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1、Emission SpectroscopyAccording to the Bohr atomic model, electrons orbit the nucleus within specific energy levels. These levels are defined by unique amounts of energy. Electrons that have the lowest energy are found in the levels closest to the nucleus. Electrons of higher energy are located furth
2、er away.If an electron absorbs enough energy to bridge the gap between energy levels, the electron may jump to a higher level. Since this change results in a vacant lower orbital, the configuration is unstable. The excited!” electron releases its newly acquired energy and falls back to its initial o
3、r ground state. Often, the excited electrons get enough energy to make several energy level transitions. When these electrons return to their ground state, several distinct energy emissions occur. The energy that electrons absorb is often of a thermal or electrical nature, and the energy that electr
4、ons emit when returning to ground state is electromagnetic radiation. Sometimes this electromagnetic radiation can be picked up by our eyeballs as it is energy in the form of visible light. In 1900, Max Planck studied visible emissions from hot glowing solids. He proposed that light was emitted in p
5、ackets of energy called quanta, and that the energy of each packet was proportional to the frequency of the light wave. According to Einstein and Planck, the energy of the packet could be expressed as the product of the frequency (n) of emitted light and Planks constant, h, 6.626 x 10-34 Js: E = hn.
6、 Wavelength is related to frequency in the equation, c = nl, where l stands for wavelength. l is the lower case Greek letter for L, and is called lambda. Frequency and wavelength have an inverse relationship, which equals c, the speed of light, 3.00 x 108 m/s. If we know the wavelength of the light,
7、 then we can calculate the frequency of the light. Once we have calculated the frequency, then we can use Planks constant to calculate how much energy, in Joules, is in the wavelength. For example, red colors tend to have wavelengths around 700 nm. So, if we do the math:c = flc=flso 3.00 x 108 mx109
8、 nm=700 nm s1 m4.29 x 1014 s-1 = the frequency of the red lightNow using E= hnwe get E = (6.626 x 10-34 Js)(4.29 x 1014 s-1)= 2.69 x 10 19 JAlthough this number really has no meaning to us, if we were to calculate various wavelengths of light, we would find out that blue light which has a shorter wa
9、velength, roughly 400 nm, has more energy. There are two wavelengths that you might be familiar with, 680 nm and 700 nm. During photosynthesis, photocenters are able to absorb light that has these two wavelengths. The chloroplast is able to take the light energy and convert it to chemical energy. No
10、t all of the energy emitted when an electron drops from the excited state to the ground state can be perceived as visible light. Only wavelengths ranging from 400 nm to 700 nm are visible to the human eye. Other wavelengths exist, but we are unable to detect them with our eyeballs. Visible light is
11、just one part of the electromagnetic spectrum. All energy that exists in the form of waves falls somewhere on the electromagnetic spectrum. At one end are gamma rays which have a very short wavelength and lots of energy, whereas at the other end of the spectrum are radio waves which have a long wave
12、length and low energy. If white light, such as daylight, passes through a prism or diffraction grating, its component wavelengths are bent at different angles. This process produces a rainbow of distinct colors known as a continuous spectrum. If, however, the light emitted from hot gases or energize
13、d ions is viewed in a similar manner, isolated bands of color are observed. These bands form characteristic patterns, unique to each element. They are known as bright line or emission spectra. Bright line spectra are not continuous like a rainbow, but instead are discrete bands. Some bands are brigh
14、ter than others. There are a few reasons this may happen. One reason is because there happens to be many electrons available to jump from a particular level to another and drop back down, whereas there may be only a few that jump a different interval. Another reason may be because other electromagne
15、tic radiation is being emitted whose waves interfere with those of the light waves.By analyzing the emission spectrum of hydrogen gas, Bohr was able to calculate the energy content of the major electron levels. Although the electron structure as suggested by his planetary atomic model has been modif
16、ied according to modern quantum theory, his description and analysis of spectral emission lines are still valid.In addition to the fundamental role spectroscopy played in the development of todays atomic model, this technique can also be used in the identification of elements. Since the atoms of each element contain unique arrangements of electrons, emission lines can b
