1. Electrons

2. Light and Quantized Energy Electrons part 1

3. The spectrum and refraction 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.

4. 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. (Goldilocks!) EMS

5. 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

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

7. *** higher frequency = higher energy *** 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 Cont…

8. 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

9. 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

10. 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 Emissions of light by atoms

11. 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 Continuous vs. emission

12. e e- are found on certain energy levels (orbitals) around the atom. - there is a maximum of seven energy levels in an atom. - an e- requires one ‘quanta’ of energy to jump to the next energy level. - e- on the energy level closest to the nucleus have the lowest energy. The 7th energy level has the highest energy. - at their lowest energy level are considered to be at the ground state (most stable). i if e- absorb a quantum or more of energy (from electricity), they can jump to higher energy levels (excited state). e e- must lose energy in order to fall from the excited state back to the ground state. - this energy is emitted in the form of visible light! Quantum of energy

13. Bohr’s model of the atom was able to correctly explain why line spectra results when atoms are heated. The energy of each orbital can be calculated using Bohr’s equation. E n = -(2.18 x 10 -18 J) / n 2 Where n = (1,2,3,ect…)

14. Quantum Theory and the Atom Electrons part 2

15. Quantum Numbers t there are four quantum numbers... he ‘address’ of an e-. Will look like this { 1, 1, 0, +1/2) 1st = n = energy level of the e-. Possible values are (1,2,3,etc…) - 1 = lowest, 7 = highest 2nd = l = shape of sub-orbital - 4 different shapes…found on each energy level - ‘the principle quantum number’ 3rd = ml = orientation of orbital in space. - which axis the orbital lies on 4th = ms = spin of the e- on its axis. - clockwise or counter-clockwise - represented by rows on the periodic table

16. Principle energy level: Maximum Number of Electrons he formula 2n2 is used to find the max. # of e- on any principle energy level. - n = the energy level = (row on P.T.) E Ex… e energy level 3… 2(3)2 = 18 e- nergy level 6… 2(6)2 = 72 e- - Corresponds to a ring on the Bohr Model. Each ring is principle energy. { 1, 1, 0, +1/2} First ring on Bohr Model from our first example.

17. The 2nd Quantum # (sublevels) l l = shape of orbital. Possible values: 0,1,2,3. n = # of possible l t there are four orbital shapes, represented by letters… ( (0) s = sphere - lowest energy e each orbital can only hold 2 e-. = 1 orbital = 2 total e- ( (1) p = dumbbell = 3 orbitals = 6 total e- 2) d = clover-leaf= 5 orbitals = 10 total e- 3) f = double clover-leaf= 7 orbitals= 14 total e- - highest energy { 1, 1, 0, +1/2) Tells us that the electron is in the p orbital.

18. The 3 rd Quantum number Called the magnetic quantum # (m l ) Possible values: between ( l to - l ) Possible values: between ( l to - l ) Ex. If 2 nd quantum number is 1 then all possible values for 3 rd quantum number is: -1, 0, +1 This is the orientation of orbital around the axis. This is the orientation of orbital around the axis. { 1, 1, 0, +1/2) Tells us the e- is in the second of three orbitals.

19. The 4 th quantum number The spin of the electron. (m s ) Can have two possible spins: clockwise, counter clockwise. Can have two possible spins: clockwise, counter clockwise. We will use +½ for clockwise. We will use +½ for clockwise. We will use -½ for counterclockwise. We will use -½ for counterclockwise. { 1, 1, 0, +1/2) Tells us the e- has a clockwise spin.

20. cont… ‘s’ orbital ‘s’ orbital ‘p’ orbitals ‘p’ orbitals 2 nd quantum number designates shape of orbitals. And each can have 2e- 3 rd quantum # is orientation. 1 st quantum number tells how large the orbital. 4 th quantum number tells the spin of e-

21. cont… cont… ‘d’ orbitals ‘d’ orbitals ‘f’ orbitals ‘f’ orbitals How many e- can each orientation have? Answer: 2e- How many total e- can the d orbitals have? 10e-

22. Principle (n) 2 nd Quant. # (sublevel l ) # of orbitals of sublevels. (- l to + l ) Total # of orbitals of (n) 1 s = 0 1 (0) 1 2 s = 0 p = 1 1 (0) 3 (-1, 0, +1) 4 3 s = 0 p = 1 d = 2 1 (0) 3 (-1,0,+1) 5 (-2,-1,0,+1,+2) 9 4 s = 0 p = 1 d = 2 f = 3 1 (0) 3 (-1,0,+1) 5 (-2,-1,0,+1,+2) 7 (-3,-2,-1,0,+1,+2,+3) 16

23. Questions about address of electron. If n=3, can l = 1? Yes Yes If n=2, can l = 2? What are the possible values of the sublevels in the 3 principle energy level? 0,1,2 ( n=3 so 3 possible and start with 0) 0,1,2 ( n=3 so 3 possible and start with 0) How many orbitals does the p sublevel have? 3 (-1,0,+1) 3 (-1,0,+1)

24. Electron Configuration Electrons part 3

25. Aufbau Principle States that each electron occupies the lowest energy orbital available. All orbitals related to an energy sublevel are of equal energy. Ex. All “p” orientations have same energy. All orbitals related to an energy sublevel are of equal energy. Ex. All “p” orientations have same energy. The energy sublevels within a principle energy level have different energies. Ex. The 2s orbitals have less energy than the 2p orbitals. The energy sublevels within a principle energy level have different energies. Ex. The 2s orbitals have less energy than the 2p orbitals. The sequence of energy is: s (lowest),p,d, then f (highest). The sequence of energy is: s (lowest),p,d, then f (highest).

26. Pauli’s Exclusion Principle P Pauli stated that no two e- in the same atom will have the same set of quantum #s. o only 2 e- can fill an orbital - each e- has a different spin (‘s’ quantum #) - even if first 3 quantum #s are the same, it is always the 4th that will define an individual e-. Hund’s Rule H Hund stated that e- will fill empty orbitals 1st before they pair up. ‘ ‘elbow room’

27. The Diagonal Rule1s2s2p 3s3p 3d 4s4p 4d 4f 5s5p5d5f 6s6p 6d 7s 7p p d s f

28. cont… P Pictorial configuration… - draw a line to represent each orbital - follow Aufbau, Pauli, & Hund’s Rule - fill orbitals in singly 1st before you pair e-. - use arrows to represent e-...point up or down to represent spin. E Ex. Use diagonol rule: Give the pictorial notation for B (Z = 5) 1s 1s 2s 2s 2p 2p G Give the pictorial notation for S (Z = 16) 1s 1s 2s 2s 2p 2p 3s 3s 3p 3p 1s2 2s2 2p1 1s2 2s2 2p6 3s2 3p4

29. What is the address of the last electron added? 1s 1s 2s 2s 2p 2p 1s 1s 2s 2s 2p 2p 3s 3s 3p 3p 1s2 2s2 2p1 1s2 2s2 2p6 3s2 3p4 {2, 1, -1, +1/2} {3,1,-1,-1/2}

30. Electron Configurations t the periodic table is laid out in order of the quantum #s. i in e- configurations, you must represent (in order), the: energy level #, orbital (s,p,d or f) and e- # for each e- in the atom. r recall that # of p+ and e- = each other in a neutral atom! m must fill in lowest energy level (1) 1st, and lowest orbital (s) 1st on each energy level.

31. Lets give the electronic configuration of Boron (B). 1s 2 2s 2 2p 1 Boron is here. Lets give the electronic configuration for Fe Fe is here. Look at the diagonal rule; after 4s, then 3d. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

32. P Practice… cont… cont… W Write out the correct e- configuration for Mg (Z = 12) 1s2 2s2 2p6 3s2 rite out the correct e- configuration for P (Z = 15) 1s2 2s2 2p6 3s2 3p3 ‘ ‘d’ orbitals are so large that they reach into the next energy level. t therefore, the 1st ‘d’ orbital belongs to the 3rd energy level…even though we don’t see it until the 4th row of the P.T.! W Write out the correct e- configuration for Br (Z = 35) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 ‘ ‘f’ orbitals are so large that they reach into the next 2 energy levels. t therefore, the 1st ‘f’ orbital belongs to the 4th energy level…even though we don’t see it until the 6th row of the P.T.!

33. Last Energy Level using Noble gas notation T To give the noble gas notation, find the element of interest, then find the previous noble gas and start from that point. Ex… Potassium (K) = Iodine (I) = P Put the noble gas in brackets which will represent all of the electrons up to the noble gas. Nickel (Ni) = [He]2s2 = [Ne]3s2 3p2 = [Kr]5s2 4d9 = Beryllium (Be) Silicon (Si) Silver (Ag) [Ar] 4s 1 [Kr]5s 2 4d 10 5p 5 [Ar] 4s 2 3d 8

34. Electron Dot Diagrams r represents all of the valence e- in an atom. v valence e- are all the e- on the outer-most energy level of an atom. - only ‘s’ and ‘p’ orbitals on highest level!!! - ‘f’ and ‘d’ orbitals are embedded in the other two. - represented by the ‘A’ columns on the P.T. e electron dot diagrams consist of the symbol of the element surrounded by dots for each valence e-. Ex… Iodine (I): Beryllium (Be): 5 s25 p 5 I 2s2 Be [Kr] [He]