2. Chapter 3 Arrangement of Electrons in the Atom

3. Spectra When white light passed through a prism= Continuous spectrum

4. Hydrogen Spectra Niels Bohr studied light emitted from a Hydrogen discharge tube and observed a spectrum consisting of a series of narrow coloured lines.(emission line spectrum) Use spectrometer or hand held spectroscope.

5. Line Spectra

6. FlameFlame Tests In this experiment we investigate the colours given off by certain salts

7. Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. Include link to web page http://www.unit5.org/christjs/flame%20tests.htm methane gas wooden splintstrontium ioncopper ionsodium ion calcium ion

8. Emission Spectrum of Hydrogen 1 nm = 1 x 10 -9 m = “a billionth of a meter” 410 nm434 nm486 nm656 nm

9. Emission Spectrum of an Element 1 nm = 1 x 10 -9 m = “a billionth of a meter” 410 nm434 nm486 nm656 nm 1 nm = 1 x 10 -9 m = “a billionth of a meter”

10. Continuous and Line Spectra 4000 A o 5000 6000 7000 light Na H Ca Hg 400 450 500 550 600 650 700 750 nm Visible spectrum  (nm)

11. Continuous and Line Spectra 4000 A o 5000 60007000 4000 A o 5000 6000 7000 light Na H Ca Hg

12. Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. Include link to web page http://www.unit5.org/christjs/flame%20tests.htm methane gas wooden splintstrontium ioncopper ionsodium ion calcium ion

13. Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

14. Bohr’s Theory Why is an emission spectrum produced? -Bohr theory 1.Electrons revolve around the nucleus in fixed paths called orbits. 2.Electron has a fixed amount of energy. 3.As long as the electron is in one energy level it neither loses nor gains energy. 4.When the electron absorbs energy- jumps from a lower to a higher energy level. 5.Energy lost as it falls from a higher energy level to a lower energy level. 6.Definite amount of energy lost.

15. Each definite amount of energy emitted gives rise to a line in the emission spectrum. Since only definite amounts of energy are emitted, therefore electrons can occupy only definite energy levels. Light emitted h= Plank’s constant f = frequency of light emitted E 2 -E 1 =hf Line spectra applet

16. Each electron transition results in a line of a particular wavelength. N=1 Lyman series Ultraviolet N=2 Balmer series visible lines N=3 Paschen series Infra Red

17. Limitations to Bohr Theory Limitations of the Bohr Theory It is mainly successful in explaining one electron system e.g. Hydrogen. The Bohr Theory does not explain the splitting of lines in the emission spectra and therefore the account for the existence of sublevels. Doesn’t take into account the wave nature of electrons The presence of atomic orbitals is also not accounted for.

18. Atomic Absorption Spectrometry. Atoms can also absorb light. If white light is passed through a gaseous sample of an element it is found that the light that comes out has certain wavelengths missing. We can analyse and correspond these to various elements.

19. Spectra

21. Atomic absorption spectrometer

22. Differences of spectra An emission spectrum consists of coloured lines against a black background. An absorption spectra consists of dark lines against a coloured background One use of Atomic absorption spectrometry is the analysis of water for metals like lead, mercury. Cadmium(heavy metals)

23. Wave nature of the electron Wave nature of the electron. When the Bohr theory is applied to atoms with more than 1 electron it failed to account for many of the lines in the emission spectrum of these atoms. De Broglei-moving electrons had a wave motion associated with them.

24. Heisenberg’s Uncertainty principle Heisenberg’s Uncertainty Principle.-It is impossible to measure at the same time both the velocity and position of electrons. This gives rise to the probability of finding electrons in a particular position inside the atom. Gives rise to atomic orbitals.

25. Quantum Mechanics Heisenberg Uncertainty PrincipleHeisenberg Uncertainty Principle –Impossible to know both the velocity and position of an electron at the same time Microscope Electron  Werner Heisenberg ~1926

26. Atomic orbital’s An atomic orbital is a region in space where there is a high probability of finding an electron. Erwin Schroedinger.— used maths equations to work out probability of finding an electron in any particular sublevel in an atom.

27. Quantum Mechanics Schrödinger Wave EquationSchrödinger Wave Equation (1926) quantized –finite # of solutions  quantized energy levels probability –defines probability of finding an electron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Erwin Schrodinger ~1926

28. Quantum Mechanics Orbital (“electron cloud”) –Region in space where there is 90% probability of finding an electron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Electron Probability vs. Distance Electron Probability (%) Distance from the Nucleus (pm) 100150200250500 0 10 20 30 40 Orbital 90% probability of finding the electron

29. Orbitals S-orbital-----spherical in shape P-orbital------ Dumbell in shape

30. Relative Sizes 1s and 2s 1s 2s Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 334

31. Quantum Numbers Principal Quantum Number n 1. Principal Quantum Number ( n ) –Energy level –Size of the orbital –n 2 = # of orbitals in the energy level Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 1s1s 2s2s 3s3s

32. Quantum Numbers pxpx pzpz pypy x y z x y z x y z

33. p-Orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 335 pxpx pypy pzpz

34. Quantum Numbers s p d f Angular Momentum Quantum # l 2. Angular Momentum Quantum # ( l ) –Energy sublevel –Shape of the orbital Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

35. Quantum Numbers Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

36. Quantum Numbers n=# of sublevels per level n 2 =# of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem n = 3 n = 2n = 1 Principallevel Sublevel Orbital ssp sp d pxpx pypy pzpz d xy d xz d yz dz2dz2 d x 2 - y 2 pxpx pypy pzpz

37. Maximum Capacities of Subshells and Principal Shells n 1 2 3 4...n l 0 0 1 0 1 2 0 1 2 3 Subshell designation designation s s p s p d s p d f Orbitals in subshell subshell 1 1 3 1 3 5 1 3 5 7 Subshell capacity capacity 2 2 6 2 6 10 2 6 10 14 Principal shell capacity capacity 2 8 18 32...2n 2 Hill, Petrucci, General Chemistry An Integrated Approach  1999, page 320

38. Quantum Numbers Magnetic Quantum Number m l 3. Magnetic Quantum Number ( m l ) –Orientation of orbital –Specifies the exact orbital within each sublevel Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

39. d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336

40. Shapes of s, p, and d-Orbitals

41. Atomic Orbitals

42. s, p, and d-orbitals A s orbitals: Hold 2 electrons (outer orbitals of Groups 1 and 2) B p orbitals: Each of 3 pairs of lobes holds 2 electrons = 6 electrons (outer orbitals of Groups 13 to 18) C d orbitals: Each of 5 sets of lobes holds 2 electrons = 10 electrons (found in elements with atomic no. of 21 and higher) Kelter, Carr, Scott,, Chemistry: A World of Choices  1999, page 82

43. Principal Energy Levels 1 and 2

44. Quantum Numbers Spin Quantum Number 4. Spin Quantum Number ( m s ) –Electron spin  +½ or -½ –An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

45. Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron Pauli Exclusion PrinciplePauli Exclusion Principle –No two electrons in an atom can have the same 4 quantum numbers. –Each electron has a unique “address”: Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

46. Electron Orbitals: Electron orbitals Equivalent Electron shells (a) 1s orbital (b) 2s and 2p orbitalsc) Neon Ne-10: 1s, 2s and 2p 1999, Addison, Wesley, Longman, Inc.

47. Feeling overwhelmed? Read Section 4-2! Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem "Teacher, may I be excused? My brain is full." Chemistry