Chapter

1. Chapter Electrons in the Atom

2. Timeline of the Models* ? *There are many more models. These are the ones we’ll cover in class.

3. John Dalton (c. 1800 AD) • English chemist, John Dalton, performed experiments with various chemicals and showed matter seemed to consist of “indivisible” particles (atoms). • Though he didn’t know about atoms’ structure, he did know about the Law of Conservation of Matter and based his theory on this. Dalton was an avid weather watcher and discoverer of color blindness among other things.

4. Dalton’s Ideas “Billiard Ball” or “Marbles” • Dalton's model says atoms are tiny, indivisible, indestructible particles • Believed each atom had a certain mass, size, and chemical behavior determined by what kind of elements they make up. • See next slide for the details….

5. John Dalton Theory --1800(a.k.a “the marble guy”) • Atoms are the smallest particles of nature-indivisible and indestructible • All atoms of the same element are identical • Atoms of different kinds can combine to form compounds • Chemical reactions are atoms recombining to form new substances

6. J.J. Thomson (c. 1897) • In 1897, English physicist J.J. Thompson discovered the electronand proposed a model for the structure of the atom. • Using a CATHODE RAY TUBE, Thomson discovered electrons have a negative charge and thought that the rest of matter must have a positive charge to offset the negative electron.

7. His Experiment

8. JJ Thompson’s Model“Plum Pudding” • Because the beam of light traveled from to the positive end of the tube he concluded that the light had a negative charge • Because the beam could push a paddle wheel he concluded that the particle had mass. • Thompson's model says atoms are positively charged spheres with negatively charged electrons randomly located throughout.

9. Side Trip…Alpha Particles! • Around this time scientists also discovered alpha rays (particles), which had a positive charge. • Some physicists thought these alpha particles were made up of the positive parts of JJ Thompson’s atom.

10. Ernest Rutherford (1911) • Rutherford as a student worked under J.J. Thompson supervision at the famous Cavendish Laboratories. • In 1911 Ernest Rutherford bombarded atoms with alpha rays to investigate the inside of the atom. • The results were, to say the least, unexpected!

11. Gold Foil Experiment • Rutherford used Radium as the source of the alpha particles and shot them at a thin gold foil like aluminum foil but made of gold • A fluorescent screen sat behind the gold foil on which he could observe the alpha particles’ impact.

12. When the particles bounced back or were deflected Rutherford reasoned that it hit something massive and positive. This mass became know as the nucleus. • When the alpha particles went straight through it hit nothing. • This happened most often so the atom is mostly empty space.

13. Rutherford’s “Planetary Model Rutherford’s model said the negative electrons orbited a positive center (NUCLEUS)like our planets orbit the sun. • The nucleus contained most of the mass of the atom • And the distance between the positive center (nucleus) and the electrons was huge-like a marble in the center of a football field. • The atom was mostly empty space!!!!

14. One little problem… • The theory of electricity and magnetism predicted that opposite charges attracteach other and the electrons should gradually lose energy and spiral inward toward the nucleus.(BOOM! No more atom.)

15. Niels Bohr (1912) • In 1912 a Danish physicist, Niels Bohr came up with a theory that said the electrons do not spiral into the nucleus and came up with some rules for what does happen. • This was a pretty radical approach, because for the first time rules had to fit the observation regardless of how they conflicted with the theories of the time!(Aristotle would have been furious).

16. Previous experiments-White light gives off all wavelengths of energy- all colors.

17. The explanation- • An electron absorbs energy it jumps farther away form the nucleus • As the electron falls back closer to the nucleus it gives off the energy as colored light.

18. How Light Relates to Electron Location • Bohr observed that only certain colors were given off • Therefore the electron could only orbit at certain distances from the nucleus

19. Bohr’s Rules • RULE 1: Electrons can orbit only at certain allowed distances from the nucleus (energy-levels). • RULE 2: An atom absorbs energy when an electron gets boosted from a low-energy orbit to a high-energy orbit. • Rule 3: Atoms radiate energy when an electron jumps down from a higher-energy orbit to a lower-energy orbit.

20. Bohr’s Model of the atom. Light can excite electrons around atoms and this gives rise to “quantum levels”. See light excite an atom!

21. Are We Done Yet? • Almost… Cliff Notes version is that Niels Bohr came really close, and when you add the works of Arnold Sommerfeld, Wolfgang Pauli, Louis de Broglie, Erwin Schrödinger, Max Born, and Werner Heisenberg, we arrive at today’s model…

22. Today’s Model!-Electron Cloud • Today's model says electrons are not confined to fixed orbits. • They occupy volumes of space outside the nucleus.

23. Models of the Atom • Dalton – indivisible sphere • Thomson Model – a ball of positive charge containing a number of electrons “plum pudding” • Rutherford Model – atom has a positively charged nucleus • Bohr – electrons travel around the nucleus in definite orbits (energy levels) • Quantum Mechanical Model – no definite shape and no definite electron orbitals

24. Models of the Atom • Energy level – the region around the nucleus where the electron is likely to be moving (discrete levels, like stair-steps) • Quantum – the amount of energy required to move an electron from its present energy level to the next higher one

25. Models of the Atom • Atomic orbital – cloud shaped regions where electrons are thought to be located • S – orbital = Spherical • P – orbital = Peanut • D – orbital = Double peanut • F – orbital = Far too complex (Flower) • Nodes – regions where the probability of finding an electron is very low (eg: No electrons)

26. Table 5.1 Summary of Principle Energy Levels, Sublevels, and Orbitals

27. Increasing energy (increasing distance from nucleus) Energy Level n 1 2 3 4 Maximum number of electrons allowed 2 8 18 32

28. Electron Arrangement in Atoms • Electron Configurations – the way electrons are arranged around the nucleus of atoms • Aufbau Principle – electrons enter orbitals of lowest energy first • Pauli exclusion Principle – An atomic orbital may describe at most two electrons

29. Electron Arrangement in Atoms • Hund’s rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins • Once all orbitals of equal energy have one electron with parallel spin, the next electron to enter the orbital has the opposite spin.

30. Electron Arrangement in Atoms • Cr = • Cu = SEE OVERHEAD

31. Physics and the Quantum Mechanical Model • Electromagnetic radiation – includes radio waves, microwaves, visible light, infrared light, ultraviolet light, X-rays, and gamma rays

32. Physics and the Quantum Mechanical Model • Amplitude – the height of the wave from the origin to the crest • Wavelength – the distance between the crests • Frequency – the number of wave cycles to pass a given point per unit of time Crest Crest

33. Physics and the Quantum Mechanical Model • Hertz – the metric unit of cycles per second • Spectrum – when sunlight passes through a prism, the light separates into a spectrum of colors • Atomic emission spectrum – passing light emitted by an element through a prism gives a spectrum of the element

34. Physics and the Quantum Mechanical Model • c = ג x v • c = speed of light • v = frequency • ג = wavelength

36. Physics and the Quantum Mechanical Model • Photons – light particles • Photoelectric effect – electrons called photoelectrons are ejected by metals when light shines on them. • Li, Na, K, Cs, and Rb are the best metals to demonstrate the effect.

37. Physics and the Quantum Mechanical Model • Ground State – the lowest energy level (n = 1) • Energy levels above the ground state are denoted N = 2,3,4,5,6,ect… • Electrons absorb energy to move up an energy level and give off energy to move down an energy level