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

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

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

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

4. Problem with Rutherford’s Model !! • 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. • 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 Matter Dalton Thomson Rutherford ? Bohr & de Broglie Einstein Plank Maxwell Newton Light

6. What is the nature of light?

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

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

9. Two Competing Theories • The theory of light • The theory of light WAVE Particle

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

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

12. Electromagnetic Wave Theory (1865) • Based on experiments of Michael Faraday • Theory developed by James Clerk Maxwell 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, c=2.998×108 m/s c=

13. Electromagnetic Spectrum 1020 Hz 1014 Hz 1010 Hz 10-6 nm 108 nm higher energy lower energy

14. 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 By the mid-1800s, the wave theory became predominant, but…… When light interacted with matter, the wave theory failed. The important examples are: • Blackbody Radiation • The Photoelectric Effect • Emission Spectra of Atoms

16. Problem #1. Blackbody Radiation blackbody “object that absorbs all the colors in the spectrum” Blackbody Simulation actual 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.

17. Planck’s Quantum Theory • 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: • 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 18th and 19th centuries to the quantum physics of the 20th century.

19. Problem #2Photoelectric Effect Animation

20. Problem #2. Photoelectric Effect • 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 • 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. • 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. Duality of Light • Wave 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) • Planck’s constant E=h E=hc/ h=6.626x10-34 J·s 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 Flame tests http://college.cengage.com/chemistry/ general/ebbing/general_chem/9e/assets/ instructors/protected/videos.html#Chapter 7 • Periodic Table of Line Spectra

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

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, 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 • 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. • 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 n= energy level number or principal quantum number Why does an electron possess energy? 1) 2)

30. Stairstep analogy How do quantized energy levels explain spectral lines? • 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”) • 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 Energy of H electron = E = -RH/n2 n = 1, 2, 3, … ∞ RH = 2.179 x 10-18 J

31. energy levels

32. H emission spectrum “quantum jump” ∆E4→2 = E2 - E4= h4→2

33. A “quantum jump” Emission ∆E = E2 - E4= h4→2 Absorption ∆E = E4 – E2= h2→4

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

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

36. 1923 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 …… electrons behaved as waves also Diffraction pattern obtained by firing a beam of electrons through a crystal.

37. Dr. Quantum video

38. Werner Heisenberg The Uncertainty Principle • 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 speed position

39. h 4m ± speed ± position Heisenberg’sUncertainty Principle Wave behavior limits what can be known! • What if the particle has a small mass? • What if the electron’s position is known very precisely? • What if the electron’s speed is known very precisely? (±x)(±vx) Can the electron’s orbit be precisely defined?

40. Erwin Schrodinger Wave Equation & Wave Mechanics • 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 • Can be applied to more complex atoms too!

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

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

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

44. “deterministic” “probabilistic”

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

46. Quantum Mechanics Model 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 anorbital. The orbital is calculated to confine 99% of electron’s range. Energy of the electron is quantized into sublevels.

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

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