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What is the brightest color a light source that only emits blue frequency light can achieve?

What is the brightest color a light source that only emits blue frequency light can achieve?


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Let's say there is a luminous light bulb that only emits blue light, and you (or a camera if you prefer) look at it from a close distance. Will the maximum brightness perceived not have any "white" to it? Because there is no white light if it's all blue. Probably the brightest it could be would be solid blue? Therefore, blue photons could produce a range of color of black to solid blue depending on the luminous or amount of photons hitting your eye per second? If that's the case then what happens if you raised the intensity of the bulb to an extreme? Is there a cut off point where even if you add more intensity to the bulb, you still see the same solid blue color?


The color that human vision perceives for a specific frequency of light (prior to saturation of any cones) is derived from the relative responses of different cones. For example, at about 480nm, the response of S and M cones are roughly equal. However, the human eye is not able to distinguish pure colors (single wavelengths) from combinations of wavelengths which produce the same response in each of the three cone types.

It is possible that a blue light source could appear white to the human eye, because the cones which inform our perception of color are not strictly monochromatic. Rather, they have a range of sensitivity across the spectrum, and all cones are at least somewhat sensitive to blue light. (See this diagram.) If the blue light is so intense that it saturates all cones, S (blue), M (green), and L (red), then it will be perceived as white light. However, in order to saturate the low-sensitivity M and L cones, the intensity of the light would be blindingly high. Prior to this point, but after saturating the S cones, the apparent color of the light would shift toward green before eventually appearing white. For intensities that are not saturating to any cones, the hue of the color blue will remain constant, while it's value (or lightness) may change.


Difference Between Daylight and Soft White LED Bulbs

That being said, the common terminology associated with energy-efficient LED lights is Daylight, Bright White, and Soft White. These are nothing but representations of the quantitative value of Kelvins. We will only discuss the Daylight and Soft White LED light bulbs in this article. Daylight is a very bright white-blue light with a very high color temperature in the range of 5000 – 6500 K. They reflect colors naturally just like Warm White LED lights creating a near-perfect natural effect. Soft White produces a yellow hue and a lower color temperature in the range of 2700 – 3000 K. Remember, the higher the Kelvin value, the brighter the light. We present an unbiased comparison between the two LED bulb types based on their color temperatures.


What is full spectrum light? Full spectrum is not directly visible or observable

Contributing to the confusion among consumers is the fact that the "fullness" of a light spectrum is not directly observable to the human eye. In other words, a non-full spectrum bulb and natural daylight could have the exact same emitted light color and appearance, despite having significantly different spectral properties.

Full spectrum typically refers to the completeness of a light source's spectral energy, particularly when compared to natural light sources such as natural daylight. The exact spectral composition of a light source can only be determined by specialized photometric equipment, such as a spectrometer.

In other words, as a consumer, you have no practical way to independently verify or confirm that a full spectrum bulb you purchased actually has a complete spectrum.

Spectral speaking, there are many ways to create the same light color, and this also holds true for the color of natural daylight (commonly called daylight white).

For starters, let's take a look at the light spectrum for natural daylight. You will notice that the light energy is distributed evenly across the entire visible spectrum, without any gaps, dips, or spikes.

What is very important to remember is that both natural daylight and this fluorescent lamp have the same apparent light color - daylight white. In other words, despite a significant spectral difference, the light color emitted from the fluorescent lamp is indistinguishable from daylight, to our eyes.


Excitation range and maximum

An excited fluorophore molecule emits lower-energy light than the light it absorbs. Therefore, there is always a shift along the spectrum between the color of the light absorbed by the fluorophore during excitation, and the color emitted.

A fluorescent dye absorbs light over a range of wavelengths—and every dye has a characteristic excitation range. However, some wavelengths within that range are more effective for excitation than other wavelengths. This range of wavelengths reflects the range of possible excited states that the fluorophore can achieve. So for each fluorescent dye, there is a specific wavelength—the excitation maximum—that most effectively induces fluorescence.

Typical representation of how fluorophore excitation range (bars) and fluorophore excitation maximum (stars) are displayed.


Interlude: the CIE X,Y,Z color space

It would be possible to parametrize a color by the value of the three cone cells' responses (normalized with respect to some reference white, say). Such is not, however, the usual path: the precise cone response functions to stimuli have been known only quite recently. Instead, some older, conventional functions, are of use: the CIE ( Commission Internationale de l'Éclairage )'s color matching functions, boringly called X, Y and Z. Function Y is supposed to be precisely the luminosity function. Function Z is (proportional to) the response of the short cones (this one is easy to determine, because it vanishes for sufficiently large wavelengths, certainly from 620nm on). And function X is a more or less arbitrary convention.

Typically one gets rid of the annoying dependence on luminosity, and moves to the two-dimensional color space, as follows: let x=X/(X+Y+Z), y=Y/(X+Y+Z) and z=Z/(X+Y+Z) (so that x+y+z=1 and z is redundant). Then a color is specified by its chromaticity value, (x,y), and its luminosity Y (if necessary). Often luminosity is unspecified or otherwise irrelevant, and we live fully in the two-dimensional (x,y) diagram. This is the very coordinate system we have used earlier to represent the two-dimensional chromaticity diagram (with x ranging from 0 to 1 on abscissa and y ranging from 0 to 1 on ordinate of course, everything “lives” in the lower-left triangle, because z=1−x−y has to be nonnegative also).

Here are a few sample data points:

xyApprox
Flat spectrum reference0.33330.3333
Illuminant D65 (sRGB white)0.31270.3290
Illuminant D550.33240.3474
Illuminant D50 (PCS white)0.34570.3585
Illuminant A0.44760.4074
Illuminant C0.31010.3161
sRGB red phosphor0.64000.3300
sRGB green phosphor0.30000.6000
sRGB blue phosphor0.15000.0600
Blackbody (infinite limit)0.23990.2340
Blackbody (9300K)0.28490.2932
Blackbody (6500K)0.31350.3236
Blackbody (5000K)0.34510.3516
Blackbody (3000K)0.43690.4041
Blackbody (0K limit)0.73470.2653
Monochromatic 420nm0.17140.0051
Monochromatic 460nm0.14400.0297
Monochromatic 490nm0.04540.2950
Monochromatic 520nm0.07430.8338
Monochromatic 532nm (“green laser”)0.17020.7965
Monochromatic 550nm0.30160.6923
Monochromatic 555nm (sensitivity peak)0.33740.6588
Monochromatic 570nm0.44410.5547
Monochromatic 589nm (orange sodium line)0.56930.4301
Monochromatic 590nm0.57520.4242
Monochromatic 610nm0.66580.3340
Monochromatic 635nm (“red laser”)0.71400.2859
Long cone pure stimulus (theoretical)0.75010.2499
Medium cone pure stimulus (theoretical)1.4669-.4669
Short cone pure stimulus (theoretical)0.1669-.0180


Red Shift & Blue Shift

A light source moving away from the listener (v is positive) would provide an fL that is less than fS. In the visible light spectrum, this causes a shift toward the red end of the light spectrum, so it is called a redshift. When the light source is moving toward the listener (v is negative), then fL is greater than fS. In the visible light spectrum, this causes a shift toward the high-frequency end of the light spectrum. For some reason, violet got the short end of the stick and such frequency shift is actually called a blue shift. Obviously, in the area of the electromagnetic spectrum outside of the visible light spectrum, these shifts might not actually be toward red and blue. If you're in the infrared, for example, you're ironically shifting away from red when you experience a "redshift."


Wavelength to Colour Relationship

A simple tool to convert a wavelength in nm to an RGB, hexadecimal or HSL colour.

Over the course of millions of years, the human eye has evolved to detect light in the range 380&mdash780nm, a portion of the electromagnetic spectrum known as visible light, which we perceive as colour. The particular range of wavelengths coincides with a window in the Earth's atmosphere, through which this light can travel. Higher frequency radiation, such as x-rays are absorbed by the atmosphere, as are lower frequencies, such as microwaves.

Sunlight appears white to us because it emits almost uniformly over all visible frequencies. However, a laser for example, emits only at a single very specific frequency. Helium-neon lasers emit at 632.8nm, which is a bright red. The lasers in your Blu-ray player emit at 405nm, which as the name suggests, is blue. We can begin to build up a picture of how frequency is related to colour.

A frequent way of referring to colour on computer screens is by using the RGB system. In this model, each colour is given a value for each red, green and blues components ranging from 0 to 255, giving a total value of 16.7 million possible colours. However, due to the very complex way in which the eye perceives colours, we can see colours which are outside of the gamut of the RGB scheme - there is no unique mapping that definitively converts a wavelength to a colour, and as such the above tool should been seen as more of an approximation than a rigorous resource.


What is the brightest color a light source that only emits blue frequency light can achieve? - Biology


Photos: Edison Tech Center / Planar: www.Planar.com

Using electric current through a phosphor or semiconductor
Commercial History
(1950s - Today)

Introduction & Statistics

How They Work

Inventors and Developments

To put it simply EL lamps or "high field electroluminescent" lamps use electric current directly through a phosphor to make light. Unlike most lamps, they can be shaped to be extremely flat, or in narrow wire-like shapes.
Electroluminescence or "EL" is the non-thermal conversion of electrical energy into light energy. This phenomenon is used in EL lamps, LEDs, and OLEDs. In this page we talk about EL devices which create light by exciting high energy electrons in phosphor materials like ZnS:Mn. This type of device uses "high field electroluminescence".

Advantages:
-Low wattage
-Long life
-No external circuitry required (no ballast needed to limit current, it can be plugged directly into AC power and will self-regulate power through it's own resistivity)
-Can be manufactured into flat flexible panels, narrow strings, and other small shapes
-Can be made into waterproof computer monitors which are more durable and light weight than LCDs or Plasma screens.
-Not directional like LCDs when used as a computer monitor, looks good at all angles
-EL displays can handle an impressive -60 C to 95 C temperature range, which LCD monitors cannot do

Disadvantages:
-Not practical for general lighting of large areas due to low lumen output of phosphors (so far)
-Poor lumens per watt rating, however typically the lamp is not used for high lumen output anyway
-Reduced lumen output over time, although newer technologies are better than older phosphors on this point
-Flexible flat EL sheets wear out as they get flexed, durability is being worked on
-The lamps can use significant amount of electricity: 60-600 volts
-Typical EL Needs a converter when used with DC sources such as on watches (to create higher frequency AC power, this is audible)

EL Statistics
*Lumens per watt: 2-6
*Lamp life: 2,000 - 50,000 hours
*CRI - N/A
*Color Temperature - N/A
*Available in 0.01 - 3 W

Left: EL background with LCD display,
marketed as "Indiglo" by Timex in the 1990s


A electroluminescent exit sign, easy to operate on low power and very long lamp life. Photo: Limelite

1. How it Works:

There are several variations on how EL works depending on whether you are talking about a flat panel light, rope light, DC EL technology, thin film EL display, or other complex design.

EL devices are monocarrier devices which give off light due to impact excitation of an optical center like the Mn atom. They do this by transporting high energy electrons in the host matrix (commonly ZnS).

For simplicity we will describe a simple EL lamp:

High voltage AC power passes through a thin layer of phosphor or semiconductor and this causes emission of light. Two layers of solid material (one being transparent) act as electrodes and a powder phosphor or semiconductor in between glows when electrons pass through it from one electrode to another. Light escapes the device on one side thanks to the development of transparent conductors like indium tin.

Thick phosphor powder EL Lamps are used in most simple lamps used for illumination including the night lights and exit/safety signs. The graphic below shows thick phosphor lamps.

Thin Film and Thick Dielectric EL (TFEL, TDEL) : this technology is used in a variety of applications, EL displays (ELDs) are the most common use. A display is not a "lamp" in the traditional sense, however we cover it here due to its importance in the development of EL. TFEL and TDEL often use rare earth materials such as Er, Tm, In, and more.

TFEL - Thin Film Electroluminescent Devices
TFEL emerged in the 1950s and it different in that it contains thinner active layers and a different construction. TFEL was an improvement over thick powder construction, it allows for small devices and precise control of pixels on a display. It was a challenge to develop ways to deposit/grow thin polycrystalline films onto a substrate (the supporting material) however many processes have been developed to allow TFEL technologies to expand. Below we highlight the basic construction of a TFEL device.
Note:
-TFEL has a maximum of 6 lumens per Watt as of 2012.
-TFEL typically requires 1.5 Megavolts per centimeter to cause the active layer to make light

Video of TFEL construction:

How TFEL works:
TFEL has a phosphor layer that emits light when a big enough electrical field is applied. This thin film of a phosphor requires such a high level of energy that there is a potential for a damaging short circuit through imperfections in the phosphor. Insulating layers are used between the electrode and the phosphor on both sides to limit current and make the TFEL work properly.

TFEL devices behave like 3 capacitors in series: voltage rises and a breakdown voltage is reached where current flows through the semiconductive layer (the phosphor) which excites the phosphor and makes light. The insulator layers act as capacitors, with voltage building and breaking through.

Before you see how the EL device works you may want to review how a capacitor works in this video:


Thin film EL uses a process of epitaxy to grow crystals on top of a substrate. This process allows one to create a "film" or ultra-thin layer of material (measured in nanometers (nm)) on glass or other flat surface (this surface provides structure and is called the 'substrate'). TFEL epitaxy creates layers about 500 nanometers thick, although the size varies depending on the product. Later on TDEL (thick dielectric EL) was developed to produce a product with higher luminosity than TFEL. TDEL uses a structure where electrodes are separated from the thicker phosphor by a thin insulative layer. Both TFEL and TDEL use epitaxy, there are many forms of epitaxy from MBE (molecular beam) to ALE (Atomic Layer Epitaxy) (which was renamed ALD (Atomic Layer Deposition)). Understanding epitaxy requires a bit of time, we recommend online lectures and web sites for this area. Read more about ALD from Tuomo Suntola here (PDF).

Transparent and Non-transparent EL displays

One way to build a non-transparent TFEL display is to use two layers for plastic film or glass, one is coated with indium tin oxide (ITO) or other semiconductor while the other flat surface has a reflective material. Light will be produced in the 'active' layer of phosphors (ZnS Mn for example). Light emitted in the wrong direction will be reflected off the back plate and go through the opposite side which has the transparent semiconductor, this way you achieve a higher luminosity. With many individually controlled units and a controlling computer you can turn the unit on or off, collectively this will make a display screen. In a multi-color display filters applied over top of the units can control whether the unit emits red, yellow or green light. Blue has not been developed yet, and it is because of this EL displays cannot currently compete with LCD technology for full-color consumer displays.

Transparent EL displays two layers of transparent conducting films (TCFs) as the electrodes with the phosphor in between. Since they do not have a reflective backing they do not currently produce the same level of brightness as standard EL displays. Despite this the display has some very interesting and unique applications which have not become widespread yet.

Transparent conducting films (TCFs) include indium-tin oxide (ITO) and fluorine doped tin or zinc oxide (FTO)(FZO). ITO is also used in the thin-film solar industry. Carbon nanotube technology is an organic conducting film which could replace expensive rare earth materials like indium. Poly(3,4-ethylenedioxythiophene) PEDOT films and other polymer films also have potential to replace ITO. Making newer cheaper materials is important for seeing the growth of EL displays and lights in the daily life of consumers.

This type of lamp makes light as electrons radioactively combined in holes of a semiconductor. Understanding how semiconductors work on a molecular level requires a long description or entire lecture. The Indian Institute of Technology Madras has a multi-video lecture starting with a 59 minute video on solid state materials.

TDEL:
TDEL or thick film dielectric EL technology is known for providing a solution to the blue problem. It provides the only full color RGB display technology available at this time.

Thick film dielectric displays have proven to be effective: they have a good brightness (luminosity) and have a decent efficiency. iFire Group and TDK Corporation currently hold the patents for this technology. The phosphor in TDEL is 10K - 20K nanometers thick. Some TDEL like that used in displays uses two layers of phosphors. The bottom thick layer is resistant to dielectric breakdown, so it can transport a higher current and make a brighter light. Above the thick dielectric layer is colored phosphors of ZnMgS:Mn (green) and BaAl2S4:Eu (blue). With this system RGB can be created.

2. Inventors and Developments:

Electroluminescence was used as early as 1936 by scientist Georges Destriau. It wasn't until the 1950s when companies began developing the technology to be used for practical applications.

1936 - Georges Destriau , who was an associate of Marie Curie in her lab in Paris began studying electroluminescence. He coined the term as he worked with ZnS powders.
Paris, France

1958 - Elmer Fridrich while working for General Electric developed EL lamps, some of which where quite sophisticated in design. Fridrich also became famous for inventing the halogen lamp and advancing fluorescent lamp technology. He was a key member of engineering teams at Nela Park, Ohio and Schenectady, New York.
Photo: General Electric

1958 - Nataliya Andreeva Vlasenko and A. Popkov : Developed the first TFEL prototype and worked on methods for boosting luminosity. They pioneered early work on DC EL lamps as well.
Kiev, Ukraine

1968 - Aron Vecht develops DC EL technology for lamps and watches. London, UK
Photo: University of Greenwich

1974 - Tuomo Suntola develops ALE Epitaxy for Thin film electroluminescent (TFEL) technologies. This method of depositing thin semiconducting films on a substrate has become a basis for TFEL production. Thin polycrystalline films are about 500 nanometers thick. Thin films allow for a more uses of EL than clumsy thick phosphor powders.
Lohja, Finland
Photo: Tuomo Suntola

1970s - Hiroshi Kobayashi worked over 30 years on inorganic electroluminescent devices with late Professor Shosaku Tanaka . His work helped with the commercialization of inorganic EL displays in Japanese industry. A great deal of work was done at Tottori University. He retired in 2003 and now lives in Tokyo.
Tottori Prefecture / Tokyo, Japan
Photo: Hiroshi Kobayashi

1974 - Toshio Inoguchi develops the first practical ELD (electroluminescent display) at Sharp Corporation. He uses TFEL to make this possible. His displays have long life and are brighter in luminosity. His work set the stage for later advancements and kept Sharp at the leading edge for the next few decades. The displays were used first as displays for medical instruments. The displays were monochromatic, but a better option than CRTs.
Osaka, Japan
Photo: Toshio Inoguchi and Sharp Corporation

1980s - Christopher N. King and team* develop advanced EL displays which use thin film technology. The team had started at Tektronix and launched spin-off Planar Systems in 1983. The new displays increase the number of available colors as time goes on. Increasing luminosity and contrast to compete with LCDs became important in the 1990s and 2000s. Since the 90s the engineers at Planar have improved the EL display, they have achieved better luminosity, contrast and efficiency. *Jim Hurd, John Laney, Eric R. Dickey (ICEBrite)
Beaverton, Oregon
Photo: Chris King

1990s - Xingwei Wu develops TDEL technology. Thick Dielectric EL displays achieve blues bright enough to be used in full color displays. TDEL is brighter than TFEL, and uses "color by blue" method to achieve good RGB. TDEL is the first full color capable EL technology. Dr. Xingwei Wu is the primary engineer at iFire Technology.

Oakville, Ontario, Canada

Photo: Xingwei Wu. iFire Technology Ltd.

2016 - You - Choose a career in engineering and be the next pioneer! LEARN MORE

2000s - EL lamps become more affordable to the average consumer and are used in decorative clothing and thin film application on various products. As a lamp for general illumination EL technology is not preferred due to limited maximum lumen production combined with low efficiency compared with LEDs. The unique spatial aspect of the EL lamp (flat and flexible) allows it to maintain a market niche.

EL displays have come a long way since 1980 however a better blue phosphor which can be used in displays is still needed. Developing a high-luminosity, high-efficiency blue would allow a red-green-blue combination that would allow the EL display to better compete with LCD.

Further reading in more detail:
A History of Electroluminescent Displays by Jeffrey A. Hart, Stefanie Ann Lenway, Thomas Murtha. 1999


Bohr's Model

In 1913, a Danish physicist, Niels Bohr (1885&ndash1962 Nobel Prize in Physics, 1922), proposed a theoretical model for the hydrogen atom that explained its emission spectrum. Bohr&rsquos model required only one assumption: The electron moves around the nucleus in circular orbits that can have only certain allowed radii. Rutherford&rsquos earlier model of the atom had also assumed that electrons moved in circular orbits around the nucleus and that the atom was held together by the electrostatic attraction between the positively charged nucleus and the negatively charged electron. Although we now know that the assumption of circular orbits was incorrect, Bohr&rsquos insight was to propose that the electron could occupy only certain regions of space.

Using classical physics, Niels Bohr showed that the energy of an electron in a particular orbit is given by

where ( Re ) is the Rydberg constant, h is Planck&rsquos constant, c is the speed of light, and n is a positive integer corresponding to the number assigned to the orbit, with n = 1 corresponding to the orbit closest to the nucleus. In this model n = &infin corresponds to the level where the energy holding the electron and the nucleus together is zero. In that level, the electron is unbound from the nucleus and the atom has been separated into a negatively charged (the electron) and a positively charged (the nucleus) ion. In this state the radius of the orbit is also infinite. The atom has been ionized.

Figure 7.3.2 The Bohr Model of the Hydrogen Atom (a) The distance of the orbit from the nucleus increases with increasing n. (b) The energy of the orbit becomes increasingly less negative with increasing n.

During the Nazi occupation of Denmark in World War II, Bohr escaped to the United States, where he became associated with the Atomic Energy Project.

In his final years, he devoted himself to the peaceful application of atomic physics and to resolving political problems arising from the development of atomic weapons.

As n decreases, the energy holding the electron and the nucleus together becomes increasingly negative, the radius of the orbit shrinks and more energy is needed to ionize the atom. The orbit with n = 1 is the lowest lying and most tightly bound. The negative sign in Equation 7.3.3 indicates that the electron-nucleus pair is more tightly bound when they are near each other than when they are far apart. Because a hydrogen atom with its one electron in this orbit has the lowest possible energy, this is the ground state (the most stable arrangement of electrons for an element or a compound), the most stable arrangement for a hydrogen atom. As n increases, the radius of the orbit increases the electron is farther from the proton, which results in a less stable arrangement with higher potential energy (Figure 2.10). A hydrogen atom with an electron in an orbit with n > 1 is therefore in an excited state. Any arrangement of electrons that is higher in energy than the ground state.: its energy is higher than the energy of the ground state. When an atom in an excited state undergoes a transition to the ground state in a process called decay, it loses energy by emitting a photon whose energy corresponds to the difference in energy between the two states (Figure 7.3.1 ).

Figure 7.3.3 The Emission of Light by a Hydrogen Atom in an Excited State. (a) Light is emitted when the electron undergoes a transition from an orbit with a higher value of n (at a higher energy) to an orbit with a lower value of n (at lower energy). (b) The Balmer series of emission lines is due to transitions from orbits with n &ge 3 to the orbit with n = 2. The differences in energy between these levels corresponds to light in the visible portion of the electromagnetic spectrum.

So the difference in energy (&DeltaE) between any two orbits or energy levels is given by ( Delta E=E_>-E_> ) where n1 is the final orbit and n2 the initial orbit. Substituting from Bohr&rsquos equation (Equation 7.3.3) for each energy value gives

If n2 > n1, the transition is from a higher energy state (larger-radius orbit) to a lower energy state (smaller-radius orbit), as shown by the dashed arrow in part (a) in Figure 7.3.3. Substituting hc/&lambda for &DeltaE gives

Canceling hc on both sides gives

Except for the negative sign, this is the same equation that Rydberg obtained experimentally. The negative sign in Equation 7.3.5 and Equation 7.3.6 indicates that energy is released as the electron moves from orbit n2 to orbit n1 because orbit n2 is at a higher energy than orbit n1. Bohr calculated the value of (Re) from fundamental constants such as the charge and mass of the electron and Planck's constant and obtained a value of 1.0974 × 10 7 m &minus1 , the same number Rydberg had obtained by analyzing the emission spectra.

We can now understand the physical basis for the Balmer series of lines in the emission spectrum of hydrogen (part (b) in Figure 2.9 ). As shown in part (b) in Figure 7.3.3 , the lines in this series correspond to transitions from higher-energy orbits (n > 2) to the second orbit (n = 2). Thus the hydrogen atoms in the sample have absorbed energy from the electrical discharge and decayed from a higher-energy excited state (n > 2) to a lower-energy state (n = 2) by emitting a photon of electromagnetic radiation whose energy corresponds exactly to the difference in energy between the two states (part (a) in Figure 7.3.3 ). The n = 3 to n = 2 transition gives rise to the line at 656 nm (red), the n = 4 to n = 2 transition to the line at 486 nm (green), the n = 5 to n = 2 transition to the line at 434 nm (blue), and the n = 6 to n = 2 transition to the line at 410 nm (violet). Because a sample of hydrogen contains a large number of atoms, the intensity of the various lines in a line spectrum depends on the number of atoms in each excited state. At the temperature in the gas discharge tube, more atoms are in the n = 3 than the n &ge 4 levels. Consequently, the n = 3 to n = 2 transition is the most intense line, producing the characteristic red color of a hydrogen discharge (part (a) in Figure 7.3.1 ). Other families of lines are produced by transitions from excited states with n > 1 to the orbit with n = 1 or to orbits with n &ge 3. These transitions are shown schematically in Figure 7.3.4

Figure 7.3.4 Electron Transitions Responsible for the Various Series of Lines Observed in the Emission Spectrum of Hydrogen. The Lyman series of lines is due to transitions from higher-energy orbits to the lowest-energy orbit (n = 1) these transitions release a great deal of energy, corresponding to radiation in the ultraviolet portion of the electromagnetic spectrum. The Paschen, Brackett, and Pfund series of lines are due to transitions from higher-energy orbits to orbits with n = 3, 4, and 5, respectively these transitions release substantially less energy, corresponding to infrared radiation. (Orbits are not drawn to scale.)

In contemporary applications, electron transitions are used in timekeeping that needs to be exact. Telecommunications systems, such as cell phones, depend on timing signals that are accurate to within a millionth of a second per day, as are the devices that control the US power grid. Global positioning system (GPS) signals must be accurate to within a billionth of a second per day, which is equivalent to gaining or losing no more than one second in 1,400,000 years. Quantifying time requires finding an event with an interval that repeats on a regular basis. To achieve the accuracy required for modern purposes, physicists have turned to the atom. The current standard used to calibrate clocks is the cesium atom. Supercooled cesium atoms are placed in a vacuum chamber and bombarded with microwaves whose frequencies are carefully controlled. When the frequency is exactly right, the atoms absorb enough energy to undergo an electronic transition to a higher-energy state. Decay to a lower-energy state emits radiation. The microwave frequency is continually adjusted, serving as the clock&rsquos pendulum. In 1967, the second was defined as the duration of 9,192,631,770 oscillations of the resonant frequency of a cesium atom, called the cesium clock. Research is currently under way to develop the next generation of atomic clocks that promise to be even more accurate. Such devices would allow scientists to monitor vanishingly faint electromagnetic signals produced by nerve pathways in the brain and geologists to measure variations in gravitational fields, which cause fluctuations in time, that would aid in the discovery of oil or minerals.

Example 7.3.1: The Lyman Series

The so-called Lyman series of lines in the emission spectrum of hydrogen corresponds to transitions from various excited states to the n = 1 orbit. Calculate the wavelength of the lowest-energy line in the Lyman series to three significant figures. In what region of the electromagnetic spectrum does it occur?

Given: lowest-energy orbit in the Lyman series

Asked for: wavelength of the lowest-energy Lyman line and corresponding region of the spectrum

  1. Substitute the appropriate values into Equation 7.3.2 (the Rydberg equation) and solve for (lambda).
  2. Locate the region of the electromagnetic spectrum corresponding to the calculated wavelength.

We can use the Rydberg equation to calculate the wavelength:

A For the Lyman series, n1 = 1. The lowest-energy line is due to a transition from the n = 2 to n = 1 orbit because they are the closest in energy.

It turns out that spectroscopists (the people who study spectroscopy) use cm -1 rather than m -1 as a common unit. Wavelength is inversely proportional to energy but frequency is directly proportional as shown by Planck's formula, E=h( u ).

Spectroscopists often talk about energy and frequency as equivalent. The cm -1 unit is particularly convenient. The infrared range is roughly 200 - 5,000 cm -1 , the visible from 11,000 to 25.000 cm -1 and the UV between 25,000 and 100,000 cm -1 . The units of cm -1 are called wavenumbers, although people often verbalize it as inverse centimeters. We can convert the answer in part A to cm -1 .

[lambda = 1.215 imes 10^<&minus7> m = 122 nm ]

This emission line is called Lyman alpha. It is the strongest atomic emission line from the sun and drives the chemistry of the upper atmosphere of all the planets producing ions by stripping electrons from atoms and molecules. It is completely absorbed by oxygen in the upper stratosphere, dissociating O2 molecules to O atoms which react with other O2 molecules to form stratospheric ozone

B This wavelength is in the ultraviolet region of the spectrum.

Exercise 7.3.1: The Pfund Series

The Pfund series of lines in the emission spectrum of hydrogen corresponds to transitions from higher excited states to the n = 5 orbit. Calculate the wavelength of the second line in the Pfund series to three significant figures. In which region of the spectrum does it lie?

Answer: 4.65 × 10 3 nm infrared

Bohr&rsquos model of the hydrogen atom gave an exact explanation for its observed emission spectrum. The following are his key contributions to our understanding of atomic structure:

  • Electrons can occupy only certain regions of space, called orbits.
  • Orbits closer to the nucleus are lower in energy.
  • Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic spectra.

Unfortunately, Bohr could not explain why the electron should be restricted to particular orbits. Also, despite a great deal of tinkering, such as assuming that orbits could be ellipses rather than circles, his model could not quantitatively explain the emission spectra of any element other than hydrogen (Figure 7.3.5). In fact, Bohr&rsquos model worked only for species that contained just one electron: H, He + , Li 2 + , and so forth. Scientists needed a fundamental change in their way of thinking about the electronic structure of atoms to advance beyond the Bohr model.

Figure 7.3.5 The Emission Spectra of Elements Compared with Hydrogen. These images show (a) hydrogen gas, which is atomized to hydrogen atoms in the discharge tube (b) neon and (c) mercury. The strongest lines in the hydrogen spectrum are in the far UV Lyman series starting at 124 nm and below. The strongest lines in the mercury spectrum are at 181 and 254 nm, also in the UV. These are not shown.

Thus far we have explicitly considered only the emission of light by atoms in excited states, which produces an emission spectrum (a spectrum produced by the emission of light by atoms in excited states). The converse, absorption of light by ground-state atoms to produce an excited state, can also occur, producing an absorption spectrum (a spectrum produced by the absorption of light by ground-state atoms). Because each element has characteristic emission and absorption spectra, scientists can use such spectra to analyze the composition of matter.

When an atom emits light, it decays to a lower energy state when an atom absorbs light, it is excited to a higher energy state.


3 Answers 3

The 100 W light bulb dissipates more energy per second (1 watt = 1 joule per second) than the 20 W light bulb, and consequently the light emanating from the 100 W bulb carries more energy than the light emanating from the 20 W bulb.

In the picture of light as an electromagnetic wave, the energy carried by the light is proportional to the square of the wave's amplitude. The technical term for this energy is "Poynting flux". (In fact we usually take the time-average over one period of oscillation as the definition of the energy in the wave.) In this model, the photo-receptors in your eye are oscillators. What is oscillating? Electric charge. Charges are accelerated in response to the electric field of the light: the greater the electric field (or amplitude), the greater the amplitude of the oscillation, and the greater the electric currents in your eye (and the greater the brightness).

In the picture of light as a particle (a photon), each particle carries with it an amount of energy proportional to its frequency: $E=h u$, where $h$ is Planck's constant, and $ u$ is the frequency of light. The energy flux is then the energy per photon multiplied by the flux of photons (# of photons per unit area per second). So the 100 W bulb emits more photons per second than the 20 W bulb. In this model, the photoreceptors in your eye undergo chemical reactions as a result of absorbing photons. The more photons absorbed per second, the brighter the light appears.

Brightness is just the number of photons per second hitting your eye - all the other properties of the light are the same.

edit: perceived brightness is the number of 'detected' photons hitting your eye per second!

Different wavelengths of light correspond to different colours. 555nm means light with a wavelength of 555 nano-meters (billions of a meter), this is roughly green light. So all this says is that you eye is most sensitive to green light and so a given number of green photons/second will appear brighter than the same number of red photons. You can see this with laser pointers, for the same power small pointers - green ones look much brighter than red.

I am nowhere near as expert as a professional, but I have a private passion for this field.

Dim and bright are perceptual terms. There are many dimensions. I will start with a simple idea and build out.

Consider that you are adapted to a monochromatic light of around 533 nanometer wavelength bathing the room such that the light from the brightest object generates approximately 1e7 photons per second on a foveal L cone of 1 micron face diameter. This is considered a well-lit but not stressful scene. The photoreceptor opsins bleach at a rate of approximately 5e3 opsins per second. The light feels neither dim nor bright because you are adapted.

If you increase the source photon rate by a factor of 10, that same photoreceptor is now bleaching at a rate of 5e4 opsins per second. This feels brighter. But over time, the photoreceptor undergoes phagocytosis, decreasing its length by 90% changing the opsin bleach rate back to 5e3 opsins per second, so you now experience this as neither dim nor bright.

If you decrease the wavelength to 430 nanometers, the bleach rate of the L cone decreases by 90% and one would think this would appear dim. However, the S cone bleach rate reaches its maximum and S cones have a stronger effect on perceived brightness than L cones, so without increasing the photon bleach rate, the light now appears to have gotten brighter.

This is the principal reason amber sunglasses make the world seem brighter and more colorful. By suppressing the short wavelength photons from reaching the eye, the adaptation of the L and M cone bias favors a greater linear range. This makes colors more discriminable and is another dimension of brightness.

I will leave these three dimensions for further discussion and, if requested, I will dive deeper yet into the wonderful world of retinal adaptation to various light regimes.


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