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Chromophores

Page history last edited by Courtney Talicska 9 years, 4 months ago

 

 

 

 

Introduction

 

Chromophores are the portions of molecules that cause them to appear colored. Molecules appear colored when electromagnetic energy in the visible region - corresponding to the energy difference between the ground state and excited state - is absorbed by a molecule1. Thus, photons of light that have the correct amount of energy can excite electrons to higher energy levels when they are absorbed by a chromophore. Molecules composed of only single bonds generally have energies of excitations that are so large they require a photon of extremely high energy to excite electrons. Such photons are found outside the visible region of the electromagnetic spectrum and are not seen by human eyes. Molecules with many conjugated bonds (which can be modeled by the infamous particle-in-a-box model) and metal complexes with d-orbitals that have split energy levels will have spacings between energy levels that correspond to the visible portion of the electromagnetic spectrum. It is precisely due to the fact that such molecules absorb photons of energy corresponding to visible light that they appear colored.

 

There are two types of chromophores: conjugated pi system and metal complex chromophores. A conjugated pi-system chromophore is recognized a molecule with a series of alternating double and single bonds, while a metal complex chromophore involves the splitting of d-orbitals by binding a transition metal to ligands2. Each type of chromophore is discussed in more detail below.

 

Conjugated Chromophores

 

Conjugated chromophores, as mentioned earlier, are formed by alternating single and double bonds in a molecule. This alternating bond pattern, known as conjugation, de-localizes electron density throughout the conjugated portion of the molecule. This causes a conjugated chromophore to have a lower energy configuration than a molecule consisting of only single bonds. Thus, conjugated chromophores require electromagnetic radiation of a lower energy to excite their electrons, and so they absorb photons of light that have an energy corresponding to the visible region of the electromagnetic spectrum. In this case, the visible wavelengths that are not absorbed by the chromophore are reflected off the compound towards your eyes and the compound appears to be colored! Some examples of molecules containing conjugated chromophores are the organic pigments that make tomatoes appear red and carrots orange, lycopene and β-carotene.

 


Molecular structures of the organic pigments lycopene and β-Carotene. The conjugated chromophore portion of each molecule is highlighted in blue.

 

Metal Complex Chromophores

 

Metal complex chromophores form when the d-orbital of a molecule is split by the binding of a transition metal to an ion or molecule. Transition metals are commonly metal complex chromophores because they have an unfilled d-level atomic orbital and bind to ions or molecules in an attempt to 'fill' this orbital. When such binding occurs, the newly filled d-orbital contains some electrons that are at a slightly higher energy than the other electrons of the d-orbital. A pictorial representation of this concept is given in the diagram below. The splitting of the d-orbital causes the molecule to appear colored because the energy spacing between the lower-energy d-orbital and the higher-energy d-orbital is small enough that a photon of light of an energy corresponding to the visible portion of the electromagnetic spectrum can excite an electron from the lower to the upper energy d-orbital3. Thus, metal complex chromophores absorb light of an energy that corresponds to visible light, and the light not absorbed by the metal complex chromophore is reflected off the compound and interpreted by your eye as color! Some common examples of metal complex chromphores include hemoglobin and colorful minerals like amethyst and malachite.

 

 

In the unexact point of view, all of the d-orbitals appear to be in the same energy level. In reality two of the d-orbitals are slightly higher in energy than the other three. This causes a splitting of the d-orbitals which produces color in metal complexes.

 

 

Quantum Mechanical Perspective

 

From all we have learned about quantum mechanics, it should come as a very little surprise that quantum mechanics govern the physical principles of our everyday lives. Chromophores are no exception to this rule. Chromophores cause objects to appear colored because they absorb electromagnetic energy corresponding to the visible region of the electromagnetic spectrum. This is due to the fact that molecules that contain chromophores have energy levels that are more closely spaced than molecules that do not contain chromophores4. This, in turn, causes the molecule to absorb electromagnetic radiation of a lower energy and frequency corresponding to certain colors of the visible spectrum.

 

Quantum mechanically it must be noted that there is a fundamental limit on the energy that objects can absorb. Rather than absorbing energy continuously, objects are limited to absorbing discrete quanta of energy in the form of photons4. Depending on its constituent atoms and number of conjugated bonds, each chromophore will absorb photons of differing energy levels as noted above. Of major focus to us is the effect of conjugation on the absorption of these photons. In other words, we want to know how conjugation affects the energy a chromophore will absorb to cause the electronic transitions described in the previous section.

 

This is most easily illustrated by the common model used in quantum mechanics for conjugation: the particle-in-a-box model! As you may already know, the particle-in-a-box model utilizes a box of the length of the conjugated chain present in the chromophore. Depending on the length of the conjugated chain and the extent of conjugation in the chromophore, the energy levels of the particle-in-a-box model change as illustrated below.

 

Molecules with only single (sigma) bonds have orbitals that are farther apart in energy than molecules with conjugated double bonds.

 

Clearly, molecules with conjugated double bounds (those that contain chromophores) have energy levels that are more closely spaced than molecules with only single (sigma) bonds. The same concept of closely space energy levels applies to metal complex chromophores, in which the energy differences between the lower and upper energy levels of the split d-orbitals is very small3 . This provides rationale for the conclusion that molecules containing chromophores will have electrons that require less energy to be excited from the ground state to the excited states and thus will absorb photons of a lower energy and frequency corresponding to a specific energy of the visible portion of the electromagnetic spectrum.

 

 

Energy Absorption

 

Electrons can absorb energy from light to be promoted to a higher energy level. Eventually an excited electron loses the energy gained from the absorption of a photon of light by emitting electromagnetic radiation, causing it to go back to a more stable energy level.

 

 

 

An electron absorbs energy from a photon of light and is excited to a higher energy level. Eventually, the electron loses this energy (emits a photon) and goes back to the ground state.

 

Electronic spectroscopy is a technique that is used to study the electronic structure of atoms and molecules. It uses absorption to provide meaningful data about the wavelengths of light a molecule does and does not absorb5. Electronic spectroscopy applies the broad method of absorption spectroscopy to the ultra-violet and visible range of the electromagnetic spectrum. Absorption spectroscopy in general relies on the ability of a molecule to absorb electromagnetic radiation of a certain energy to study the structure of complex molecules. The experimental set up for electronic spectroscopy is simple, a source of electromagnetic energy is passed through a sample towards a detector. Only the electromagnetic radiation that is not absorbed by the molecule reaches the detector and is recorded6 . From this, the wavelength(s) of light that did not reach the detector (the light that the molecule absorbed) is determined and this data is used to classify the molecule. Absorption spectroscopy can be applied to any region of the electromagnetic spectrum: ultra-violet, visible, x-ray, and more depending on the type of transition being studied.

 

The production of color on an object occurs when an object is illuminated by light. Some of the photons of light hitting the object are absorbed. The photons of light absorbed by the molecule correspond to the wavelength of electromagnetic radiation required for electrons to be excited to higher energy levels. The complement to the color of light absorbed is reflected back to our eyes, which then detect the light as color. The color of a metal complex is produced when a photon of light excites an electron in a split d-orbital to the higher energy level of the split d-orbital.  A certain wavelength of light is absorbed by the metal complex. The color of the object as it appears to us is the complementary color of the wavelength of light that is absorbed by the metal complex. This is why many transition metals that are bound to ions or molecules appear intensely colored.

 

 

Electron Transition

 

When light of the right energy hits a chromophore, the energy from the light can be used to promote an electron from a bonding orbital into an empty antibonding orbital.  Since bigger jumps require more energy, it can only be caused by light that transfers more energy, which correlates to means light with a higher frequency and therefore a lower wavelength.  The distance that the electron jumps directly corrolates with the wavelength of light needed to excite the electron that distance.

An absorption spectrometer works sends out waves that vary from UV to infrared wavelengths, and only the jumps allowed from that range of wavelengths occur.  Electrons that are excited in this way can only go from pi orbitals to pi antibonding orbitals; from nonbonding orbitals to pi antibonding orbitals; and to a lesser extent, nonbonding orbitals to sigma antibonding orbitals.  

 

When light hits the chromophore area of the molecule, the energy from the light can be used to promote and electron from a bonding orbital into an empty antibonding orbital7.  Since bigger jumps require more energy, it can only be caused by light that transfers more energy, which corresponds to light with a higher frequency and therefore a lower wavelength.  

The point of highest absorption by the molecule is given by Electron Absorption Spectroscopy, and from this we can tell what colors are given off.  Electron Absorption Spectroscopy is a process in which light waves ranging from UV to infrared are shined on a material, to determine which wavelengths are the most heavily absorbed8.  For example, beta carotene has its highest absorption from 440 to 520 nanometers, and its peak is at 470 nanometers9.  This corresponds to the blue-cyan region shown on the color wheel below, which means beta carotene absorbs blue-cyan light, meaning that the color it gives off is the complementary color to blue-cyan, which is orange. 

 

The colors corresponding to specific ranges of wavelengths in the electromagnetic spectrum.

 

The basic color wheel, the complement of a given color is the color located directly across from it on the color wheel.

 

Colors opposite each other on the wheel are complementary, for example, if a chromophore were to absorb blue light, it would give off yellow light. And in fact, if you mix two complementary colors, they would give you white light.  Note that this is not like paint, as mixing blue and yellow would certainly not give you white paint!

 

 

So...Why Should I Care?

 

Chromophores are present in organic molecules like paints, dyes, and food colorings. Chromophores are part of the human body in the red coloring of hemoglobin and inside of our eyes, where they detect light. Some chromophores even absorb and emit different wavelengths of light based on different levels of pH in their surrounding environment. Such variation can be based on whether the molecule is protonated or not. In short, chromophores are present all around the world and we encounter them every day. In fact, without chromophores many molecules would not appear colored. Imagine living in a world without the vibrant color changes of leaves in the fall or the bright purple of an amethyst crystal. A world without chromophores would be a dull, colorless mystery!

 

 

 

References

1) "Chromophores Summary | BookRags.com." BookRags.com | Study Guides, Lesson Plans, Book Summaries and More. World Of Chemistry, 2011. Web. 3 Dec. 2011. <http://www.bookrags.com/research/chromophores-woc/>.

2) Schwartz, Benjamin. "Conjugated Polymers. What Makes a Chromophore?" News and Views. Nature Publishing Group, 2011. Web. 2 Dec. 2011. <http://www.nature.com/nmat/journal/v7/n6/full/nmat2191.html>.

3) "What Is a Chromophore?" InnovateUs Inc., 2008. Web. 2 Dec. 2011. <http://www.innovateus.net/science/what-chromophore>. 

DePalma, Angela. "UV-VIS Spectroscopy: It's All in the Chromophores." Lab Manager Magazine®. 7 May 2010. Web. 2 Dec. 2011. <http://www.labmanager.com/?articles.view/articleNo/3659/article/UV-VIS-Spectroscopy--It-s-all-in-the-Chromophores>.

4) Lloyd, Seth. "A Bit of Quantum Hanky Panky." Digital Publishing Company. Massachusetts Institute of Technology. Web. 3 Dec. 2011. <http://mag.digitalpc.co.uk/Olive/ODE/physicsworld/LandingPage/LandingPage.aspx?href=UEhZU1dvZGUvMjAxMS8wMS8wMQ..>.

5) "Electronic Spectroscopy." Massachusetts Institute of Technology. Web. 2 Dec. 2011. <http://web.mit.edu/5.33/www/lec/spec6.pdf>.

6) "Molecular Expressions: Science, Optics and You - Electron Absorption and Emission: Interactive Java Tutorial." Molecular Expressions: Images from the Microscope. 15 June 2006. Web. 4 Dec. 2011. <http://micro.magnet.fsu.edu/primer/java/scienceopticsu/exciteemit/index.html>.

7) Casado, Juan. "Vibrational and Quantum Study of Non-Linear Optical Chromophores." 2004. Web. 3 Dec. 2011. <http://www.biblioteca.uma.es/bbldoc/articulos/16494532.pdf>.

8) Nighswander-Rempel, Stephen. "Quantum Yield Calculations for Strongly Absorbing Chromophores." Centre for Biophotonics and Laser Science. University of Queensland, 2010. Web. 2 Dec. 2011. <http://arxiv.org/ftp/physics/papers/0601/0601147.pdf>.

9) "Beta-carotene." Laser Photomedicine and Biomedical Optics at the Oregon Medical Laser Center. Web. 2 Dec. 2011. <http://omlc.ogi.edu/spectra/PhotochemCAD/html/beta-carotene.html>.

 

Pictures:

Color Wheel: http://sbaratz.home.mindspring.com/project/2.htm

Colored Beakers: https://sites.google.com/a/ucdavis.edu/straughnchembio/home

Energy Absorption: http://micro.magnet.fsu.edu/primer/java/scienceopticsu/exciteemit/index.html

Lycopene and B-Carotene: http://www.orthomolecular.org/library/jom/2000/articles/2000-v15n02-p103.shtml

Wavelengths and Color: http://earthsci.org/education/teacher/basicgeol/change/change.html 

 

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