1. Arrangement of Electrons in Atoms Chapter 4

2. Section 4.1 Wave-Particle Nature of Light 1. Electromagnetic Radiation -a form of energy that exhibits wavelike behavior as it travels through space - moves at 3.0 x 10 10 cm/s in a vacuum - wavelength ( λ ): the distance between corresponding pointson adjacent waves

3. - frequency ( ν ): the number of waves that pass a given point in a specific amount of time - c = λν where: c is the speed of light 3.0 x 10 8 m/s λ is the wavelength ν is the frequency

5. - continuous spectrum: a spectrum in which all wavelengths within a given range are included - electromagnetic spectrum: consists of all electromagnetic radiation, arranged according to increasing wavelength

7. 2. Light as particles  material does not burn a. glows red b. changes to yellow then white c. visible colors  photoelectric effect : the emission of electrons by certain metals when light shines on them

9. Quantum: a finite quantity of energy that can be gained or lost by an atom Planck 1. proposed a relationship between a quantum of energy and the frequency of radiation 2. E = hv where: E = energy of radiation h = Plank’s constant 6.626 x 10 -34 J  s

10. 3. photon: an individual quantum of light 4. ground state: state of lowest energy of an atom 5. excited state: a state in which it has a higher potential energy atom Ex. When atoms in a gaseous state are heated; the potential energy increases, return, give off the added energy in the electromagnetic radiation

11. Ex. Continued: neon signs and fireworks We can show this on different line spectrum or line- emission spectrum. Each of the colored lines is produced by light of a different wavelength.

13. The different representative lines

14. Bohr Model 1. First model of the electron structure 2. Gives levels where an electron is most likely to be found 3. Incorrect today, but a key in understanding the atom

15.  Section 4.2  Quantum Model of the Atom 1. Describes mathematically the wave properties of electrons and other very small particles 2. dual wave-particle nature of light 3. orbital: a three-dimensional region about the nucleus in which a particular electron can be located

16. Review characteristics of electrons 1. extremely small mass 2. located outside the nucleus 3. moving at extremely high speeds in a sphere 4. have specific energy levels 5. when atoms are heated: bright lines appear 6. arranged in discrete levels 7. absorbs energy to “jump: to a higher energy level 8. falls to a lower level, energy is emitted

17. 9. gain loss of energy of energy

18. Electron levels (shells) Contain electrons that are similar in energy and distance from nucleus Low energy electrons are closest to the nucleus Identify by numbers 1,2,3,4,5,6…… The first shell (1) is lowest in energy, 2 nd level is next 1<2<3<4…..

19. Number of electrons Maximum number of electrons in any electron level = 2n 2 when the level # total electrons n = 12 n = 2 8 n = 3 18

20. Orbitals Orbitals shape # of orbitals # of electrons (sublevels) s sphere 1 2 p dumbbell 3 6 d 4-leaf clover 5 10 f dragon-fly 7 14

21. Electron configuration

22. Notation Longhand configuration: Examples: S 16 e - Ca 20 e -

23. Shorthand configuration: Use the preceding Noble gas in [ ] Examples: S 16 e - Ca 20 e -

24. Hund’s Rule Orbitals of equal energy are each occupied by one electron before any one orbital is occupied by a second electron  Pauli Exclusion Principle No two electrons in the same atom can have the same set of four quantum numbers.

25. Examples: S 16 e - Ca 20 e -

26. Exceptions s 1 : one electron leaves the “s” and goes to the “d” : Nb, Cr, Mo, Tc, Ru, Rh, Cu, Ag, Au, Pt : examples: Cr #24 [Ar] 4s 2 3d 4 unstable __ __ __ __ __ __ [Ar] 4s 1 3d 5 more stable __ __ __ __ __ __

27. Section 4.3 Quantum numbers Four quantum numbers “Specifies the address of each electron in an atom” 1. Principal Quantum Number (n)  Energy level  Size of the orbital

28. 2. Angular Momentum Quantum Number (l)  energy sublevel  shape of the orbital  0 1 2 3 s p d f

29. 3. Magnetic quantum number (m)  Indicates the orientation of an orbital about the nucleus  s ____ p ___ ____ ____ 0 -1 0 +1 d ___ ___ ___ ___ ___ -2 -1 0 +1 +2

30. f ___ ___ ___ ___ ___ ___ ___ -3 -2 -1 0 +1 +2 +3 4. Spin quantum number (s)  Indicates two possible states of an electron in an orbital.   = +1/2  = -1/2

31. Ex. Write the configuration for P #15, determine the quantum numbers for the third and seventh electron. State the element whose last electron has the following quantum number’s n = 5 l = 2 m= 0 s = +1/2

32. Magnetism: 1. Diamagnetism: the property of a substance whereby it is weakly repelled by a magnetic field ex. Full electrons 2. Paramagnetism: a weak attraction between magnetic fields and substances whose atoms have uneven electron distribution ex. Not full electrons

33. Section 4.4 Electron Dot Structure 1. Keeps track of valence electron 2. Valence electrons: outermost electrons 3. Octet rule: has eight valence electrons : it is stable 4. symbol: X

34. Examples: 1. Na #112. Cl #17 3. Ba #56 4. P #15