorted if
animal blood is present in the room that is being tested.
Luminol reacts with fecal matter, causing the same glow as if it were blood.
Luminol's presence may prevent other tests from being performed on a piece of evidence.
However, it has been shown that DNA can be successfully extracted from samples treated with
luminol reagent.[4]
Bioluminescence (voluntary)
Chemiluminescence takes place in numerous living organisms, the American firefly being a widely
studied case of bioluminescence.
The firefly reaction has the highest known quantum efficiency, QC of 88%, for chemiluminescence
reactions. ATP (adenosine tri-phosphate), the ubiquitous biological energy source, reacts with
luciferin with the aid of the enzyme luciferase to yield an intermediate complex. This complex
combines with oxygen to produce a highly chemiluminescent compound.
Phosphorus
Rearranged from Wikipedia, the free encyclopedia; S. Lang
Phosphorus (IPA: /ˈfɒsfərəs/) is the chemical element that has the symbol P and atomic number
15. The name comes from the Greek: φώς (meaning "light") and φόρος (meaning "bearer"). A
multivalent nonmetal of the nitrogen group, phosphorus is commonly found in inorganic phosphate
rocks.
Phosphorus is a component of DNA, RNA, ATP, and also the phospholipids which form all cell
membranes. It is thus an essential element for all living cells. The most important commercial use
of phosphorus-based chemicals is the production of fertilizers.
Phosphorus compounds are also widely used in explosives, nerve agents, friction matches,
fireworks, pesticides, toothpaste and detergents.
Precautions
White phosphorus should be kept under water at all times as it presents a significant fire hazard due
to its extreme reactivity with atmospheric oxygen, and it should only be manipulated with forceps
since contact with skin can cause severe burns. Chronic white phosphorus poisoning leads to
necrosis of the jaw called "phossy jaw". Ingestion of white phosphorus may cause a medical
condition known as "Smoking Stool Syndrome". [20]
Glow from white phosphorus
Phosphorus (Greek. phosphoros, meaning "light bearer") was discovered by German alchemist
Hennig Brand in 1669. Working in Hamburg, Brand attempted to distill some kind of "life essence"
from his urine, and in the process produced a white material that glowed in the dark. The
phosphorus had in fact been produced from inorganic phosphate, which is a significant component
of dissolved urine solids. White phosphorus is highly reactive and gives-off a faint greenish glow
upon uniting with oxygen. The glow observed by Brand was actually caused by the very slow
burning of the phosphorus, but as he saw no flame nor felt any heat he did not recognize it as
burning.
It was known from early times that the glow would persist for a time in a stoppered jar but then
cease. Robert Boyle in the 1680s ascribed it to "debilitation" of the air; in fact, it is oxygen being
consumed. By the 18th century, it was known that in pure oxygen phosphorus does not glow at
all;[1] there is only a range of partial pressure at which it does. Heat can be applied to drive the
reaction at higher pressures.[2]
In 1974, the glow was explained by R. J. van Zee and A. U. Khan.[3] A reaction with oxygen
takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived
molecules HPO and P2O2 that both emit visible light. The reaction is slow and only very little
of the intermediates is required to produce the luminescence, hence the extended time the
glow continues in a stoppered jar.
Although the term phosphorescence is derived from phosphorus, the reaction which gives
phosphorus its glow is properly called chemiluminescence (glowing due to a cold chemical
reaction), not phosphorescence (re-emitting light that previously fell onto a substance and
excited it).
Phosphorescence is the slow decay of a metastable electronic state to a lower energy state
through emission of light. The decay is slow because the transition from the excited to the
lower state requires a spin flip, making it classically forbidden.
In contrast, chemiluminescence occurs when the product molecules of a chemical reaction
(HPO and P2O2 in this case) leave the reaction in an electronically excited state. These excited
molecules then release their excess energy in the form of light. The frequency (color) of the
light emitted is proportional to the energy difference of the two electronic states involved.
Phosphorescence
Explanations of the phenomenon
Phosphorescence is a specific type of photoluminescence related to fluorescence. Unlike
fluorescence, a phosphorescent material does not immediately re-emit the radiation it absorbs.
Figure 5: Phosphorescence
First, let's distinguish phosphorescence from fluorescence. A fluorescent paint glows under a UV
lamp, but stops glowing as soon as the lamp is turned off. A phosphorescent paint keeps glowing
for a while.
Phosphorescent substances have the ability to store up light and release it gradually. The notion of a
metastable state explains this. If the molecules of the substance can get from the ground state to a
metastable state, and if the metastable state can slowly decay back to the ground state via photon
emission, then we have phosphorescence.
Typically, the metastable state is a triplet state, and the ground state is a singlet state. Ground state
molecules absorb photons and go to excited singlet states (see figure 5). Most of them immediately
hop right back to the ground state, emitting a photon, but non-radiative processes take a few to a
less energetic triplet state. Once these molecules get to the lowest triplet state, they are stuck there,
at least for a while. Some low probability process accomplishes the triplet-singlet conversion, and
the molecules slowly leak out light.
Phosphorescent powder under visible light, ultraviolet light, and total darkness.
Confusion with chemiluminescence
Some examples of "glow-in-the-dark" materials do not glow because they are phosphorescent. For
example, "glow sticks" glow due to a chemiluminescent process which is commonly mistaken for
phosphorescence. In chemi-luminescence, an excited state is created via a chemical reaction. The
excited state will then transfer to a "dye" molecule, also known as a (sensitizer, or fluorophor), and
subsequently fluoresce back to the ground state.
[edit] Common components in phosphorescent materials
Common pigments used in phosphorescent materials include zinc sulfide and strontium aluminate.
Use of zinc sulfide for safety related products dates back to the 1930s. However, the development
of strontium oxide aluminate, with a luminance approximately 10 times greater than zinc sulfide,
has relegated most zinc sulfide based products to the novelty category. Strontium oxide aluminate
based pigments are now used in exit signs, pathway marking, and other safety related signage.
Strontium aluminate based afterglow pigments are marketed under brandnames like SuperLumiNova[1][2] or NoctiLumina.[3]
Fluorescence
From Wikipedia, the free encyclopedia
Fluorescent minerals
Fluorescence is a luminescence that is mostly found as an optical phenomenon in cold bodies, in
which the molecular absorption of a photon triggers the emission of a photon with a longer (less
energetic) wavelength. The energy difference between the absorbed and emitted photons ends up as
molecular rotations, vibrations or heat.
Sometimes the absorbed photon is in the ultraviolet range, and the emitted light is in the visible
range, but this depends on the absorbance curve and Stokes shift of the particular fluorophore. The
term 'fluorescence' was coined by George Gabriel Stokes in his 1852 paper; the name was given as
a description of the essence of the mineral fluorite, composed of calcium fluoride, which gave a
visible emission when illuminated with "invisible radiation" (UV radiation).
Contents
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1 Equations
o 1.1 Photochemistry
o 1.2 Quantum yield
o 1.3 Lifetime
2 Rules
3 Applications
o 3.1 Lighting
o 3.2 Analytical chemistry
o 3.3 Biochemistry and medicine
o 3.4 Gemology, mineralogy, geology, and forensics
o 3.5 Organic liquids
4 See also
5 External links
6 Further reading