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Spectrophotometry Chapter 18 Review of Electron Energies • The electrons of an atom have different energies. • Not all energies exist, only certain allowed energy levels. • Electrons with more energy are able to get farther away from the nucleus and its + charges. • Therefore, electrons in higher energy levels spend more time farther away from the nucleus. • The higher energy levels are larger so they can hold more electrons. • Electrons are not orbiting the nucleus like planets! • The energy levels of the electrons are divided into sublevels and orbitals. Light as a Wave • Light is also known as electromagnetic radiation. – energy radiated in the form of a WAVE caused by an electric field interacting with a magnetic field – result of the acceleration of a charged particle – does not require a material medium and can travel through a vacuum Properties of Waves All waves have: • wavelength (, “lambda”). – The distance between 2 crests of a wave. – In units of length (m, cm, nm) • frequency (, “nu”). – The # of waves that oscillate per second. – In units of 1/sec or Hertz, Hz • speed (for light waves speed = c) – In units of m/s • Amplitude (A) – The “height” of the wave The Wave Nature of Light (electromagnetic radiation or EMR) • The speed of a wave, v, is given by its frequency multiplied by its wavelength: • v = (units are m/s) • For light in a vacuum, speed v = c =2.9979 x 108 m/s and this is the fastest anything could travel. • c = (units are m/s) – Note that the longer the wavelength the smaller frequency. – Or the shorter the wavelength the higher (larger) the frequency. Visible Light: The Light our Eyes Can Detect • Wavelengths: ~ 400 to 750 nanometers. – Each wavelength corresponds to a different color of light. – Red light: ~ 620 - 750 nm – Violet light: ~ 390 – 450 nm • White light is a mixture of all the colors of the rainbow. – A prism can split white light into its colors. Electromagnetic Radiation (EMR) Spectrum • Visible light is only one type of EMR. • Other types include: – Gamma rays – X-rays – Ultraviolet – Infrared – Microwaves – Radio and TV waves • All of these have different wavelengths. What is the frequency of blue light with = 440 nm? • c= • = c/ • = (3.00 x 108 m/s)/(440 x 10-9 m) • = 6.8 x 1014 1/s • = 6.8 x 1014 Hz Wave Number • The relationship between energy and frequency of light is: E = h where h is Planck’s constant (6.626 10-34 Js). • Using the above equation and the relationship between wavelength and frequency we get: E = h c/ • Since h and c are constants, it follows that 1/ is proportional to E. • The quantity 1/ is given the symbol and is called the wave number. Line Spectra • Each element emits (gives off) its own, individual set of colors when energized. • This is called its atomic emission spectrum. (or line spectrum) – The colors show up as a series of lines when viewed through a prism. – Called spectral lines. • Like a fingerprint for that element. Ground and Excited States • Ground State: The state of least possible energy in a physical system, as of elementary particles. – Also called ground level. • Excited State: Being at an energy level higher than the ground state Absorption • A photon of light hits a molecule or atom. • If the energy of the photon is “right”, it is absorbed. • • An electron excited to a higher energy level. • The energy of the photon must be exactly equal to the difference of the electron energy levels. Photons that aren’t absorbed are scattered (or transmitted). Types of Transitions • Light in the visible and UV may cause electronic transitions. • IR radiation effects vibrational states of molecules (more later). • Microwave radiation effects the rotational states of molecules (more later). Emission • An electron is excited to a higher energy level by an outside energy. – Heat, electricity, etc. • The electron returns to a lower energy level. • A photon (tiny packet) of light is emitted. • The energy of the photon is equal to the difference between the energy levels of the electron. Spectrophotometry • In a spectrophotometry experiment, we measure the light absorbed or emitted. • This week in lab: absorbance in the visible region. Absorption Methods: Transmittance T = P/Po where T => transmittance P => power of transmitted radiation Po => power of incident radiation %T = (P/Po)*100 where %T => percent transmittance Absorption Methods, Absorbance A = - log10T = - log10 (P/Po) where A => absorbance Relation Between Transmittance and Absorbance P/Po %T A 1 100 0 0.1 10 1 0.01 1 2 Absorption Methods, Beer’s Law A = abc = bc where a => absorptivity b => path length c => concentration => molar absorptivity Homework Chp 18 Exercises: 18A, B, C Problems: 1 – 4, 6 – 12, 16, 18