Chapter 5. The Electromagnetic Spectrum. What is visible light? What is visible light? Which of the following would you consider dangerous?

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1 Which of the following would you consider dangerous? X-rays Radio waves Gamma rays UV radiation Visible light Microwaves Infrared radiation Chapter 5 Periodicity and Atomic Structure 2 The Electromagnetic Spectrum One of the early uses for X-rays was for shoe fitting. It was phased out in the late 1950 s. What is visible light? Refraction of light makes separation possible. Angle of refraction varies with color Suggest that light is a wave EM waves What is visible light? Young s double-slit experiment Constructive interference Destructive interference Interference pattern suggests that light is made of waves

2 Behavior of Light Light is typically described as traveling in waves (similar to water) EM Spectrum Different colors of light correspond to different wavelengths in the visible portion of the EM spectrum. Two wavelengths are shown below. 750 c = ν λ Blue Light Red Light c = speed of light ( x 10 8 m/s) λ (lambda) = wavelength (m), ν (nu) = frequency (1/s, s -1, or Hz) Remember, although they have different wavelengths, they still travel the same speed (c) The Electromagnetic Spectrum Absorption vs Emission Spectra EM waves Continuous spectrum Lines correspond to all wavelengths of light within the light source Excited Gas Cold Gas Emission spectra Lines correspond to photons of discrete energies that are emitted when excited atoms release energy and electrons drop to a lower energy levels. Absorption spectra occur when light passes through a cold, dilute gas; Certain wavelengths are absorbed. These correspond to dark lines (absence of light) in the spectrum. Absorption vs Emission Spectra Emission Spectra of Elements 11

3 Bohr s Model of the Atom Absorption and emission of energy Also referred to as the planetary model Energy absorbed (light, heat, etc.) Energy released as light Electrons circle the nucleus in quantized (discrete) orbits Each orbit is a different energy level. Electrons in different orbits have different energies n=1 n=2 Energy Levels: 1<2<3 n=3 Electrons cannot exist between energy levels. There are only specific discrete energy levels within an atom n=1 n=2 n=3 When an atom absorbs energy, an electron can jump from a lower energy level to a higher energy level. n=1 n=2 n=3 Once an electron is excited, it can drop back down to a lower energy level. This releases energy in the form of ER. Emission Spectra How CFLs Work The emission spectra seen is a result of electrons dropping to lower energy level and emitting light Each line represents an electron with a specific energy within the atom - Fluorescent Lights Electrons excite Hg atoms. Hg electrons become excited, and then return to the ground state, giving off mostly UV photons (invisible) - - Hg UV UV The UV photons excite the electrons in the phosphor coating. Those electrons go to the ground state, which in turn gives off visible light Withou the phosphor coating, you would just have a black light. P h o s p h o r C o a t i n g If emissions spectra are dependent on the identity of the element, why does tungsten (W) in incandescent light bulbs give off white light? Emissions Spectra Given off in the gasstate Metals can be ionized into the gas state Discrete wavelengths given off Depends on element Blackbody Radiation Classical physics (assuming light as waves) could not accurately predict the intensity of blackbody radiation. Black Body Radiation Given off in the solidstate All wavelengths given off Depends only on temperature Sodium Lamp (gas phase) Tungsten Lamp (solid phase) Actual observations If light behaved as waves

4 Photoelectric Effect Scientist observed that electrons could be ejected from metals using electromagnetic radiation Not matter the intensity of the light, each metal had a minimum wavelength need to eject and electron. Wave-Particle Duality Based on photoelectric effect, light acts as a stream of particles called photons. These are the packets or quanta of energy Energy of photons is proportional to frequency, inversely proportional to wavelength Einstein was able to explain this assuming that light was a particle. E = hν Since ν = c/λ h c E = λ Planck s constant h = x Js (J = kg m 2 / s 2 ) Describing Atoms Classical descriptions: Dalton: atoms are hard particles, all atoms of the same element are the same Thomson: atoms are divisible (electrons in atoms) Rutherford: positive nucleus New view of atomic behavior: Planck: Blackbody radiation heat solids to red or white heat, matter did not emit energy continuously; in whole-number multiples of certain quantities Einstein: Photo electric effect minimum frequency needed to eject electrons Matter absorbs or emits energies in packets - quanta Wavelike Properties of Matter If light can behave like a wave and a particle, then maybe matter (i.e., electrons) can behave as a wave and a particle also. If an electron behaves like a standing wave, then it can only have specific wavelengths Therefore, an electron can only have specific frequencies and energies Waves: E = mc 2 Matter: E = hν de Broglie combined the two equations: m = h / λ c Can calculate wavelength for matter if we know its velocity (v instead of c): λ = h / m v The double-slit experiment revisited The double-slit experiment revisited How will electrons behave?? How particles behave no interference pattern How waves behave interference pattern Interference Pattern electrons, which are particles, behaving as waves

5 Heisenberg Uncertainty Principle If electrons have wavelike properties, then we can t know both its position and path. In order to determine position we can hit it with a photon of light, but this will change its position and momentum. Quantum Mechanical Model The Bohr model worked well for hydrogen, but failed for elements with more than one proton and one electron. Quantum Mechanics was developed (by Schrödinger in the 1920 s) to describe motions of subatomic particles Did not attempt to describe positions of particles; used mathematical equations to describe the probability of finding the particles Schrödinger used differential equations to find an electron s wave function (Ψ) which determines the probability of find the electron. The probability density (map of likely locations) is the electron cloud Atomic Orbitals The region of highest probability for finding an electron is an electron cloud. This region of high probability is called an atomic orbital. Each orbital holds at most 2 electrons. Where an electron will be found 90% of the time Shapes of Orbitals s orbitals are spherical 1 s orbital in a subshell Nodes: zero probability of finding an electron Shapes of Orbitals p orbitals are dumb-bells (2 lobes) 3 p orbitals in a subshell Shapes of Orbitals d orbitals are intersecting dumb-bells (4 lobes) 5 d orbitals in a subshell 31 32

6 Energies of Orbitals In hydrogen, all shells are equivalent in energy. This is because there is only 1 electron in the system. No other electrons with which to interact. Energies of Orbitals In many-electron models, the energy levels depend on the shell and subshell. Electron Configurations Aufbau principle: start with the nucleus and empty orbitals, then build up the electron configuration using orbitals of increasing energy. This is because there are >1 electrons in the system and the electrons interact with one another. This is why the Bohr model did not work for multi-electron systems. 36 Electron Configurations Aufbau principle: start with the nucleus and empty orbitals, then build up the electron configuration using orbitals of increasing energy. Electron Configurations (Figure 5.17) 37

7 Electron Configuration Writing Electron Configurations Every electron in every atom is found in an orbital. We can use the four quantum numbers to assign electrons to orbitals. No more than 2 electrons to one orbital! 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f 5s, 5p, 5d, 5f, (5g) 6s, 6p, 6d, (6f, 6g, 6h) 7s, 7p, (7d, 7f, 7g, 7h) Another way to conceptualize the order in which orbitals are filled with electrons (In theory) H He Li Be B N O Ne Na Al S Ar K Sc Ti Zn Br 39 Shells and Subshells shells are energy levels with the same number subshells are energy levels with different letters within the same shell Shorthand Notation Rather than writing out complete electron configurations, we can use the previously filled shell (noble gas) and show the valence electrons: P: 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne] 3s 2 3p 3 The 1 st shell has 1 subshell (s) The 2 nd shell has 2 subshells (s and p) The 3 rd shell has 3 subshells (s, p, and d) The 4 th shell has 4 subshells (s, p, d, and f) For example, the 3 rd energy levels is comprised of 3 subshells 41 Write the shorthand notation for: Ca Cl Sr Fe Electron Configuration Some exceptions to the Aufbau order What are the expected electron configurations for Cr and Cu? Filled and half-filled d subshells seem to be especially stable. Cr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Also true for Mo Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Also true for Ag and Au Quantum Numbers an electron s address The Principal Quantum Number, n describes distance of the electron from the nucleus; called shells n = 1, 2, 3, etc; larger number is farther from nucleus n corresponds to a row in the periodic table The Angular Momentum Quantum Number, l describes the shape of the orbital; labeled s, p, d, f; called subshells For n = 1, only the 1s subshell exists For n = 2, only the 2s and 2p subshells exist For n = 3, 3s, 3p, and 3d subshells exist 43 44

8 Quantum Numbers an electron s address The Magnetic Quantum Number, m l describes the orientation of the orbital with respect to x, y, and z axes s orbitals only have 1 orientation (spherical) p orbitals have 3 orientations (along x-axis, along y-axis, or along z-axis); p x, p y, and p z d orbitals have 5 orientations; d xy, d xz, d yz, d x2-y2, d z2 Electrons per shell Any shell (energy level) can hold a maximum of 2n 2 electrons So the third shell (n=3) can hold a maximum of (2)3 2 electrons.or 18 electrons The Spin Quantum Number, m s describes the spin of an electron in an orbital (shown as up and down arrows in orbital diagrams) 45 Things that make you go hmmm. How many electrons in the 2p x orbital? How many subshells in the 4 th shell? How many electrons in the 3d subshell? How many electrons in the 4d subshell? How many electrons in the 3p subshell? How many electrons in the 3 rd shell? Orbital Diagrams Orbital diagrams are pictorial representation s of electron configurations. Electron Configurations 48 Hund s Rule If two or more orbitals (i.e., a p or d orbital) with the same energy are available, one electron goes into each orbital until they have to pair up. Fighting sibling analogy: 3-bedroom house, 6 siblings For example, an atom with 2 p electrons: 1 electron will go into the p x orbital, the next electron will go into the p y orbital. Pauli Exclusion Principle Pauli Exclusion Principle: no two electrons can have the same values of all 4 quantum numbers Describes what happens when electrons share an orbital. Only two electrons can occupy a single orbital and they must have opposite spin (i.e., the 4 th quantum number). The first electron is designated as positive spin, the second electron in that orbital has negative spin. 50

9 Calculation Practice c = λν E = hν 1) Which has a higher frequency: light from a red stoplight with a wavelength of 750 nm or a yellow light with a wavelength of 600 nm? 2) What is the wavelength of a radio station s waves transmitting at a frequency of MHz (megahertz)? (FM radio waves range from MHz.) 3) Red lights at traffic stops have wavelengths of about 650 nm. What is the frequency (in Hz) of this light? 4) Calculate the energy (in Joules) of a photon with a wavelength of 5.00 x 10 4 nm (infrared region). Calculation Practice The energy of a photon is 5.87 x J. What is the frequency of the photon? What is the wavelength of an electron that travels at 34.7 m/s and has a mass of 9.11 x kg? A kg baseball is thrown at a velocity of 42.5 m/s. Calculate the wavelength of the baseball. How does the baseball s wavelength compare to the electron from the example above?