1. Chapter 7 Beyond Rutherford to “The Most Successful Theory of the 20 th Century”

2. http://www.learnerstv.com/animatio n/chemistry/ruther14.swf View Rutherford’s experiment

3. Rutherford’s Atomic Model (“planetary model”) e- Nucleus Diameter = 10 -15 m Empty space Orbiting electron (fixed radius) Diameter of atom = 10 -10 m

4. Problem with Rutherford’s Model !! It did not obey the classical laws of physics It did not obey the classical laws of physics But atoms don’t collapse, yet Rutherford’s experiment showed that electrons can be located a distance away from the nucleus. But atoms don’t collapse, yet Rutherford’s experiment showed that electrons can be located a distance away from the nucleus. So, the model of the _______ behavior is flawed. So, the model of the _______ behavior is flawed. According to Newton’s classical laws, electrons orbiting the nucleus should radiate energy, slow down, and be pulled into the nucleus & collapse the atom

5. Collision of Ideas Dalton Thomson Rutherford Newton Maxwell Plank Einstein Matter Light ? Bohr & de Broglie

6. What is the nature of light?

7. Isaac Newton: “Light is a particle” Newton’s prism By the 17th century, light was found to travel in straight lines travel in straight lines reflect & refract reflect & refract transmit energy from one place to another transmit energy from one place to another

8. The WAVE THEORY, advocated by Robert Hooke Christian Huygens argued that light is a wave. The PARTICLE THEORY, advocated by Isaac Newton Isaac Newton and Pierre Laplace, argued that Isaac Newton and Pierre Laplace, argued that light was made up of a stream of tiny particles (“corpuscles”).Pierre Laplace Isaac Newton Pierre Laplace

9. Two Competing Theories The theory of light The theory of light

10. “white light” light energy composed of a continuous spectrum of visible electromagnetic radiation ultraviolet infrared

11. Wavelength = distance between wave crests (m) Frequency = cycles per second (Hz) Basics of wave theory

12. Electromagnetic Wave Theory (1865) Electromagnetic waves have a variety of wavelengths, but all travel at the speed of light, Based on conservation of energy, Maxwell derived the wave equation,  Based on experiments of Michael Faraday Michael FaradayMichael FaradayMichael Faraday  Theory developed by James Clerk Maxwell James Clerk MaxwellJames Clerk MaxwellJames Clerk Maxwell c = c = c = 2.998 × 10 8 m/s

13. Electromagnetic Spectrum 10 20 Hz 10 14 Hz10 10 Hz 10 -6 nm 10 8 nm higher energy lower energy

14. ROY G BIV ROY G BIV low energy  high energy Colors in the visible spectrum: Red, Orange, Yellow, Green, Blue, Indigo & Violet

15. Problems with the Wave Theory of Light 1. Blackbody Radiation 2. The Photoelectric Effect 3. Emission Spectra of Atoms By the mid-1800s, the wave theory became predominant, but…… When light interacted with matter, the wave theory failed. The important examples are:

16. Problem #1. Blackbody Radiation blackbody “object that absorbs all the colors in the spectrum” When heated to a high enough temperature, the blackbody radiates white light. The wave theory predicts a continuous spectrum of emitted light, but the theory fails to match experiment. Blackbody Simulation Blackbody Simulation actual spectrum

17. Planck’s Quantum Theory Measured blackbody radiation did not produce a continuous spectrum, as wave theory predicted Measured blackbody radiation did not produce a continuous spectrum, as wave theory predicted In 1900, German Physicist Max Planck proposed a new quantum theory of light: In 1900, German Physicist Max Planck proposed a new quantum theory of light: Light is taken up and given off by a blackbody not as a continuous wave, but in little “packets” of light energy of specific values Planck called these packets “quanta” (singular is quantum) of energy

18. Quantum Theory of Light and Quantum Physics Plank’s quantum theory of light was a historical turning point in physics, transitioning classical physics from the 18 th and 19 th centuries to the quantum physics of the 20 th century. Plank’s quantum theory of light was a historical turning point in physics, transitioning classical physics from the 18 th and 19 th centuries to the quantum physics of the 20 th century.

19. Problem #2 Photoelectric Effect Animation

20. Problem #2. Photoelectric Effect Imagine shining light of various wavelengths (energies) on the surfaces of different metals Imagine shining light of various wavelengths (energies) on the surfaces of different metals Only light energies above a certain threshold cause electrons to be ejected from the metal surface Only light energies above a certain threshold cause electrons to be ejected from the metal surface This conflicts with predictions of the wave theory This conflicts with predictions of the wave theory Animation

21. Einstein’s Photons In 1905, a Swiss patent clerk proposed that light consists of particles called photons. In 1905, a Swiss patent clerk proposed that light consists of particles called photons. As Planck proposed, Einstein’s photons have a certain quanta of energy (based on wavelength) As Planck proposed, Einstein’s photons have a certain quanta of energy (based on wavelength) His model of light solved the problem of the photoelectric effect. His model of light solved the problem of the photoelectric effect. Duality of Light Wave behavior Wave behavior Particle behavior Particle behavior

22. Solar Sail (based on Einstein’s photon theory) - Light reflecting off a mirror imparts momentum - Yet light has no mass (experiment by Compton in 1923) Cosmos 1 concept

23. Energy of Photons At a specific frequency (or wavelength) photons possess a specific quantity of energy (E ) At a specific frequency (or wavelength) photons possess a specific quantity of energy (E ) Planck’s constant Planck’s constant E = h E = h h = 6.626 x 10 -34 J·s E = h c/ E = h c/ Question: Is 400 nm light (violet light) more or less energetic than 750 nm light (red light)?

24. Concept Check The energy required to dislodge electrons from sodium metal via the photoelectric effect is 275 kJ/mol. What wavelength (in nm) has sufficient energy per photon to dislodge an electron from the surface of sodium? sodium

25. Concept Check Which photons have the highest energy? A) Cell phone operating at 1900 MHz B) A laser pointer using 635 nm light

26. Problem #3. Atomic Line Spectra Periodic Table of Line Spectra Periodic Table of Line Spectra Periodic Table of Line Spectra Periodic Table of Line Spectra Flame tests http://college.cengage.com/chemistry/ general/ebbing/general_chem/9e/assets/ instructors/protected/videos.html#Chapter 7

27. Problem #3. Atomic Line Spectra Periodic Table of Line Spectra Periodic Table of Line Spectra Periodic Table of Line Spectra Periodic Table of Line Spectra Emission spectra for pure elements Fireworks

28. Niels Bohr (1885-1962) Danish physicist who worked with J.J. Thomson at Cambridge University in 1911. He didn’t agree with Thomson’s atomic model, Danish physicist who worked with J.J. Thomson at Cambridge University in 1911. He didn’t agree with Thomson’s atomic model, so worked for Rutherford in 1912. In 1912, in a bold step, he suggested that the classical laws of physics cannot be applied to matter as small as atoms and electrons. Instead, new laws are needed In 1912, in a bold step, he suggested that the classical laws of physics cannot be applied to matter as small as atoms and electrons. Instead, new laws are needed Bohr sought to solve the problem with Rutherford’s atomic model and explain the phenomenon of atomic spectra, by applying the quantum theory of light to atoms and electrons Bohr sought to solve the problem with Rutherford’s atomic model and explain the phenomenon of atomic spectra, by applying the quantum theory of light to atoms and electrons

29. Bohr’s Quantum Atomic Model Postulated that the energy of the electron must be quantized. Only certain electron energies are possible. Postulated that the energy of the electron must be quantized. Only certain electron energies are possible. Orbit radii (energy levels) correspond to definite energies Orbit radii (energy levels) correspond to definite energies Energy is emitted or absorbed by the electron only as the electron changes from one allowed energy level to another Energy is emitted or absorbed by the electron only as the electron changes from one allowed energy level to another n = n = energy level number or principal quantum number Why does an electron possess energy? 1)2)

30. How do quantized energy levels explain spectral lines? Atoms “place” electrons in lowest possible energy levels (“ground state”) Atoms “place” electrons in lowest possible energy levels (“ground state”) When electrons are provided with enough energy, they “jump” to higher energy levels, where they are unstable (“excited state”) When electrons are provided with enough energy, they “jump” to higher energy levels, where they are unstable (“excited state”) The electrons then fall back down to the lower possible energy levels, releasing absorbed energy as a photon of light The electrons then fall back down to the lower possible energy levels, releasing absorbed energy as a photon of light We see these photons as the spectral lines emitted by excited atoms We see these photons as the spectral lines emitted by excited atomsStairstepanalogy Energy of H electron = E = -R H /n 2 n = 1, 2, 3, … ∞ R H = 2.179 x 10 -18 J

31. energy levels

32. “quantum jump” ∆E 4→2 =  E 2 - E 4  = h 4→2 H emission spectrum

33. A “quantum jump” Emission ∆E =  E 2 - E 4  = h 4→2Absorption ∆E =  E 4 – E 2  = h 2→4

34. Simulations of Bohr Model Visible emission spectral lines of hydrogen Visible emission spectral lines of hydrogen Visible emission spectral lines of hydrogen Visible emission spectral lines of hydrogen

35. Success & Limitation of Bohr’s Quantum Model Explained the existence of spectral lines Explained the existence of spectral lines Solved the problem with Rutherford’s model of the hydrogen atom Solved the problem with Rutherford’s model of the hydrogen atom But, the mathematics only worked for atoms with 1 electron! But, the mathematics only worked for atoms with 1 electron! How can this model be made to work for all elements?

36. de Broglie’s Novel Notion Light was “known” (thought) to be a wave, but Einstein showed that it also acts particle-like. Electrons were “known” to be particles mass & charge. French physicist: What if …… What if …… 1923 electrons behaved as waves also Diffraction pattern obtained by firing a beam of electrons through a crystal.

37. Dr. Quantum video Dr. Quantum video

38. Werner Heisenberg In 1927, German physicist, proposed that the dual nature of the electron places limitations on how precisely we can know both the location and speed of the electron In 1927, German physicist, proposed that the dual nature of the electron places limitations on how precisely we can know both the location and speed of the electron Instead, we can only describe electron behavior in terms of probability Instead, we can only describe electron behavior in terms of probability The Uncertainty Principle speed position

39. Heisenberg’s Uncertainty Principle Wave behavior limits what can be known! What if the particle has a small mass? What if the particle has a small mass? What if the electron’s position is known very precisely? What if the electron’s position is known very precisely? What if the electron’s speed is known very precisely? What if the electron’s speed is known very precisely? (±x)(±v x )  Can the electron’s orbit be precisely defined? ± position± speed h h4m4m h h4m4m

40. Erwin Schrodinger In 1926, Austrian physicist, proposed an equation that incorporates both the wave and particle behavior of the electron In 1926, Austrian physicist, proposed an equation that incorporates both the wave and particle behavior of the electron When applied to hydrogen’s 1 electron atom, solutions provide the most probable location of finding the electron in the first energy level When applied to hydrogen’s 1 electron atom, solutions provide the most probable location of finding the electron in the first energy level Can be applied to more complex atoms too! Can be applied to more complex atoms too! Wave Equation & Wave Mechanics

41. Extremely small mass Located outside the nucleus Moving at very high speeds Have specific energy levels Standing wave behavior Extremely small mass Located outside the nucleus Moving at very high speeds Have specific energy levels Standing wave behavior Electron Characteristics

42. A baseball behaves as a particle and follows a predictable path. BUT An electron behaves as a wave, and its path cannot be predicted. All we can do is to calculate the probability of the electron following a specific path. Baseball v. Electron

43. What if a baseball behaved like an electron? Characteristic wavelength ( ) baseball  10 -34 m electron  0.1 nm So, all we can predict is….. = h /(mu) = h /(mu) mass speed

44. “deterministic” “probabilistic”

45. Bohr Model v. Quantum Mechanics Energy Electron Position/Path Elements Bohr Quantum Mechanics

46. The electron's movement cannot be known precisely. We can only map the probability of finding the electron at various locations outside the nucleus. The probability map is called an orbital. The orbital is calculated to confine 99% of electron’s range. Energy of the electron is quantized into sublevels. Quantum Mechanics Model

47. Quantum Mechanics Model Describes the energy, arrangement and space occupied by electrons in atoms Quantum Mechanics Electron’s energy is quantized Mathematics of waves to define orbitals (wave mechanics)

48. “Most Successful Theory of the 20 th Century” Dalton Thomson Rutherford Newton Maxwell Plank Einstein Matter Light Schrödinger Heisenberg Wave Mechanics Quantum Mechanics Bohr & de Broglie