1. Electrons in Atoms and the Periodic Table Chapter 9 Tro, 2 nd ed.

2. A Brief History of Atomic Theory Greeks were the first to suggest that matter was made up of small particles called atoms Early chemists performed experiments to find these particles Their experiments led to Dalton’s Atomic Theory Dalton’s Atomic Theory was revised further to the Thompson and Rutherford “nuclear” models of the atom Further work has led to the modern nuclear atomic theory, which is still being revised

3. A Brief History of Atomic Theory ~1900 Max Planck: energy emitted in bursts called quanta ~1905 Einstein used quanta to explain properties of light Called quantum of light a photon ~1911 Rutherford proposed “nuclear atom” Nucleus is a tiny center in the atom: Contains most of the mass Contains all of the positive charge (Thought electrons in orbit like mini-solar system - not!)

4. WHAT YOU NEED TO KNOW ABOUT ENERGY AND LIGHT (& e-s & photons) Light is part of the electromagnetic spectrum of radiation Continuous spectrum is like a “rainbow” Discrete spectrum is lines of specific colors of light Each line of color has its own energy associated with it Light in a vacuum moves at constant speed c = 2.998 X 10 8 m/sec (memorize!)

5. WHAT YOU NEED TO KNOW ABOUT ENERGY AND LIGHT (& e-s & photons) Electromagnetic Radiation - Examples light from the sunx-rays microwavesradio waves television wavesradiant heat All show wavelike behavior. Each travels at the same speed in a vacuum:2.998 x 10 8 m/s

6. visible light is part of the electromagnetic spectrum X-rays are part of the electromagnetic spectrum Infrared light is part of the electromagnetic spectrum The Electromagnetic Spectrum Know visible light as ROYGBIV.

7. Wavelength, (measured from peak to peak) Wavelength, (measured from trough to trough) Light has the properties of a wave. Characteristics of a Wave

8. Frequency,, is the number of wavelengths that pass a particular point per second. Characteristics of a Wave

9. Speed is how fast a wave moves through space. Characteristics of a Wave

10. WHAT YOU NEED TO KNOW ABOUT ENERGY AND LIGHT (& e-s & photons) Each specific “color” of visible light has a specific wavelength called and a specific frequency called and both relate to speed of light c = Wavelength,, is length unit: meters/wavelength Frequency,, is number of wavelengths/second Speed of light:c = meters* wavelengths = m/s wavelength second Energy (Joules) of one photon of light of specific wavelength or frequency: E = h = Joules/photon h = Planck’s constant (6.626 x 10 -34 J. s)

11. The Bohr Atom At high temperatures or voltages, elements in the gaseous state emit light of different colors. When the light is passed through a prism or diffraction grating a line spectrum results. Niels Bohr, a Danish physicist, in 1912-1913 carried out research on the hydrogen atom.

12. Line spectrum of hydrogen. Each line corresponds to the wavelength of the energy emitted when the electron of a hydrogen atom, which has absorbed energy falls back to a lower principal energy level. These colored lines indicate that light is being emitted only at certain wavelengths. Each element has its own unique set of spectral emission lines that distinguish it from other elements.

13. Spectra Tro: Figure 9.8: each element produces its own unique emission spectrum.

14. Energy is absorbed by e-, then emitted as a photon of light. Tro: Figure 9.11

15. An electron can have one of several possible energies depending on its orbit. E2E2 E3E3 E1E1 The Bohr Atom

16. Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus. An electron has a discrete energy when it occupies an orbit. The Bohr Atom

17. When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom. The color of the light emitted corresponds to one of the lines of the hydrogen spectrum. The Bohr Atom Different lines of the hydrogen spectrum correspond to different electron energy level shifts.

18. Light is not emitted continuously, but is emitted in discrete packets called photons The Bohr Atom

19. CALCULATIONS FOR  AND E OF A PHOTON EMITTED BY AN ELECTRON When e- in H atom moves from energy level 1 up to energy level 4 and then drops back down to energy level 2, we see a photon of light emitted that has a wavelength,, of 4.86 x 10 -7 m. Calculate the frequency of the light and the energy of the photon emitted. c = Rearrange to get frequency, = c/ = (2.998x10 8 m/s)/(4.86x10 -7 m) = 6.17 x 10 14 /s E = h = 6.626x10 -34 J. s * 6.17 x 10 14 /s = 4.09 x 10 -19 J/photon of light

20. MODERN ATOMIC THEORY Bohr’s calculations succeeded very well for the hydrogen atom. Bohr’s methods did not succeed for heavier atoms. So…more theoretical work on atomic structure was needed.

21. MODERN ATOMIC THEORY Thompson had shown that light, which is photons of energy, had the properties of matter as well. In 1924, Louis De Broglie suggested that all matter must also have wave properties. De Broglie showed that the wavelength of ordinary sized objects, such as a baseball, are too small to be observed. For objects the size of an electron the wavelength can be detected.

22. MODERN ATOMIC THEORY In 1926, Schröedinger created a mathematical model that showed electrons as waves. Schröedinger’s work led to a new branch of physics called wave or quantum mechanics. Using Schröedinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined. The actual location of an electron within an atom cannot be determined (Heisenberg Uncertainty Principle).

23. MODERN ATOMIC THEORY Based on wave mechanics it is clear that electrons are not revolving around the nucleus in orbits. Instead of being located in orbits, the electrons are located in orbitals. An orbital is a region around the nucleus where there is a high probability of finding an electron.

24. According to Bohr the energies of electrons in an atom are quantized. The wave-mechanical model of the atom also predicts discrete principal energy levels within the atom.

25. Energy Levels of Electrons Principal energy levels, n = 1,2,3...(also called principal quantum number) Each energy level contain(s) sublevel(s) called s, p, d, and f (the angular momentum quantum number, l ) Within sublevels are orbitals (designated by orientation in 3-D space, called the magnetic quantum number, m l ) Each orbital can hold 2 e - s max (each electron is assigned a spin quantum number, m s )

26. The first four principal energy levels of the hydrogen atom As n increases, the energy of the electron increases Each level is assigned a principal quantum number, n Energy levels in atoms

27. Each principal energy level is subdivided into sublevels. The number of sublevels equals the assigned energy level. For n=1, there is one sublevel, s. For n=2, there are two sublevels, s & p, etc. The sublevels have the quantum designation, l. The maximum value is n -1. Energy levels in atoms

28. Within sublevels the electrons are found in orbitals. An s orbital is round and soft, like a nerf ball. The shape represents the highest probability where the electron might be found.

29. An electron can spin in one of two possible directions represented by ↑ or ↓. The two electrons that occupy an atomic orbital must have opposite spins. This is known as the Pauli Exclusion Principal. An atomic orbital can hold a maximum of two electrons.

30. Each p orbital has two lobes and can hold a maximum of two electrons. Since there are three orbitals, a p sublevel can hold a maximum of 6 electrons. A p sublevel is made up of three p-type orbitals.

31. The three p orbitals all center at the atom’s nucleus… pxpx pypy pzpz …and occupy one of the three axes of 3-D space.

32. The five d orbitals lie in different planes and point in different directions. Each d orbital can hold a maximum of two electrons. A d sublevel can hold a maximum of 10 electrons. A d sublevel is made up of five orbitals.

33. Shells & Subshells Tro Figure 9.19: The number of subshells within a shell is equal to the value of n, the principal quantum number. Etc, etc, etc: What subshells exist for n = 5?

34. Energy Levels of Electrons Pauli Exclusion Principle: Each orbital can hold a max of 2 e-s, so possibilities are 0, 1 or 2 e-s. s has only 1 orbital  2 e-s MAX p has 3 orbitals  6 e-s MAX d has 5 orbitals  10 e-s MAX f has 7 orbitals  14 e-s MAX (g has ? Orbitals  ?? e-s MAX)

35. Energy Levels of Electrons Max is 2 e-s because of a property of e-s called spin. (Pauli Exclusion Principle) Each e- is spinning on its axis like the earth. Any spinning charge creates a magnetic field, with N and S poles. The direction of spin determines which is North. The e-s will pair up in an orbital so their N poles are opposite each other.

36. Atomic Structure of the First 18 Elements: Use these guidelines The ground state of the electron is the lowest energy orbital it can occupy. Higher energy orbitals are excited states. The distribution of electrons into the various energy shells and subshells in an atom in its ground state is called its electron configuration. Each energy shell and subshell has a maximum number of electrons it can hold: s = 2, p = 6, d = 10, f = 14 Place electrons in the energy shells and subshells in order of energy, from low energy up: the Aufbau Principal.

37. Atomic Structure of the First 18 Elements: Use these guidelines No more than two electrons can occupy one orbital Electrons occupy the lowest energy orbitals available, the ground state. They enter a higher energy orbital only after the lower orbitals are filled. (Aufbau again.) For the atoms beyond hydrogen, orbital energies vary as “s<p<d<f” for a given value of n. Each orbital in a sublevel is occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital. (Hund’s Rule)

38. Energy 1s 7s 2s 2p 3s 3p 3d 6s 6p 6d6d 4s 4p 4d 4f 5s 5p 5d 5f After 3p is filled, the next lowest energy is 4s, not 3d. See all the overlap beyond 3p.

39. Order of Subshell Filling in Ground State Electron Configurations 1s1s 2s2s2p2p 3s3s3p3p3d3d 4s4s4p4p4d4d4f4f 5s5s5p5p5d5d5f5f 6s6s6p6p6d6d 7s7s Start by drawing a diagram putting each energy shell on a row and listing the subshells, (s, p, d, f), for that shell in order of energy, (left-to-right) Next, draw arrows through the diagonals, looping back to the next diagonal each time

40. Writing Electron Configurations Arrangement of electrons within their respective sublevels. 2p62p6 Principal energy level Type of orbital Number of electrons in sublevel orbitals

41. Order of Subshell Filling in Ground State Electron Configurations Try using the drawing with the arrows to do aluminum, phosphorus and chlorine. Al 1s 2 2s 2 2p 6 3s 2 3p 1 P 1s 2 2s 2 2p 6 3s 2 3p 3 Cl 1s 2 2s 2 2p 6 3s 2 3p 5

42. Orbital Box Notation In the following diagrams boxes represent orbitals. Electrons are indicated by arrows: ↑ or ↓ for spin. Electron configurations can be spdf or orbital box. See both as follows.

43. ↑ 1s21s2 ↓ H↑1s11s1 Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. He Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins. Filling the 1s sublevel

44. Li 1s22s21s22s2 The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital. ↑ ↓ ↑ ↓ 1s1s2s2s ↑ 1s22s11s22s1 Be ↑↓ The 2s orbital fills upon the addition of beryllium’s third and fourth electrons. 1s1s2s2s Filling the 2s sublevel

45. B1s22s22p11s22s22p1 1s1s2s2s2p2p ↑↓ ↑↓ ↑ Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. C 1s1s2s2s2p2p ↑↓ ↑↓ The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. 1s22s22p21s22s22p2 ↑ ↑↑ N 1s1s2s2s2p2p ↑ ↓ ↓↑ The third p electron of nitrogen enters a different p orbital than its first two p electrons to give nitrogen the lowest possible energy. 1s22s22p31s22s22p3 ↑↑

46. ↑ 1s22s22p41s22s22p4 ↑↑ 1s1s2s2s2p2p ↑↓↑ ↓ O There are four electrons in the 2p sublevel of oxygen. One 2p orbitals is has a second electron, with a spin opposite to the electron already in the orbital. ↑↓ ↑↓↑ 2p2p F 1s1s2s2s ↑↓↑ ↓ There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. 1s22s22p51s22s22p5 ↓

47. ↑↓ ↑↓↑↓ 2p2p Ne 1s1s2s2s ↑↓↑ ↓ There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. 1s22s22p61s22s22p6

48. Na1s22s22p63s11s22s22p63s1 1s1s2s2s2p2p 3s3s ↑↓↑ ↓ ↑↓ ↑↓↑↓↑ The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium. Mg 1s1s2s2s2p2p 3s3s ↑↓↑ ↓ ↑↓ ↑↓↑↓↑ The 3s orbital fills upon the addition of magnesium’s twelfth electron. 1s22s22p63s21s22s22p63s2 ↓

50. After element 18, the 4s sublevel is filled before the 3d is filled.

51. Writing Electron Configurations Try writing the electron configurations for calcium, vanadium, iron, and arsenic. Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 V 1s 2 2s 2 sp 6 3s 2 3p 6 4s 2 3d 3 Fe 1s 2 2s 2 sp 6 3s 2 3p 6 4s 2 3d 6 As 1s 2 2s 2 sp 6 3s 2 3p 6 4s 2 3d 10 4p 3

52. Electron Structures and the Periodic Table In 1869, Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of the elements based on increasing atomic masses. Mendeleev’s arrangement is the precursor to the modern periodic table.

53. Horizontal rows are called periods. Period numbers correspond to the highest occupied energy level.

54. Elements in the A groups are designated representative elements. Elements in the B groups are designated transition elements. Elements with similar properties are organized in groups or families. Groups are numbered with Roman numerals.

55. SHORTHAND FOR ELECTRON CONFIGURATIONS: The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets. B1s 2 2s 2 2p 1 [He]2s 2 2p 1 Cl1s 2 2s 2 2p 6 3s 2 3p 5 [Ne]3s 2 3p 5 Na1s 2 2s 2 2p 6 3s 1 [Ne]3s 1

56. VALENCE ELECTRONS Valence electrons determine how atoms of the representative elements will combine with each other. (By forming ions or by sharing e-s.) Valence electrons are in highest principal energy level. FIND valence e-s: - by looking at where atom is on periodic table - or by looking at its written e- config and underlining the e-s in the highest # principal energy level Notice that Group 1A elements all have ns 1 in last place, Group 2A is ns 2, etc. What is Group 7A “generic” valence arrangement? ns 2 np 5

57. VALENCE ELECTRONS EXAMPLE: IODINE is element #53. Identify the number of valence e’s and their location. It is in group 7A, so it has 7 valence e-s. Or underline in the configuration: [Kr]5s 2 4d 10 5p 5.

58. For “A” family elements the valence electron configuration is the same in each column. The chemical behavior and properties of elements in a family are associated with the electron configuration of its elements.

59. With the exception of helium which has one filled s orbital, the nobles gases have filled all p orbitals.

60. VALENCE ELECTRONS TRY: Write electron configuration and use periodic table for arsenic, magnesium, argon, silicon, gallium. Determine which are valence electrons and how many there are. As[ 18 Ar] 4s 2 3d 10 4p 3 5 Mg[ 10 Ne] 3s 2 2 Ar[ 10 Ne] 3s 2 3p 6 8 Si[ 10 Ne] 3s 2 3p 2 4

61. 10.16 d orbital filling d orbital numbers are 1 less than the period number Arrangement of electrons according to sublevel being filled.

62. 10.16 f orbital filling f orbital numbers are 2 less than the period number Arrangement of electrons according to sublevel being filled.

63. 10.17 Period number corresponds with the highest energy level occupied by electrons in that period. Memorize the Cr and Cu difference!

64. Electron Configuration Using the Periodic Table See page 284 in your textbook for complete instructions. Locate the element on the Periodic Table and locate its row and column. The outermost electrons are determined by this. Then fill in the “inner” electrons by looking at the noble gas just before the element and filling in up to the atom. Make sure the electrons add up to the element’s atomic number. Example: Bromine. Row 4, column 7A or 17. Outermost electrons will end with 4p 5. Argon is just before Br. From Argon, go to row 4, s block, cross the 3d’s, and you’re at the 4p block. Bromine is [Ar]4s 2 3d 10 4p 5

65. GROUP PRACTICE: Try to write Ti, Fe, Pd, and U abbreviated configuration using the rules & using the Periodic Table Ti: [Ar] 4s 2 3d 2 Fe: [Ar] 4s 2 3d 6 Pd: [Kr] 5s 2 4d 8 U: [Rn] 7s 2 5f 4 Do the complete e- config and abbrev e- config for germanium, and underline its valence e-s.

66. The Noble Gas Electron Configuration Noble gases have 8 valence electrons - except for He, which has only 2 electrons Noble gases are especially nonreactive - He and Ne are practically inert Noble gases are nonreactive because their electron configurations are especially stable

67. Alkali Metals Alkali metals have one more electron than the previous noble gas Alkali metals tend to lose their extra electron, resulting in the same electron configuration as a noble gas, forming a cation with a 1+ charge

68. Halogens Electron configurations for halogens have one fewer electron than the next noble gas In reactions with metals, the halogens tend to gain an electron and attain the electron configuration of the next noble gas, forming an anion with charge 1- In reactions with nonmetals, they tend to share electrons with the other nonmetal, so that each attains the electron configuration of a noble gas

69. Stable Electron Configuration and Ion Charge Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas Isoelectronic series: N 3- O 2- F - Ne Na + Mg 2+ Al 3+ Make a table as above for the isoelectronic series around argon.

70. Periodic Trends in Atomic Properties Characteristic properties and trends of the elements are the basis of the periodic table’s design. These trends allow us to use the periodic table to accurately predict properties and reactions of a wide variety of substances.

71. Atomic radii increase down a group. For each step down a group electrons enter the next higher energy level. 11.2 Atomic Radius

72. Radii of atoms tend to decrease from left to right across a period. For representative elements within the same period, the energy level remains constant as electrons are added. This increase in positive nuclear charge pulls all electrons closer to the nucleus. 11.2 Each time an electron is added, a proton is also added to the nucleus.

73. Example 9.6 – Choose the Larger Atom in Each Pair C or O Li or K C or Al Se or I

74. Example 9.6 – Choose the Larger Atom in Each Pair C or O Li or K C or Al Se or I?

75. IONIZATION ENERGY The first ionization energy of an atom is the energy required to remove the first electron from an atom. Na + ionization energy → Na + + e - The second ionization energy is the amount of energy required to remove the second electron from an atom. Na + + second ionization energy → Na 2+ + e -

76. Periodic relationship of the first ionization energy for representative elements in the first four periods. 11.3 Ionization energies gradually increase from left to right across a period. IA IIA IIIA IVA VA VIA VIIA Noble Gases 1 2 3 4

77. Periodic relationship of the first ionization energy for representative elements in the first four periods. 11.3 Ionization energies of Group A elements decrease from top to bottom in a group. IA IIA IIIA IVA VA VIA VIIA Noble Gases Distance of Outer Shell Electrons From Nucleus nonmetalsmetals nonmetals have higher ionization potentials than metals

78. Example 9.7 – Choose the Atom with the Highest Ionization Energy in Each Pair Mg or P As or Sb N or Si O or Cl

79. Example 9.7 – Choose the Atom with the Highest Ionization Energy in Each Pair Mg or P As or Sb N or Si O or Cl?