CEM 141, notes, week 5
CEM 141, notes, week 5 Cem 141
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This 6 page Class Notes was uploaded by Leah DiCiesare on Friday October 7, 2016. The Class Notes belongs to Cem 141 at Michigan State University taught by J. Hu in Fall 2016. Since its upload, it has received 8 views. For similar materials see General Chemistry in Chemistry at Michigan State University.
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Date Created: 10/07/16
Lecture Notes: 10/310/7 Electromagnetic Radiation (EM Radiation) o Ex. Radio waves, microwaves, infrared, visible, ultraviolet, xrays, Gamma rays o Light can be described as a wave or a particle (waveparticle duality) o EM radiation can be described as a wave compound of oscillating electric and magnetic fields. The fields oscillate in perpendicular planes. o Wavelength ( λ is the distance between any two identical points in neighboring waves, measured in meters o Amplitude is the distance from the average value to the maximum value, the intensity o Frequency (v) is the number of oscillations from the average value to the made by the wave per second (Hz) 8 o c=λv c=velocity of light in vacuum (3.0*10 m/s) λ and v have an inverse relationship o Energy increases as frequency increases (and wavelength decreases) no a typical wave (for electromagnetic wave) o ROY G. BIV order of visible light in spectrum (Red, Orange, Yellow, Green, Blue, Indigo, Violet) R lowest frequency and highest wavelength V highest frequency and lowest wavelength o In order to detect something using EM radiation we must use a wavelength smaller than the objects we are observing Properties of a wave o Diffraction a wave will pass through a barrier with a slit and be small and then spread out again o Diffraction pattern through one barrier slit, strongest intensity in middle through two slits interference patterns with intense bright spots and darks spots that alternate o Interference if two waves meet in the same space in the same phase and they join it will have doubles amplitude (brighter light) constructive; destructive if waves are in opposite phase, they cancel each other out; usually somewhere in between Photoelectric Effect o Many metals emit electrons when electromagnetic radiation (light) shines on the surface Lecture Notes: 10/310/7 o The light is transferring energy to the electrons at the metal surface where it is transformed into kinetic energy that gives the electrons enough energy to "leave" the atoms in metal o Need energy high enough to overcome the electrostatic force between electrons and nucleus o Evidence When a short wavelength (high energy) light shines on metal surface electrons are emitted measureable current If light intensity is increased more electrons are emitted There is a threshold frequency below which no electrons are emitted no matter how bright (intense) the light is If light were just a wave increasing the intensity should increase the energy and eject more electrons no threshold in frequency Only light above a certain threshold frequency (energy) will result in ejected electrons Below the threshold increasing the intensity has no effect Einstein postulated that light must come in as packets (or particle or quanta) called photons Paradigm shift fundamental change in the basic concepts and experimental practices of scientific discipline EM Radiation is a Particle o Energy is transferred as a particle (photon) that has a definable energy o PlanckEinstein relation: E=hv (h is Planck's constant) One photon can interact and eject one electron, if it has enough energy. If the photon does not have enough energy then no electron is ejected o Planck: light must have quantized energy related with frequency and proposed the equation: E=hv (h=6.626*10 Js) But failed to provide physical explanation o Einstein: photoelectric effect threshold of frequency (energy) "photon" (particle) Nobel Price 1921 o Increasing intensity = increasing photon number Decreasing frequency = lower energy of each photon o o Decreasing frequency below threshold = energy in photo is too low Summary of Electromagnetic Radiation o Can be described as either a particle or a wave These are models not reality o Truly difficult to imagine these ideas how can one phenomenon be two different things? o Waveparticle duality is important at very small scales Matter and energy don't behave the same as in our macroscopic world Absorption and Emission Spectra o Visible Spectrum: light from sun (white light) can be separated by a prism (Newton did this first) only a very small part of the EM spectra Lecture Notes: 10/310/7 o Atomic Emission Spectrum: light from one particular atom does not contain all the colors of the spectrum has only a few wavelengths use hot gas to test o Atomic Absorption Spectrum: light that is absorbed in an atomized sample of gas (cold); continuous spectrum with dark lines that show which light was absorbed by the sample o Emission + Absorption Spectra = continuous spectrum; they are complementary o Each element has characteristic wavelengths that it can absorb or emit o Spectra show light only of specific wavelengths/energies the spectrum of element is the same whether that element is on Earth, in the Sun, of in a galaxy light years away Niels Bohr Model o Electrons move in orbits around nucleus o These orbits have definite energies and are at definite distances from the nucleus o So the energies are quantized o Explained emission and absorption spectra by invoking discrete energy levels characterized by quantum numbers (n) Has to have the exact same energy to change electron orbit levels or emit electrons o Photons of electromagnetic energy are emitted or absorbed by atoms as electrons move from one energy level to another o The energy of the photons corresponds to the difference in energy levels of the electrons o Only works for hydrogen so it is misleading o Introduced quantized energy levels of electrons Better to use energy diagrams Lecture Notes: 10/310/7 o Each energy level has a quantum number o The higher the number, the higher the energy o Energy levels are not orbits o Electrons transition between energy levels by absorbing or emitting photons Textbook Notes: 2.12.4 2.1 Light and Getting Quantum Mechanical o Evidence that light is both a wave and a particle o James Clerk Maxwell (18311879) developed the electromagnetic theory of light, in which visible light and other waves were viewed in terms of perpendicular electric and magnetic fields o A light wave can be described by defining its frequency (v) and its wavelength (λ) λv =c c=velocity of light o Max Planck (18581947), German physicist An object heated to a particular temperature emits radiation (infrared) Studied how the color of the light emitted changed as a function of an object's temperature Result known as the ultraviolet catastrophe reproducible data that didn't follow the theory but also the theory could not be modified to accept this data Matter absorbs and emits energy only in discrete chunks called quanta; quanta occurred in multiples of E (energy)=hv; h is Planck's constant and v is frequency of light 2.2 Taking Quanta Seriously o 1905 Einstein used idea of quanta to explain the photoelectric effect which was described and patented by Nikola Tesla o Photoelectric effect occurs when light shines on a metal plate and electrons are ejected, creating a current o There is a threshold wavelength (energy) of light that is a characteristic for the metal used, beyond which no electrons are ejected Explained by the idea that light comes in particle form known as photons which also have a wavelength and frequency o Intensity of light is related to the number of photons that pass by per second o Energy per photon is dependent upon its frequency or wavelength c=3.0*108 o o λv=c; higher frequency (v), shorter wavelengλh ( ) and greater energy per photon Energy is directly related to frequency but inversely related to wavelength E=hc/λ o Once the wavelength is short enough, or the energy is high enough to eject electrons, increasing the intensity of the light now increases the number of electrons ejected There has to be enough energy to overcome the attraction between the electro and the nucleus o All matter has a wavelength (a wavelike property) Louis de Broglie (18921987)λ =h/mv, mv is the momentum (mass*velocity) 2.3 Exploring Atomic Organization Using Spectroscopy o What makes rainbows possible is the face that sunlight is composed of photons with an essentially continuous distribution of visible wavelengths Textbook Notes: 2.12.4 o When a dense body, like the sun, is heated it emits light of many wavelengths visible colors; when a sample of an element or mixture is heated it emits light of only very particular wavelengths o If white light were to pass through a cold gaseous element the same wavelengths that were emitted (when it was heated) would be absorbed while all the other wavelengths would pass through o Emission and absorption wavelengths for each element are unique o Making sense of the Spectra Niels Bohr (18851962) proposed a new model for the atom 1st hypothesis: electrons within an atom can only travel along certain orbits at a fixed distance from the nucleus, each orbit corresponding to a specific energy 2nd idea: electrons can jump from one orbit to another, but this requires the absorption or emission of energy, in the form of a photon Photon has to be exactly the right amount of energy Lower, more stable orbits are visualized as being closer to the nucleus while higher, less stable, and more energetic orbits are further away When enough energy is supplied all at once an electron is removed completely, leaving a positivelycharged ion Only could predict emission/absorption spectrum for hydrogen 2.4 Beyond Bohr o Louis de Broglie used Planck's relationship between energy and frequency (E=hv), the relationship between frequency and wavelength λc= v), and Einstein's relationship between energy and mass (E=mc ) to derive a relationship between mass and wavelength for any particlλ ( =h/mv) At atomic scale, wavelengths associated with particles are similar to their size o Certainty and Uncertainty Have to use electromagnetic radiation of a wavelength similar to the size of an electron to see it When wavelengths of that small interact with an electron it changes the electron's position and motion The uncertainty created by trying to measure where something is gets greater the closer we get to the atomicmolecular scale Idea by Werner Heisenberg (19011976) Heisenberg Uncertainty Principle Can estimate uncertainty usinΔ mv*Δx>h/2 π Δmv=uncertainty in momentum of particle,Δ x is uncertainty in position
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