Electrons in Atoms

1. Electrons in Atoms

2. Problems with Rutherford’s Model (5-1) p. 127 Chlorine # 17 Reactive Argon # 18 Not reactive Potassium # 19 Very reactive Q. 22 P. 149

3. The Quest for a Better Model (5-3)P. 138-140 • Electromagnetic radiation behaves like a wave. (video)

4. Characteristics of a Wave Wavelength = λ Frequency = v (number of waves that pass a point per second) 1 Hertz (Hz) = 1 wave per second (SI Unit for frequency) P. 149 - Q 41 and 42

5. Speed and Frequency of Light c = λv c = speed of light (3.0 x 108 m/s) ↑ wavelength ↓ frequency ↓ wave length ↑frequency P. 149 - 45

6. What is the relationship between energy and frequency?

7. Problems An FM radio station broadcasts at a frequency of 98.5 MHz. What is the wavelength of the station’s broadcast signal? A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light? P. 150 - Q 55, 56, P. 151 – Q 74

8. Light: Particle or Wave? Wave model doesn’t address: Why heated objects emit only certain frequencies of light at a given temperature? Why some metals emit electrons when a colored light of a specific frequency shines on them?

9. Iron • Dark gray = room temp • Red = hot temp • Blue = extremely hotter temp • ↑ temp, ↑ kinetic energy, emit different colors of light • Wave model could not explain this

10. Max Planck – 1900P. 128 Matter gains or loses energy only in small, specific amounts called quanta quantum is the minimum amount of energy that can be gained or lost by an atom Equantum = hv h – Planck’s constant – 6.626 x 10-34 J·s J = joule, SI Unit for energy

11. Photoelectric Effect – The problem with wave theory. - Simulation Only certain frequencies of light could emit an electron from a plate of Ag. Accumulation of low frequencies couldn’t

12. Einstein and the Dual Nature of EMR (1900) P. 144 • EMR acts as a wave of individual particles (photon) Ephoton = hv

13. Calculating the energy in a Photon Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 s-1? Ephoton = hv E = (6.626 x 10-34 J·s) x (7.23 x 1014 s-1) E = 4.79 x 10-19 J P. 151 Q - 73 a,b,d - Q 76

14. Atomic Emission Spectra P.141 The frequencies of the EMR emitted by atoms of the element. Unique to each element

15. Niels Bohr – 1913(5-1)P. 128 • Worked in Rutherford’s lab • Proposed a quantum model of the atom • Explain why emission spectra were discontinuous • Predicted frequencies of light in Hydrogen’s atomic emission spectra

16. Bohr’s Explanation • Ground state – lowest energy state of an atom • Excited state – when an atom gains energy • Electrons move in circular orbits • Smaller orbit – lower energy state, “energy level” • Larger orbit – higher energy state, “energy level” P. 149 - Q - 23, 24

17. An explanation for the Emission Spectra Atoms absorb energy and are excited. As the electron returns to the ground state they give off energy “photon” equal to the difference in energy levels. Excited States Ground state

18. photon Electron absorbs energy e- Electron in ground state P. 149 - Q 47

19. Energy released “photon” Electron in Excited State e- Returns to Ground state

20. Problem: Bohr’s Model Only explains Hydrogen P. 144 • Louis de Broglie (1924) – proposed that the energy levels are based on the wave like nature of electrons

21. Heisenberg Uncertainty Principle p.145 • It is impossible to know the velocity and position of a particle at any given time (Video) Photon and electron are about the same mass.

22. Erwin Schrodinger – 1926 P. 130 • Developed the quantum mechanical model of the atom • Assigns electrons to energy levels like Bohr • Does not predict the path of the electron • It predicts the probability of finding an electron • An electron’s “atomic orbital” (Video – Quantum Atom)

23. Each dot is a picture of an electron during a given amount of time. Where does the electron spend most of the time? Boundary represents the location of an electron 90% of the time. P. 149 - Q 25

24. Principle Energy Levels p. 131 • 7 energy levels • Lowest energy is 1 – greatest energy 7 • Each level consists of sublevels The second energy level is larger and the electrons are farther from the nucleus.

25. Types of Sublevels p. 131 Same energy

26. Putting it together: P. 149 Q - 26, 27, 29, 35, 36, 37

27. Ground State Electron Configurations p. 133 • Most stable – lowest energy • 3 principals to follow

28. Aufbau Principle – each electron must occupy the lowest energy state p. 133 • Orbitals in an energy sublevel have equal energy • The energy sublevels in a principle energy level have different energies. • The sublevels increase in energy from s,p,d,f • Principal energy levels can overlap

29. Aufbau Diagram Energy levels overlap Equal energies – 2 p Sublevels have different energy levels

30. Pauli Exclusion Principle p. 134 • Each electron spins • Electrons must spin in opposite directions • 2 electrons per orbital Written as

31. Hund’s Rule p. 134 Electrons must occupy each orbital before additional electrons can be added. Pauli Exclusion Principle P. 151 Q - 68

32. Representing Electron Configurations P. 134

33. Electron Configurations Sub level diagram – indicates the order that orbitals are filled What are the orbital diagrams and electron configuration notation for Al and Cl? P. 149 Q - 32ab, 34abcd, 38ab

34. Electron Configuration Shorthand • Substitute noble gases from preceding energy levels in the notation Li – [He] 2s1 C – [He] 2s2 2p2 P. 207 – Q 35

35. Valence Electrons • Electrons in the outer most energy levels S [Ne] 3s2 3p4 Sulfur has 6 valence electrons How many valence electrons do Al, Ne, and Cl have?

36. Electron Dot Structures Valence electrons are used in reactions and are represented by an electron dot structure.

37. Writing Electron Dot Structures • Fill the valence electrons 1 at a time in any particular order. Ca C O * * * * * * * * * * * * What are the electron dot diagrams for K, Ar and F?