1. Chapter 5 Notes Electron Models

2. EMR and Quantum Theory

3. How did scientist find out about energy levels and sublevels? When we previously found the electron configuration for elements, it was for electrons at ground state (the lowest energy possible).

4. How did scientist find out about energy levels and sublevels? As energy is added to atoms, they absorb the energy by electrons going from ground state to an excited state, where electrons are no longer in the lowest energy orbitals.

5. How did scientist find out about energy levels and sublevels? Electrons can then only go back to ground state by releasing the energy, usually in the form of light in discreet packets called photons.

6. How did scientist find out about energy levels and sublevels? If electrons could orbit anywhere, when they went from excited state to ground state they would emit light of all wavelengths. This doesn’t happen! Only certain wavelengths of light are emitted, which is different for each atom.

7. How did scientist find out about energy levels and sublevels? This is called an atomic emission spectra, and is different for every element!

8. How did scientist find out about energy levels and sublevels? A scientist named Max Planck studied the cooling of metal and how its color changes, and tied the idea of frequency of light to energy. The bands of light correspond to the specific energy levels.

9. Obj. 1…Light Through a Prism W White light (sunlight) is a blend of all colors (ROY G BIV) combined together. T The wavelength (λ) and frequency (υ) for each color are unique to that color. A As light passes through a prism… - the different wavelengths of the colors are separated. - individual colors can be detected by the eye. - a rainbow appears.

10. Obj. 2…The Electromagnetic Spectrum ll substances (radioactive or not) emit electromagnetic radiation. light: O Only part of the spectrum that human eyes can detect is visible (ROY G BIV) A All other radiations have wavelengths that are either too long or too short for our eyes to detect.

11. Obj. 2 cont… Cosmic rays Microwaves Increasing danger too long to see too short to see 7.4 x 10-6 m longest 4.2 x 10-6 m shortest

12. Obj. 3…Wavelength vs. Frequency W Wavelength (λ): distance b/n crests of a wave. Wavelength crest trough Wavelength long wavelength short wavelength

13. *** higher frequency = higher energy *** Obj. 3 cont… F Frequency (ν): # of wavelengths that pass a certain point in a given amount of time. - units are Hertz (Hz) T These 2 waves are traveling at = speeds… which wave will have more crests cross the ‘finish line’ in a matter of one min.? FINISH low frequency high frequency = short wavelength = long wavelength

14. Obj. 4…Wave Calculations w wavelength (λ) and frequency (υ) are inversely related. a all waves on the EM spectrum travel at the speed of light (c). c = λ(ν) c c = speed of light = 3.00 x 108 m/s. t to solve for λ … ν ν c λ = ν o solve for υ … c ν = λ λ λ

15. Obj. 4 cont… P Practice… Calculate the υ of a wave that has a wavelength of 5.00 x 10-6 m. c ν = λ ν = 3.00 x 10 8 5.00 x 10 -6 ν 6.00 x 10 13 Hz = What is the λ of radiation with a frequency of 1.50 x 1013 Hz? c λ = ν λ = 3.00 x 10 8 1.50 x 10 13 λ = 2.00 x 10 -5 m Does this radiation have a shorter or longer λ than red light? longer…red light ~ 7.4 x 10-6 m

16. Obj. 4 cont… M Max Planck explained that light is emitted in packets (quanta) called ‘photons’ which are distinct bundles of energy. P Planck also assumed that energy of a photon is directly proportional to the frequency of the light. E = h(ν) E E = energy (in Joules) υ υ = frequency h h = Planck’s constant (6.626 x 10-34 J/Hz) P Practice… C Calculate the energy of radiation with a υ of 5.00 x 1015 Hz. E = h(ν) E = 6.626 x 10-34 (5.00 x 1015) E = 3.31 x 10 -18 J 3.31 x 10 -18 J ν h h = E h

17. Obj. 5…Atomic Spectra a all elements will emit light when excited (i.e. by electricity). toms absorb energy and then emit an = amount of energy in the form of light. - atoms emit a characteristic color - Ne = orange - red - Na = bright yellow i if we pass this light through a prism (separate the λ) we get an atomic emission spectrum. ex. of wavelengths emitted set-up

18. Obj. 5 cont… e emission spectra are unique to particular elements. o only show certain lines of the continuous spectrum (white light). h have helped us gather a lot of info. about our universe! a atomic absorption spectra shows colors missing from the continuous spectrum (missing λ were absorbed by the element). continuous absorption emission

19. Evolution of Electron Models The first model of the electron was given by J.J. Thompson—the electron’s discoverer. His was the “plum pudding” model.

20. The Rutherford Model With Rutherford’s discovery of the nucleus of an atom, the atomic model changed.

21. The Bohr Model Niels Bohr introduced his model, which answered why electrons do not fall into the nucleus. He introduced the concept of energy levels, where the electrons orbited similar to the way the planets orbit the sun.

22. Bohr Model and Energy Levels In the Bohr model, electrons are in energy levels, or regions where they most probably are orbiting around the nucleus. The analogy is that energy levels are like the rungs of a ladder—you cannot be between rungs, just like an electron cannot be between energy levels. A quantum of energy is the amount of energy it takes to move from one energy level to the next.

23. Quantum Mechanical Model In 1926, Erwin Schrodinger used the new quantum theory to write and solve mathematical equations to describe electron location.

24. The Quantum Mechanical Model, cont. Today’s model comes from the solutions to Schrodinger’s equations. Previous models were based on physical models of the motion of large objects. This model does not predict the path of electrons, but estimates the probability of finding an electron in a certain position. There is no physical analogy for this model!

25. Where are the electrons? In an atom, principal energy levels (n) can hold electrons. These principal energy levels are assigned values in order of increasing energy (n=1,2,3,4...). Within each principal energy level, electrons occupy energy sublevels. There are as many sublevels as the number of the energy level (i.e., level 1 has 1 sublevel, level 2 has 2 sublevels, etc.)

26. Where are the electrons? There are four types of sublevels we will talk about—s,p,d and f. Inside the sublevel are atomic orbitals that hold the electrons. Every atomic orbital can hold two electrons. S has 1 orbital, P has 3, D has 5 and F has 7. How many electrons can each one hold? Levels: 1, 2, 3, 4... Sublevels: s, p, d, f Orbitals: 1 for s, 3 for p, 5 for d, 7 for f. Electrons: 2 for each orbital.

27. Orbital Shapes s orbital = s sublevel (spherical) + + p x orbitalp y orbital p z orbital = p sublevel (peanut)

28. D orbitals – fyi only (daisy)

29. Where are the electrons? So how many electrons can each energy level hold? –Level 1 has an s sublevel=2 e - –Level 2 has an s and a p sublevel=8e - –Level 3 has an s, p and d sublevel=18e - –Level 4 has an s, p, d and f sublevel=32e -

30. Electron Configuration

31. In the atom, electrons and the nucleus interact to make the most stable arrangement possible. The ways that electrons are arranged around the nucleus of an atom is called the electron configuration.

32. Aufbau Principle Electrons occupy orbitals of the lowest energy first. Pauli’s Exclusion Principle An atomic orbital may have no more than two electrons. Hund’s Rule When electrons occupy the same kind of sublevel, one electron goes in each orbital before the second one goes in.

33. Pauli Exclusion Principle –Each orbital can hold only TWO electrons with opposite spins.

34. Aufbau Principle “Sports Spectator Rule” (fill the lower stands first) Electrons fill the lowest energy orbitals first.

35. RIGHT WRONG Hund’s Rule “Empty Bus Seat Rule” Within a sublevel, place one electron per orbital before placing a second electron.

36. Heisenberg Uncertainty Principle Heisenberg concluded that it is impossible to make any measurement on an object with out disturbing the object (at least a little). The principle states: “It is fundamentally impossible to know precisely both the velocity and the position of a particle at the same time.”

37. O 8e - »Orbital Diagram Electron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

38. Abbreviated S 16e - S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Unabbreviated

41. © 1998 by Harcourt Brace & Company s p d (n-1) f (n-2) 12345671234567 6767 C. Periodic Patterns

42. s-block1st Period 1s 1 1st column of s-block C. Periodic Patterns Example - Hydrogen

43. C. Periodic Patterns Shorthand Configuration –Core e - : Go up one row and over to the Noble Gas. –Valence e - : On the next row, fill in the # of e - in each sublevel.

44. [Ar]4s 2 3d 10 4p 2 C. Periodic Patterns Example - Germanium

45. Electron Configuration Exceptions –Copper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10 D. Stability Cr - [Ar]4s1 3d5 Cu - [Ar]4s1 3d10

46. Valence Electrons Electrons in the atom’s outermost orbitals. These are the electrons that determine the atom’s chemical properties. The fewer valence electrons an atom holds, the less stable it becomes and the more likely it is to react.

47. Octets When an atom has 8 electrons in its largest energy level it is said to have an octet. This is the most stable and least reactive atom.

48. Lewis Dot Structures An atom’s dot structure consists of the element’s symbol, (represents the atomic nucleus and inner-level electrons), surrounded by dots representing the atom’s valence electrons. G. N. Lewis devised this method while teaching a college chemistry class in 1902. Example: Lithium Atomic Number: 3 Electron Configuration: 1s 2 2s 2 Electron Dot Structure: Li·

49. Practice Carbon: Fluorine: Neon: