1. Topic 2.2: Electrons Honors Chemistry 2014-15 Mrs. Peters 1

2. 2.2: Electron Configuration Essential Idea: The electron configuration of an atom can be deduced from its atomic number. Nature of Science: Developments in scientific research follow improvements in apparatus – the use of electricity and magnetism in Thomson’s cathode rays. (1.8) Theories being superseded – quantum mechanics is among the most current models of the atom (1.9) Use theories to explain natural phenomena – line spectra explained by the Bohr model of the atom (2.2) 2

3. 2.2: Electron Configuration Understandings: Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level. The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n 2. 3

4. 2.2: Electron Configuration Understandings: (Continued) A more detailed model of the atom describes the division of the main energy level into s, p, d, and f sub-levels of successively higher energies. Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron. Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin. 4

5. 2.2: Electron Configuration Applications and Skills: Description of the relationship between colour, wavelength, frequency, and energy across the electromagnetic spectrum. Distinction between a continuous spectrum and a line spectrum. 5

6. 2.2: Electron Configuration Applications and Skills: (Continued) Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels. Recognition of the shape of an s atomic orbital and the p x, p y, and p z atomic orbitals. Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z=36. 6

7. Review of Topic 2.1 Let’s Review! List how many protons, neutrons and electrons the following elements have: o Li o C o O o Mg o P o Ar o Ca 7

8. Review of Topic 2.1 Let’s Review! List how many protons, neutrons and electrons the following elements have: o Li: p = 3, n = 4, e = 3 o C: p = 6, n = 6, e = 6 o O: p = 8, n = 8, e = 8 o Mg: p = 12, n = 12, e = 12 o P: p = 15, n = 16, e = 15 o Ar: p = 18, n = 22, e = 18 o Ca: p = 20, n = 20, e = 20 8

9. Review of Topic 2.1 Atomic Structure Review: Protons and Neutrons are located in the nucleus. Electrons are found in the electron cloud outside the nucleus. This unit will focus on the electron cloud and where to find electrons. 9

10. Bohr Model Bohr Model for the Atom: Nucleus in the center with protons and neutrons Electrons in layers or levels around nucleus 10 Driver.layer.com

11. Bohr Model Bohr Model for the Atom: Electrons are arranged in energy levels (layers) Shows the number of electrons in each energy level. Electron orbits are circular paths 11 Chemistry.tutorcircle.com

12. Bohr Model Bohr Model for the Atom: Which element is this the electron arrangement for? How do you know? 12 Chemistry.tutorcircle.com

13. Bohr Model Bohr Model for the Atom: Useful for explaining and predicting chemical properties Based on the fundamental idea that electrons exist in definite, discrete energy levels Electrons can move from one energy level to another 13

14. Bohr Model Bohr Model for the Atom: Limitations of this model: Assumes all orbits are fixed Assumes all energy levels are circular Suggests incorrect scale for atom 14

15. U4. Quantum Mechanical Model U4. Quantum Mechanical Model Quantum Mechanical Model: Sophisticated mathematical theory that incorporates wave-like nature of electrons Based on 2 key ideas: o Schrodinger’s Equation o Heisenberg’s uncertainty principle 15

16. U4. Quantum Mechanical Model Heisenberg’s Uncertainty Principal It is impossible to determine accurately both the momentum and the position of a particle simultaneously. It is not possible to state precisely the location of an electron and its exact momentum, we can calculate the probability of finding an electron in a given region of space 16

17. U4. Quantum Mechanical Model Schrodinger’s Equation o Formulated in 1926 by Austrian physicist Erwin Schrodinger o Equation integrates the dual wave-like and particle nature of the electron o Describe atomic orbitals: a region in space where there is a high probability of finding an electron. 17

18. U4. Quantum Mechanical Model Sublevels There are 4 types: s, p, d, and f Each type has a characteristic shape, specific number of orbitals and associated energy. Each orbital holds a maximum of 2 electrons 18

19. U5. Sublevels & A4. S Sublevels Sublevels s: spherical shape, 1 orbital, holds 2 e- 19

20. U5. Sublevels & A4. P Sublevels Sublevels P: dumbbell shaped, 3 orbitals, holds 6 e- Draw the three sublevel shapes 20

21. U5. Sublevels U5. Sublevels Sublevels d: double dumbbell shaped, 5 orbitals, holds 10 e- f: funky shaped, 7 orbitals, holds 14 e- 21

22. U3. Quantum Mechanical Model The Quantum Mechanical Model Electrons are not found at certain distances from the nucleus but are located in a region in space that is described by a set of 4 quantum numbers. The exact location and path of the electron can’t be determined. It estimates the probability of finding an electron within a certain volume of space surrounding the nucleus. Electron positions can be represented by a fuzzy cloud surrounding the nucleus (electron cloud). 22

23. U3. Quantum Mechanical Model The Quantum Mechanical Model 4 Quantum Numbers: n: energy level (called the principal quantum number) l: sublevels m l : orbital m s : spin 23

24. U3. Quantum Mechanical Model The Quantum Mechanical Model Each energy level can hold a maximum of electrons based on 2n 2 Ex: if n is 3, there can be up to 2(3) 2 electrons =18 e- Argon has 18e- and is at the end of energy level 3 24

25. A & S 5: Electron Configuration Three principles (rules) must be followed when representing electron configurations: 1. Aufbau Principle: electrons fill the lowest energy orbital first 2. Pauli Exclusion Principle: any orbital can hold a maximum of 2 electrons and those electrons have opposite spin. 3. Hund’s rule of maximum multiplicity: when filling orbitals of equal energy, electrons fill all orbitals singly before occupying in pairs. 25

26. Order of Energy Levels for Electrons This order must be followed every time! Each level must be filled before moving to the next level 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p 26

27. A&S 5. Electron Configuration Diagramming Electron Arrangement There are two methods for diagramming electron arrangement Orbital Filling Diagram Electron Configuration 27

28. A&S 5. Electron Configuration Orbital Filling Diagrams 1.Draw a box or line for each orbital and sublevel. (s = 1 line, p = 3 lines, d= 5 lines) 2.Place arrows to denote electrons. Maximum of 2 electrons per box. The first arrow is pointing up, the second arrow is pointing down to represent opposite spins. 3.Within a sublevel, each space must get an electron before the second electron is added ( ie: each p sublevel gets one before doubling up) 28

29. A&S 5. Electron Configuration Orbital Filling Draw orbital filling diagrams for the following atoms. H Be O 29

30. A&S 5. Electron Configuration Orbital Filling Draw orbital filling diagrams for the following atoms. H__ 1s Be __ __ 1s 2s O __ __ __ __ __ 1s 2s 2p 30

31. A&S 5. Electron Configuration Orbital Filling Draw orbital filling diagrams for the following atoms. H__ 1s Be __ __ 1s 2s O __ __ __ __ __ 1s 2s 2p 31

32. A&S 5. Electron Configuration Orbital Filling Draw orbital filling diagrams for the following atoms. Al Ca 32

33. A&S 5. Electron Configuration Al __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p Ca __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p4s 33

34. A&S 5. Electron Configuration Electron Configuration Start with 1s and follow the order, filling each orbital with the maximum number of electrons until all the electrons in the atom have a place. Write electron configurations for the following elements: H1s 1 Be1s 2 2s 2 O1s 2 2s 2 2p 4 Al Ca 34

35. A&S 5. Electron Configuration Electron Configuration Start with 1s and follow the order, filling each orbital with the maximum number of electrons until all the electrons in the atom have a place. Write electron configurations for the following elements: H1s 1 Be1s 2 2s 2 O1s 2 2s 2 2p 4 Al1s 2 2s 2 2p 6 3s 2 3p 1 Ca1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 35

36. Review of Topic 2.1 Let’s Review Ions! What is an ion? What are the two types of ions? List how many protons, neutrons and electrons do the following ions have: o Li +1 o O -2 o Mg +2 o P -3 o Al +3 36

37. Review of Topic 2.1 List how many protons, neutrons and electrons do the following ions have: o Li +1 p = 3, n = 4, e = 2 o O -2 p = 8, n = 8, e = 10 o Mg +2 p = 12, n = 12, e = 10 o P -3 p = 15, n = 16, e = 18 o Al +3 p = 13, n = 14, e = 10 37

38. A & S 5. Ion Configuration Cations lose electrons, Anions gain electrons For orbital filling : add or subtract the number of electrons in the charge, draw in electrons For Electron Configuration : add or subtract the number of electrons in the charge, fill in sublevels and orbitals 38

39. A & S 5. Ion Configuration Practice: Compare Mg and Mg +2 Number of electrons: Electron configuration: Compare O and O -2 Number of electrons: Electron configuration: 39

40. A & S 5. Condensed Configuration This is a short cut! Full electron configurations become lengthy and cumbersome with increasing atomic number Condensed Configurations are more convenient 40

41. A & S 5. Condensed Configuration Core Electrons: electrons that are in the inner energy levels Valence Electrons: electrons that are in the outer energy level Condensed = [nearest noble gas] + valence electrons 41

42. A & S 5. Condensed Configuration Condensed = [nearest noble gas core] + valence electrons Ex: Oxygen: has 8 e-, 2 in core and 6 valence, nearest noble gas is He (which has 2 e-) [He] 2s 2 2p 4 42

43. A & S 5. Condensed Configuration Condensed = [nearest noble gas core] + valence electrons Ex: Cobalt: [Ar} 4s 2 3d 7 43

44. A & S 5. Condensed Configuration Condensed = [nearest nobel gas core] + valence electrons Practice: Write the condensed configuration Cl: Zn: Mn: 44

45. Periodic Table Notes Let’s Label the periodic table! 45 Commons.wikimedia.org

46. A & S 1: Wavelength, Frequency and Energy Light consists of electromagnetic waves that can travel through space and matter All electromagnetic waves travel in a vacuum at the speed of light (c) C = 3.0 x 10 8 m/s 46

47. A & S 1: Wavelength, Frequency and Energy Three components of electromagnetic waves Amplitude (y): Height from the origin to the crest Wavelength (λ) : Distance between the crests Frequency (ν): Number of wave cycles to pass a given point per unit time 47

48. A & S 1: Wavelength, Frequency and Energy Wavelength is related to the frequency of the radiation by the equation c = λν Draw and label the wave diagram. 48

49. A & S 1. Electromagnetic spectrum The electromagnetic spectrum is an arrangement of all of the types of electromagnetic radiation in increasing order of wavelength or decreasing frequency 49

50. A & S 1. Electromagnetic spectrum The higher the frequency, the shorter the wavelength and the higher the energy E = hf E = energy (Joules) h = Planck’s constant (6.63 x 10 -34 J s) f = frequency (s -1 ) 50

51. A & S 1. Electromagnetic spectrum Electromagnetic Spectrum Range: Radiowaves: long wavelength, low energy radiation Microwaves Infrared radiation (IR) 51

52. A & S 1. Electromagnetic spectrum Electromagnetic Spectrum Range: Visible Light: what we can see (ROYGBIV) Ultraviolet waves (UV) X-Rays Gamma Rays: high energy radiation, short wavelength 52

53. A & S 1. Electromagnetic Spectrum 53

54. A & S 2. Continuous and Line Spectrums Continuous Spectrum Emission showing a continuous range of wavelengths and frequencies; all the colors together without any space between them 54

55. A & S 2. Continuous and Line Spectrums Line Spectrum Emission of specific elements showing a series of discrete lines; individual lines of color with space between each line On your paper, draw the difference between a continuous and line spectrum 55

56. U1. Emission spectrum Emission Spectrum: a series of lines against a black background (type of line spectrum) 56 Wps.prenhall.com

57. U1. Emission spectrum Absorption Spectrum: a pattern of dark lines against a colored background 57 www.astronomyknowhow.com

58. A & S2. Continuous and line spectrum 58

59. A & S 2. continuous and line spectrum 59 We use an instrument called a spectroscope to detect the emission spectrum for a given source of light.

60. U2. Line Emission Spectrum for Hydrogen Why do we see different colors of light in line spectrums? When electrons of a gaseous atom get excited, they are raised to a higher energy level. The extra energy is released as light when they drop back down to lower energy levels. 60

61. U2. Line Emission Spectrum for Hydrogen Why do we see different colors of light in line spectrums? The energy is provided by thermal or electrical energy Each element has its own unique line spectrum 61 Images.tutorvista.com

62. U2. Line Emission Spectrum for Hydrogen Must know the Hydrogen line series. Draw this in your notes or on the back of the Electromagnetic Spectrum paper. Indicate the colors and the wavelengths, notice the space between colors. 62

63. A & S 3 Emission Spectrum of Hydrogen The line spectrum that we see from visible light is called the Balmer series. Similar sets of lines can be seen in ultraviolet (Lyman Series) and Infrared (Paschen Series) 63

64. A & S 3 Emission Spectrum of Hydrogen Electrons move in orbits around the nucleus of the atom. Each orbit has a fixed amount of potential energy. The farther from the nucleus the orbit is, the more potential energy it has. 64

65. A & S 3 Emission Spectrum of Hydrogen When electrons absorb energy they can move out to higher energy levels (the excited state). When they fall back to the lower energy level (the ground state) they emit a photon, a discrete amount of energy. Photon energy is seen as light. 65

66. A & S 3 Emission Spectrum of Hydrogen Depending on how far they fall, different colors of light are given off. Red = short fall Violet = long fall The Balmer series of lines (visible light) is formed when the electrons fall back to the second energy level (n=2) 66

67. A & S 3 Emission Spectrum of Hydrogen Balmer Series: Visible Transition of electrons from outer levels to n=2. Spectral lines converge at increased values of n due to closer spacing of energy levels Red: n=3 to n=2 Blue-green: n=4 to n=2 Blue: n=5 to n=2 Violet: n=6 to n=2 67

68. STOP HERE! 68