IB Chemistry on Line Emission Spectrum, Bohr Model and Electromagnetic Spectrum
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IB Chemistry on Line Emission Spectrum, Bohr Model and Electromagnetic Spectrum

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IB Chemistry on Line Emission Spectrum, Bohr Model and Electromagnetic Spectrum

IB Chemistry on Line Emission Spectrum, Bohr Model and Electromagnetic Spectrum

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    IB Chemistry on Line Emission Spectrum, Bohr Model and Electromagnetic Spectrum IB Chemistry on Line Emission Spectrum, Bohr Model and Electromagnetic Spectrum Presentation Transcript

    • Tutorial on Electromagnetic Radiation, Emission Line spectrum and Bohr Model. Prepared by Lawrence Kok http://lawrencekok.blogspot.com
    • Electromagnetic Spectrum Electromagnetic spectrum ranges from Radiowaves to Gamma waves. - Form of energy - Shorter wavelength -> Higher frequency -> Higher energy - Longer wavelength -> Lower frequency -> Lower energy
    • Electromagnetic Spectrum Electromagnetic spectrum ranges from Radiowaves to Gamma waves. - Form of energy - Shorter wavelength -> Higher frequency -> Higher energy - Longer wavelength -> Lower frequency -> Lower energy Wavelength, λ - long  Frequency, f - low  Wavelength, λ - short  Inverse relationship between- λ and f Frequency, f - high 
    • Electromagnetic Spectrum Electromagnetic spectrum ranges from Radiowaves to Gamma waves. - Form of energy - Shorter wavelength -> Higher frequency -> Higher energy - Longer wavelength -> Lower frequency -> Lower energy Wavelength, λ - long  Frequency, f - low  Wavelength, λ - short  Inverse relationship between- λ and f Frequency, f - high  Electromagnetic radiation • Travel at speed of light, c = fλ -> 3.0 x 108 m/s • Light Particle – photon have energy given by -> E = hf • Energy photon - proportional to frequency Plank constant • proportionality constant bet energy and freq Excellent video wave propagation Click here to view.
    • Electromagnetic Wave propagation. Electromagnetic radiation • • • Moving charges/particles through space Oscillating wave like property of electric and magnetic field Electric and magnetic field oscillate perpendicular to each other and perpendicular to direction of wave propagation. Electromagnetic radiation Electromagnetic wave propagation Click here to view video
    • Electromagnetic Wave propagation. Electromagnetic radiation • • • Moving charges/particles through space Oscillating wave like property of electric and magnetic field Electric and magnetic field oscillate perpendicular to each other and perpendicular to direction of wave propagation. Electromagnetic radiation Electromagnetic wave propagation Click here to view video Wave Wave – wavelength and frequency - travel at speed of light
    • Electromagnetic Wave propagation. Electromagnetic radiation • • • Moving charges/particles through space Oscillating wave like property of electric and magnetic field Electric and magnetic field oscillate perpendicular to each other and perpendicular to direction of wave propagation. Electromagnetic radiation Electromagnetic wave propagation Click here to view video Violet λ = 410nm f = c/λ = 3 x 108/410 x 10-9 = 7.31 x 1014 Hz E = hf = 6.626 x 10-34 x 7.31 x 1014 = 4.84 x 10-19 J Wave Wave – wavelength and frequency - travel at speed of light Red λ = 700nm f = c/λ = 3 x 108/700 x 10-9 = 4.28 x 1014 Hz E = hf = 6.626 x 10-34 x 4.28 x 1014 = 2.83 x 10-19 J
    • Electromagnetic Wave propagation. Electromagnetic radiation • Moving charges/particles through space • Oscillating wave like property of electric and magnetic field • Electric and magnetic field oscillate perpendicular to each other and perpendicular to direction of wave propagation. Electromagnetic radiation Is it a particle or Wave? Click to view video -Wave-particle duality Wave Wave – wavelength and frequency - travel at speed of light
    • Electromagnetic Wave propagation. Electromagnetic radiation • Moving charges/particles through space • Oscillating wave like property of electric and magnetic field • Electric and magnetic field oscillate perpendicular to each other and perpendicular to direction of wave propagation. Electromagnetic radiation Is it a particle or Wave? Click to view video -Wave-particle duality Wave Wave – wavelength and frequency - travel at speed of light Simulation on Electromagnetic Propagation Click here to view simulation Click here to view simulation Click here to view simulation
    • Electromagnetic Wave Violet Red λ = 410nm λ = 700nm f = c/λ = 3 x 108/410 x 10-9 = 7.31 x 1014 Hz f = c/λ = 3 x 108/700 x 10-9 = 4.28 x 1014 Hz Wavelength – Distance bet two point with same phase, bet crest/troughs – unit nm Frequency – Number of cycle/repeat per unit time (cycles in 1 second) – unit Hz
    • Electromagnetic Wave Violet Red λ = 410nm λ = 700nm f = c/λ = 3 x 108/410 x 10-9 = 7.31 x 1014 Hz f = c/λ = 3 x 108/700 x 10-9 = 4.28 x 1014 Hz Wavelength – Distance bet two point with same phase, bet crest/troughs – unit nm Frequency – Number of cycle/repeat per unit time (cycles in 1 second) – unit Hz Which wave have higher frequency, if both have same speed reaching Y same time? Violet X Y Red
    • Electromagnetic Wave Violet Red λ = 410nm λ = 700nm f = c/λ = 3 x 108/410 x 10-9 = 7.31 x 1014 Hz f = c/λ = 3 x 108/700 x 10-9 = 4.28 x 1014 Hz Wavelength – Distance bet two point with same phase, bet crest/troughs – unit nm Frequency – Number of cycle/repeat per unit time (cycles in 1 second) – unit Hz Which wave have higher frequency, if both have same speed reaching Y same time? Violet X Click here on excellent video red /violet wave Click here to view video energy photon Y Light travel same speed Red flippers – long λ - less frequent Violet shoes – short λ - more frequent Red
    • Continuous Spectrum Vs Line Spectrum Continuous Spectrum : Light spectrum with all wavelength/frequency Emission Line Spectrum : • Spectrum with discrete wavelength/ frequency • Emitted when excited electrons drop from higher to lower energy level Absorption Line Spectrum : • Spectrum with discrete wavelength/frequency • Absorbed when ground state electrons are excited
    • Continuous Spectrum Vs Line Spectrum Continuous Spectrum : Light spectrum with all wavelength/frequency Emission Line Spectrum : • Spectrum with discrete wavelength/ frequency • Emitted when excited electrons drop from higher to lower energy level Absorption Line Spectrum : • Spectrum with discrete wavelength/frequency • Absorbed when ground state electrons are excited Atomic Emission Electrons from excited state Excited state Emit radiation when drop to ground state Radiation emitted Emission Spectrum Ground state http://www.astrophys-assist.com/educate/orion/orion02.htm
    • Continuous Spectrum Vs Line Spectrum Continuous Spectrum : Light spectrum with all wavelength/frequency Emission Line Spectrum : • Spectrum with discrete wavelength/ frequency • Emitted when excited electrons drop from higher to lower energy level Absorption Line Spectrum : • Spectrum with discrete wavelength/frequency • Absorbed when ground state electrons are excited Atomic Emission Vs Atomic Absorption Spectroscopy Electrons from excited state Excited state Electrons in excited state Emit radiation when drop to ground state Radiation absorbed Radiation emitted Absorb radiation to excited state Emission Spectrum Ground state http://www.astrophys-assist.com/educate/orion/orion02.htm Electrons from ground state
    • Line Emission Spectroscopy Line Emission Spectra for Hydrogen Energy supplied to atoms • Electrons excited - ground to excited states • Electrons exist fixed energy level (quantum) • Electrons transition from higher to lower, emit energy of particular wavelength/frequency - photon • Higher the energy level, smaller the difference in energy bet successive energy level. • Spectrum converge (get closer) with increase freq. • Lines spectrum converge- energy levels also converge • Ionisation energy determined (Limit of convergence) UV region Lyman Series n=∞ → n= 1 Visible region Balmer Series n=∞ → n= 2 IR region Paschen Series n=∞ → n= 3
    • Line Emission Spectroscopy Line Emission Spectra for Hydrogen Energy supplied to atoms • Electrons excited - ground to excited states • Electrons exist fixed energy level (quantum) • Electrons transition from higher to lower, emit energy of particular wavelength/frequency - photon • Higher the energy level, smaller the difference in energy bet successive energy level. • Spectrum converge (get closer) with increase freq. • Lines spectrum converge- energy levels also converge • Ionisation energy determined (Limit of convergence) UV region Lyman Series n=∞ → n= 1 Visible region Balmer Series n=∞ → n= 2 IR region Paschen Series n=∞ → n= 3 Line Emission Spectra • Energy supplied • Electrons surround nucleus in allowed energy states (quantum) • Excited electron return to lower energy level, photon with discrete energy/wavelength (colour) given out. • Light pass through spectroscope (prism/diffraction grating) to separate out diff colours N= 6-2 410nm N= 5-2 434nm N= 4-2 486nm N = 3-2, 656nm Visible region- Balmer Series
    • Line Emission Spectroscopy Line Emission Spectra for Hydrogen Energy supplied to atoms • Electrons excited - ground to excited states • Electrons exist fixed energy level (quantum) • Electrons transition from higher to lower, emit energy of particular wavelength/frequency - photon • Higher the energy level, smaller the difference in energy bet successive energy level. • Spectrum converge (get closer) with increase freq. • Lines spectrum converge- energy levels also converge • Ionisation energy determined (Limit of convergence) UV region Lyman Series n=∞ → n= 1 Visible region Balmer Series n=∞ → n= 2 IR region Paschen Series n=∞ → n= 3 Line Emission Spectra • Energy supplied • Electrons surround nucleus in allowed energy states (quantum) • Excited electron return to lower energy level, photon with discrete energy/wavelength (colour) given out. • Light pass through spectroscope (prism/diffraction grating) to separate out diff colours Videos on line emission N= 6-2 410nm Click here to view video Click here to view video N= 5-2 434nm N= 4-2 486nm N = 3-2, 656nm Visible region- Balmer Series
    • Hydrogen Emission Spectroscopy – Visible region (Balmer Series) Line Emission Spectra for Hydrogen Excited state 5 4 3 2 Visible region Balmer Series n=∞ → n= 2 Ground state Click here for detail notes 1 Click here video line emission spectrum
    • Hydrogen Emission Spectroscopy – Visible region (Balmer Series) Line Emission Spectra for Hydrogen Hydrogen discharge tube Excited state 5 4 3 2 Visible region Balmer Series n=∞ → n= 2 Ground state Click here for detail notes 1 Click here video line emission spectrum Hydrogen Emission Spectroscopy
    • Hydrogen Emission Spectroscopy – Visible region (Balmer Series) Line Emission Spectra for Hydrogen Hydrogen discharge tube Excited state Hydrogen Emission Spectroscopy 5 4 3 n= 5-2 n = 3-2 n= 4-2 2 λ = 434nm Visible region Balmer Series n=∞ → n= 2 Ground state 1 f = c/λ = 3 x 108/434 x 10-9 = 6.90 x 1014 Hz λ = 486nm λ = 656nm f = c/λ = 3 x 108/656 x 10-9 = 4.57 x 1014 Hz E = hf = 6.62 x 10-34 x 6.90 x 1014 = 4.56 x 10-19 J More energetic violet line Click here for detail notes Click here video line emission spectrum E = hf = 6.62 x 10-34 x 4.57 x 1014 = 3.03 x 10-19 J Less energetic red line
    • Bohr Model for Hydrogen Atom – Ionization Energy Bohr Model Energy level Electronic Transition bet levels Niels Bohr Model (1913) • • • Electrons orbit nucleus. Orbits with discrete energy levels – Quantized. Transition electron bet diff levels by absorb/emit radiation with frequency, f determined by energy diff bet levels -ΔE = hf • Energy light emit/absorb equal to diff bet energy levels
    • Bohr Model for Hydrogen Atom – Ionization Energy Bohr Model Energy level Electronic Transition bet levels Niels Bohr Model (1913) • • • Electrons orbit nucleus. Orbits with discrete energy levels – Quantized. Transition electron bet diff levels by absorb/emit radiation with frequency, f determined by energy diff bet levels -ΔE = hf • Energy light emit/absorb equal to diff bet energy levels Light emitted equal to difference bet energy levels, -ΔE = hf Ionization energy Transition electron from 1 ->∞ ∞ Plank equation Higher energy level n, smaller the difference in energy bet successive energy level. 5 4 3 ΔE = hf Light given off 2 Light energy - ΔE = hf Frequency = ΔE/h 1
    • Bohr Model for Hydrogen Atom – Ionization Energy Energy level Bohr Model Electronic Transition bet levels Niels Bohr Model (1913) • • • Electrons orbit nucleus. Orbits with discrete energy levels – Quantized. Transition electron bet diff levels by absorb/emit radiation with frequency, f determined by energy diff bet levels -ΔE = hf • Energy light emit/absorb equal to diff bet energy levels Light emitted equal to difference bet energy levels, -ΔE = hf Ionization energy Transition electron from 1 ->∞ ∞ Plank equation Higher energy level n, smaller the difference in energy bet successive energy level. 5 4 3 ΔE = hf 2 Light given off Light energy - ΔE = hf Frequency = ΔE/h 1 line converge UV region Lyman Series n=∞ → n= 1 Increase freq  Line spectrum converge (get closer) with increase freq Ionisation energy determined (Limit of convergence) line converge Visible region Balmer Series n=∞ → n= 2 Increase freq  Line spectrum converge (get closer) with increase freq Lines in spectrum converge- energy levels also converge
    • Energy Level/Ionization Energy Calculation ∞ Formula - energy level, n (eV) n = energy level 5 5 4 4 3 3 1 2 2 Energy level, n= 3 = -13.6/n2 = -13.6/32 = -1.51 eV 3 Energy level, n= 2 = -13.6/n2 = -13.6/22 = -3.4 eV 4 Energy level, n= 1 = -13.6/n2 = -13.6/1 = -13.6 eV 2 constant 10-19 J 1eV – 1.6 x h = 6.626 x 10-34 Js 1 1
    • Energy Level/Ionization Energy Calculation ∞ Formula - energy level, n (eV) n = energy level 5 5 4 4 3 3 1 2 2 Energy level, n= 3 = -13.6/n2 = -13.6/32 = -1.51 eV 3 Energy level, n= 2 = -13.6/n2 = -13.6/22 = -3.4 eV 4 Energy level, n= 1 = -13.6/n2 = -13.6/1 = -13.6 eV 2 constant 10-19 J 1eV – 1.6 x h = 6.626 x 10-34 Js 1 1 Higher energy level, n - more unstable electron - More + ve ( less negative) - More energetic 5 6 Ionization energy Transition electron from 1 ->∞ Lower energy level, n - more stable electron - more – ve (-13.6eV) - Less energetic
    • Energy Level/Ionization Energy Calculation Energy difference bet level 3 to 2 ∞ Formula - energy level, n (eV) n = energy level 5 1 4 4 3 Energy difference, n= 3-2 = -1.51 – (-3.4) eV = 1.89 eV = 1.89 x 1.6 x 10-19 J = 3.024 x 10-19 J 5 3 1 2 Energy level, n= 3 = -13.6/n2 = -13.6/32 = -1.51 eV 3 Energy level, n= 2 = -13.6/n2 = -13.6/22 = -3.4 eV 4 Energy level, n= 1 = -13.6/n2 = -13.6/1 = -13.6 eV Light given off 2 Light energy - ΔE = hf Frequency, f = ΔE/h 2 2 constant 3 10-19 J Frequency, f = ΔE/h f = 3.024 x 10-19 /6.626 x 10-34 = 4.56 x 1015 Hz 4 λ = c/f = 3 x 108/4.56 x 1015 = 657 x 10-9 = 657nm 1eV – 1.6 x h = 6.626 x 10-34 Js 1 1 Higher energy level, n - more unstable electron - More + ve ( less negative) - More energetic 5 Light given off 6 Ionization energy Transition electron from 1 ->∞ Lower energy level, n - more stable electron - more – ve (-13.6eV) - Less energetic
    • Ionization Energy for Hydrogen Atom ∞ 1 n = energy level 5 ∞ 5 4 4 3 Ionization energy Min energy to remove 1 mole electron from 1 mole of element in gaseous state M(g)  M+ (g) + e 3 2 Ionization energy Transition electron from 1 ->∞ Energy Absorb 2 2 3 Energy level, n= ∞ = -13.6/n2 = -13.6/∞ = o eV 4 Energy level, n= 1 = -13.6/n2 = -13.6/1 = -13.6 eV electron Light/photon ABSORB by electron 1 1
    • Ionization Energy for Hydrogen Atom ∞ 1 n = energy level 5 ∞ 5 4 4 3 Ionization energy Min energy to remove 1 mole electron from 1 mole of element in gaseous state M(g)  M+ (g) + e 3 2 Ionization energy Transition electron from 1 ->∞ Energy Absorb 2 2 3 Energy level, n= ∞ = -13.6/n2 = -13.6/∞ = o eV 4 Energy level, n= 1 = -13.6/n2 = -13.6/1 = -13.6 eV electron Light/photon ABSORB by electron 1 1 5 6 Energy difference, n= 1-> ∞ = 0 – (-13.6) eV = 13.6 eV = 13.6 x 1.6 x 10-19 J = 2.176 x 10-18 J for 1 electron Energy absorb for 1 MOLE electron - 2.176 x 10-18 J - 1 electron - 2.176 x 10-18 x 6.02 x 1023 J - 1 mole - 1309kJ mol-1
    • Light given off, electronic transition from high -> low level Light given off Energy Released ∞ Ionization Energy for Hydrogen Atom 1 n = energy level 5 Energy difference, n= 3-2 = -1.51 – (-3.4) eV = 1.89 eV = 1.89 x 1.6 x 10-19 J = 3.024 x 10-19 J 2 4 4 3 Energy difference bet level 3 to 2 1 ∞ 5 Ionization energy Min energy to remove 1 mole electron from 1 mole of element in gaseous state M(g)  M+ (g) + e 3 2 Ionization energy Transition electron from 1 ->∞ Light given off Energy Absorb 2 2 3 3 Frequency, f = ΔE/h f = 3.024 x 10-19 /6.626 x 10-34 = 4.56 x 1015 Hz 4 Energy level, n= 1 = -13.6/n2 = -13.6/1 = -13.6 eV Light energy - ΔE = hf Frequency, f = ΔE/h 4 Energy level, n= ∞ = -13.6/n2 = -13.6/∞ = o eV electron Light/photon ABSORB by electron 1 5 1 λ = c/f = 3 x 108/4.56 x 1015 = 657 x 10-9 = 657nm Light given off 5 6 Energy difference, n= 1-> ∞ = 0 – (-13.6) eV = 13.6 eV = 13.6 x 1.6 x 10-19 J = 2.176 x 10-18 J for 1 electron Energy absorb for 1 MOLE electron - 2.176 x 10-18 J - 1 electron - 2.176 x 10-18 x 6.02 x 1023 J - 1 mole - 1309kJ mol-1
    • Energy Level/Ionization Energy Calculation n = energy level Energy/Wavelength – Plank/Rydberg Equation ∞ Formula – Plank Equation 5 5 4 4 ΔE = hf 3 3 ∞ Rydberg Equation to find wavelength 2 2 R = Rydberg constant R = 1.097 x 107 m-1 1 1 Nf = final n level Ni = initial n level
    • Energy photon- high -> low level 1 Electron transition from 3 -> 2 Energy Level/Ionization Energy Calculation n = energy level Energy/Wavelength – Plank/Rydberg Equation Light given off Formula – Plank Equation 5 Rydberg Eqn find wavelength emit ∞ 5 4 4 ΔE = hf 3 3 ∞ Rydberg Equation to find wavelength 2 2 2 nf = 2, ni = 3 R = 1.097 x 107 3 R = Rydberg constant R = 1.097 x 107 m-1 4 5 λ = 657 x 10-9 = 657 nm f = c/λ = 3 x 108/657 x 10-9 = 4.57 x 1014 Hz Light given off 1 1 Nf = final n level Ni = initial n level
    • Energy photon- high -> low level 1 Electron transition from 3 -> 2 Energy Level/Ionization Energy Calculation n = energy level Energy/Wavelength – Plank/Rydberg Equation Light given off Formula – Plank Equation 5 Rydberg Eqn find wavelength emit ∞ 5 4 4 ΔE = hf 3 3 ∞ Rydberg Equation to find wavelength 2 2 2 nf = 2, ni = 3 R = 1.097 x 107 3 R = Rydberg constant R = 1.097 x 107 m-1 4 5 λ = 657 x 10-9 = 657 nm f = c/λ = 3 x 108/657 x 10-9 = 4.57 x 1014 Hz 1 1 Click here on energy calculation Light given off Click here to view video Click here to view video Nf = final n level Ni = initial n level
    • Light given off, high -> low level Energy photon-electronic transition from high -> low level 1 Electron transition from 3 -> 2 ∞ n = energy level 5 Light given off 4 4 3 3 2 Rydberg Eqn find wavelength emit ∞ 5 2 2 nf = 2, ni = 3 R = 1.097 x 107 3 4 5 λ = 657 x 10-9 = 657 nm f = c/λ = 3 x 108/657 x 10-9 = 4.57 x 1014 Hz Light given off 1 1
    • Light given off, high -> low level Energy photon-electronic transition from high -> low level 1 Electron transition from 3 -> 2 ∞ Ionization Energy for Hydrogen Atom 1 n = energy level 5 Rydberg Eqn find wavelength emit Light given off ∞ 5 4 4 3 Ionization energy Min energy to remove 1 mole electron from 1 mole of element in gaseous state M(g)  M+ (g) + e 3 Ionization energy Transition electron from 1 -> ∞ 1 2 Energy Absorb 2 2 nf = 2, ni = 3 R = 1.097 x 107 Rydberg Eqn find ionization energy 3 3 electron Light/photon ABSORB by electron 4 λ = 657 x 10-9 = 657 nm 1 nf = ∞, ni = 1 R = 1.097 x 107 1 4 5 f = c/λ = 3 x 108/657 x 10-9 = 4.57 x 1014 Hz Light given off
    • Light given off, high -> low level Energy photon-electronic transition from high -> low level 1 Electron transition from 3 -> 2 ∞ Ionization Energy for Hydrogen Atom 1 n = energy level 5 Rydberg Eqn find wavelength emit Light given off ∞ 5 4 4 3 Ionization energy Min energy to remove 1 mole electron from 1 mole of element in gaseous state M(g)  M+ (g) + e 3 Ionization energy Transition electron from 1 -> ∞ 1 2 Energy Absorb 2 2 nf = 2, ni = 3 R = 1.097 x 107 Rydberg Eqn find ionization energy 3 3 electron Light/photon ABSORB by electron 4 1 λ = 657 x 10-9 = 657 nm nf = ∞, ni = 1 R = 1.097 x 107 1 4 5 f = c/λ = 3 x 108/657 x 10-9 = 4.57 x 1014 Hz λ = 9.11 x 10-8 7 Light given off Energy absorb for 1 MOLE electron - 2.179 x 10-18 J - 1 electron - 2.179 x 10-18 x 6.02 x 1023 J - 1 mole - 1312kJ mol-1 6 Energy, E = hf = 6.626 x 10-34 x 3.29 x 1015 = 2.179 x 10-18 J for 1 electron 5 f = c/λ = 3 x 108/9.11 x 10-8 = 3.29 x 1015 Hz
    • Continuous Spectrum Vs Line Spectrum Emission Line Spectrum • Spectrum with discrete wavelength/ frequency • Excited electrons drop from higher to lower energy level Continuous Spectrum Light spectrum with all wavelength/frequency Excellent simulation on emission spectrum Click here to view excellent simulation Click here to view simulation Emission line spectrum for different elements Click here spectrum for diff elements Click here spectrum for diff element Click here to view simulation Video on quantum mechanics Click here on quantum mechanic, structure of atom
    • Acknowledgements Thanks to source of pictures and video used in this presentation Thanks to Creative Commons for excellent contribution on licenses http://creativecommons.org/licenses/ Prepared by Lawrence Kok Check out more video tutorials from my site and hope you enjoy this tutorial http://lawrencekok.blogspot.com