1. Chapter 9 Electrons in Atoms and the Periodic Table

2. Homework  Assigned Problems (odd numbers only)  “Questions” (page 310-11)  “Problems” 31 to 89 (page 311-14)  “Cumulative Problems” 91-113 (page 315-17)  Highlight Problems 115 (optional)

3. Light: Electromagnetic Radiation  Energy is the capacity to do work  The process of moving matter against an opposing force.  Forms of energy include heat, electrical, and light  One way energy is transmitted through space is by electromagnetic radiation  A form of energy that travels through space at the speed of light  Transmits from one place to another in the form of a wave  Given off by atoms when they have been excited by any form of energy  Electromagnetic radiation carries (radiant) energy through space and travels in waves at the speed of light  Waves are periodic: The pattern of peaks and troughs repeats itself at regular intervals

4. Light: Electromagnetic Radiation  The waves have three basic characteristics: wavelength, frequency, and speed  Wavelength (  is the distance (in nm) between neighboring peaks in a wave  The highest point on the wave is a peak  Shorter wavelengths are higher in energy  Longer wavelengths, are lower in energy

5. Light: Electromagnetic Radiation  Frequency (  ) is the number of waves that pass a fixed point in one unit of time  measured in Hertz (Hz),  1 Hz = 1 wave/sec = 1 sec -1  Velocity (v = how fast the wave is moving)  c = speed of light  3.00 x 10 8 m/s  Amplitude the height of the wave. It is the distance from the rest position to crest position or from rest position to trough position amplitude

6. Wavelength and Frequency  Because all EM radiation travels at the speed of light (c), a relationship exists between wavelength and frequency  This is an inverse relationship so that if the wavelength doubles, the frequency is halved. If the wavelength is halved, the frequency doubles (and vice-versa) C = λ ѵ C = 2 λ · ½ ѵ C = ½λ · 2 ѵ C = ½ λ · 2 ѵ

7. Waves C = speed of light frequency frequency wavelength wavelength C = λ ѵ

8. The Electromagnetic Spectrum  Light (radiant) energy is the energy of electromagnetic waves and it is classified into types according to the frequency of the wave  Sunlight, visible light, radio waves, microwaves (ovens), X-rays, and heat from a fire (infrared), are all forms of this radiant energy  These forms of radiant energy exhibit the same wavelike characteristics  The electromagnetic spectrum ranges from high- energy gamma and X-rays to very low-energy radio and TV waves

9. The Electromagnetic Spectrum  EM radiation is classified by wavelength:  Lower energy (longer wavelength, lower frequency)  Higher energy (shorter wavelength, higher frequency)  Radiowaves: AM/FM/TV signals, cell phones, low frequency and energy  Microwaves: Microwave ovens and radar  Infrared (IR): Heat from sunlight, infrared lamps for heating  Visible: The only EM radiation detected by the human eye ROYGBIVROYGBIVROYGBIVROYGBIV  Ultraviolet: Shorter in wavelength than visible violet light, sunlight  X-rays: Higher in energy than UV  Gamma rays: Highest in energy, harmful to cells

10. Wavelengths of EM Radiation  The electromagnetic spectrum ranges from high-energy gamma and X-rays to very low-energy radio and TV waves  The visible region of light is a narrow range of wavelengths between these two extremes

11. Light Emission by Different Elements  When white light passes through a prism it separates and produces a continuous rainbow of colors from (red, orange, yellow, green, blue, indigo, and, violet)  From red light to violet light the wavelength becomes shorter (700 nm to 400 nm)

12. Light Emission by Different Elements  When an element is heated its atoms absorb energy and re-emits that energy  Light is produced  If this light is passed through a prism, it does not produce a continuous rainbow, only certain colors

13. Emission Spectra  Only specific colors are produced in the visible region. This is called a “bright-line spectrum”  Each line produced is a specific color, and thus has a specific energy  Each element produces a unique set of lines (colors) which represents energy associated with a specific process in the atom  Lines are also produced in the infrared and ultraviolet regions White light produces a continuous spectra Each element produces a different discontinuous spectra

14. Emission Spectra  Scientists first detected the line spectrum of hydrogen (mid-1800’s) which produced only four lines

15. Emission Spectra  Scientist could not explain why atoms excited with energy produced discontinuous spectra  After the discovery of the nuclear structure of the atom (Rutherford, 1911), scientist thought of the atom as a microscopic solar system with electrons orbiting the nucleus  To explain the bright line spectrum of hydrogen, Bohr’s theory of the hydrogen atom began with this idea and assumed the electrons move in circular orbits around the nucleus Light emitted from hydrogen produces only specific wavelengths of light

16. Emission Spectra for Hydrogen: The Bohr Model  In 1913 Bohr developed a quantum model based on the emission spectrum for hydrogen  The proposal was based on the electron in hydrogen moving around the nucleus in a circular orbit

17. The Bohr Model: Atoms with Orbits  The Bohr atom has several orbits with a specific radius and specific energy  Each orbit or energy level is identified by “n” the principal quantum number  The values of n are positive, whole numbers 1, 2, 3, etc.  The principal energy level (n =1) has the lowest energy and the smallest radius  Electrons can be “excited” to a higher energy level with absorption of energy  The energy absorbed and released is equal to the energy difference between the two states nucleus

18. The Bohr Model: Atoms with Orbits  The different lines in an emission spectrum are associated with changes in an electron’s energy  Each electron resides in a specific E level called it’s principal quantum number (n, where n=1, n=2…)  Electrons closer to nucleus have lower energy (lower n values)  Electrons farther from the nucleus have higher energy (higher n values)

19. The Bohr Model: Excitation and Emission  Scientists associated the lines of an atomic spectrum with changes in an electrons energy (“Bohr Model”)  An electron excited to a higher energy state will return to a lower energy state  The energy that is given off (emitted) is a photon of light that corresponds to the energy difference between the higher and lower energy states  This precise amount of energy is called a quantum A photon (of light)

20. The Bohr Model: Excitation and Emission  The energy of a photon is related by the equation:  “The energy of a photon is directly proportional to its frequency”  “The energy of a photon is inversely proportional to its wavelength”  Energy transitions between orbits closer together produce photons of light with longer wavelengths (lower energy) E = h ѵ c = λ ѵ ѵ = c/λ E = hc/λ

21. The Bohr Model: Electron Energy Levels  Electrons possess energy; they are in constant motion in the large empty space of the atom  The arrangement of electrons in an atom corresponds to an electron’s energy  The electron resides outside the nucleus in one of seven fixed energy levels  Energy levels are quantized: Only certain energy values are allowed

22. The Bohr Model: Electron Energy Levels  Electrons can be “excited” to a higher E level with the absorption of E  The energy absorbed is equal to the difference between the two E states  When an electron loses E and falls to a lower E level, it emits EM radiation (photon)

23. The Bohr Model: Electron Energy Levels  If the EM radiation wavelength is in the visible spectrum a color is seen

24. The Bohr Model: Electron Energy Levels  The energy levels calculated by the Bohr model closely agreed with the values obtained from the hydrogen emission spectrum  The Bohr model did not work for other atoms  Energy levels were OK but model could not predict emission spectra for an element with more than one electron  Shrodinger in 1926 (DeBroglie, Heisenberg) developed the more precise quantum-mechanical model  The quantum (wave) mechanical model is the current theory of atomic structure

25. The Quantum-Mechanical Model: From Orbits to Orbitals  The quantum-mechanical model gives a new way to view electronic structure  This model combines the wavelike and particle-like behavior of the electron  For the hydrogen atom, the allowed energy states are the same as that predicted by the Bohr model  The Bohr model assumes the electron is in a circular orbit of some distance from the nucleus  In the quantum-mechanical model, the electron’s location cannot be described exactly  The electron’s location is described as region of space (probability) where the electron will be at any given instant

26. The Quantum-Mechanical Model: From Orbits to Orbitals  The electron is treated not as a particle but as a wave bound to the nucleus  The electron does not move around the nucleus in a circular path (orbit)  Instead, the electron is found in orbitals. It is not a circular path for the electron  An orbital indicates the probability of finding an electron near a particular point in space  An orbital is a map of electron density in 3-D space  Each orbital is characterized by a series of numbers called quantum numbers

27. The Quantum-Mechanical Model: Electron Energy Levels  Electrons with higher E will tend to be farther from the nucleus than those of lower E  The energy of an electron and its various distances from the nucleus can be grouped into levels or shells  Principal quantum number “n” is the major energy level in the atom: It has values of n =1, 2, 3, etc.  As “n” increases the size of the principal energy level (shell) increases Principal shell electron capacity = 2n 2

28. The Quantum-Mechanical Model: Electron Sublevels  All electrons in a principal shell (E level) do not have the same energy  The energy of electrons in the same shell have energies close in magnitude, but not identical  The range of energies for electrons in a shell is due to the existence of electron subshells (or energy sublevels)  An electron subshell is an energy level within an electron shell in which electrons all have the same energy

29. The Quantum-Mechanical Model: Electron Sublevels  The number of subshells (sublevels) within a principle shell (E level), n, varies  Each principal shell is divided into 1, 2, 3, or 4 subshells  Subshells are identified by a number and a letter: s, p, d, and f  Each principal shell contains the same number of subshells as its own principal shell number: No. of subshells in a principal shell = n Two electrons per subshell

30. The Quantum-Mechanical Model: Electron Sublevels  The order of the increasing energy for subshells (within an shell)  The subshells with the lowest to highest energy:  s subshell (holds up to 2 electrons)  p subshell (holds up to 6 electrons)  d subshell (holds up to 10 electrons)  f subshell (holds up to 14 electrons) s < p < d < fLowestenergyHighestenergy

31. Quantum-Mechanical Orbitals  The third term used to describe electron arrangement about the atomic nucleus (shells, subshells) is the orbital  Since the electron location cannot be known exactly, the location of the electron is described in term of probability, not exact paths  The orbital is a region of space where an electron assigned to that orbital is likely to be found  Region in space around the nucleus where there is a high (90%) probability of finding an electron of a specific energy

32. Quantum-Mechanical Orbitals  Each orbital can hold up to 2 electrons  Each subshell is composed of one or more orbitals  One orbital in an s-subshell  Three orbitals in a p-subshell  Five orbitals in a d-subshell  Seven orbitals in an f-subshell  Orbitals within the same subshell differ mainly in orientation

33. Quantum-Mechanical Orbitals  The orbitals in each of the four subshells (sublevels) have characteristic shapes  Orbitals in an s-subshell do not have the same shape as orbitals in a p-subshell, etc.  Orbitals of the same type, but in different principal shells/E levels (e.g. 1s, 2s, 3s) have the same general shape, but differ in size  The nucleus is located at the center of each orbital

34. Quantum-Mechanical Orbitals: s-Orbitals  There is one s-orbital in each s-subshell  Every principal shell contains only one s-orbital within an s-subshell  S-orbitals are spherical in shape  The larger the principal shell (energy level), the larger the sphere  An s-sublevel can hold a total of two electrons within the s-orbital

35. Quantum Mechanical Orbitals: s-Orbitals  The spherical s-orbital gets larger as n increases nucleus

36. Quantum Mechanical Orbitals: p-Orbitals  The p-orbitals come in sets of three within each p-subshell  All of equal energy  The three p-orbitals first occur in the n=2 (or higher) levels  P-orbitals are dumb-bell in shape  The three orbitals within a p-sublevel are oriented at right angles to one another and labeled as (p x, p y and p z )  p-subshell can hold a total of six electrons, two electrons in each of the p-orbitals (p x, p y and p z )

37. Quantum Mechanical Orbitals: p-Orbitals p-orbitals have a two-lobe, dumbbell shape. The nucleus is at the point where the two lobes meet nucleus

38. Quantum Mechanical Orbitals: d-Orbitals  d-orbitals come in sets of five within each d-subshell  All of equal energy  The five orbitals first occur in the n=3 shell  Odd shapes (don’t need to know them)  d-subshell can hold a total of 10 electrons, 2 electrons in each of five d- orbitals

39. d- Orbitals

40. f-Orbitals  f-orbitals come in sets of seven within each f-subshell  All of equal energy  The seven orbitals first occur in the n=4 shell  Shapes are very difficult, so you don’t need to know them either  f-subshell can hold a total of 14 electrons, 2 electrons in each of seven f-orbitals

41. Electron Configurations: How Electrons Occupy Orbitals  Two ways to show how the electrons are distributed in the principal shells within an atom  Orbital diagrams  Electron configurations  The most stable arrangement of electrons is one where the electrons are in the lowest energy subshells possible

42. Electron Configurations: How Electrons Occupy Orbitals Electron Configurations: How Electrons Occupy Orbitals  The most stable arrangement of electrons is called “ground-state electronic configuration”  The most stable, lowest energy arrangement of the electrons  The GS configuration for an element with many electrons is determined by a building-up process

43. Writing Orbital Diagrams and Electron Configurations  For the building-up process, begin by adding electrons to specific principal shells (E levels) beginning with the 1s subshell  Continue in the order of increasing subshell energies: 1s→2s →2p →3s →3p →4s →3d →4p →5s →4d →etc.

44. Writing Orbital Diagrams and Electron Configurations  The notation illustrates the electron arrangement in terms of which energy levels (shells) and sublevels (subshells) are occupied  The orbital diagram uses the building-up principal  Hund’s Rule: When electrons are placed in a set of orbitals of equal energy, the orbitals will be occupied by one electron each before pairing together

45. Electron Spin  Electrons behave as if they are spinning on an axis  A spinning electron behaves like a small bar magnet with north and south poles  Small arrows (pointed up or downward) are used to indicate the two orientations of spin  Two electrons in the same orbital must spin in opposite directions  Pauli Exclusion Principle: No more than two electrons can be placed in a single orbital and must be paired (have spins in opposite directions) orbital

46. Orbital Diagrams  Orbital Diagram Notation:  Draw a box to represent each orbital  Use an arrow up or down to represent an electron  Two electrons in the same orbital (box) must have spins in opposite directions: Only one up and one down arrow is allowed in a box (paired electrons) 1s2s2p

47. Orbital Diagrams  In General:  Begin filling from the lowest to the highest energy level  If there is more than one orbital possible, e.g., p x, p y, p z, place electrons alone before pairing up (Hund’s Rule)  Once each orbital is filled with one electron they will pair up but must have opposite spins (Pauli Exclusion Principal)

48.  s-orbitals  Only one per n  Can hold two electrons for a total of two electrons in an s-sublevel  p-orbitals  Three per n  Each can hold two electrons for a total of 6 electrons in a p-sublevel Orbital Diagrams

49.  d-orbitals  Five per n  Each can hold two electrons for a total of 10 electrons in a d-sublevel  f-orbitals  Seven per n  Each can hold two electrons for a total of 14 electrons in an f-sublevel

50.  hydrogen  Only one electron  Occupies the 1s orbital  helium  Two electrons  Both occupy the 1s orbital  lithium  Three electrons  Two occupy the 1s orbital, one occupies the 2s orbital Orbital Diagrams 1s 2s

51. Electron Configurations and the Periodic Table  The elements in the periodic table are arranged in order of increasing atomic number  The basic shape and structure of the table is consistent with (and can be explained by) the sequence used to build electron configurations  The table is divided into sections based on the type of subshell (s, p, d, or f) that receives the last electron in the building-up process

52. Electron Configurations and the Periodic Table  You can “build-up” atoms by reading across the periods from left to right  It is not necessary to memorize the filling order of the electron, just use the periodic table  Follow a path (left to right) across each period (row) of the table and note the various subshells as they are encountered  The atomic numbers are increasing across each period and this corresponds to increasing subshell energy  Since atomic numbers are increasing, each box in the table (across a period) is also an increase in one electron

53. Electron Configurations and the Periodic Table  The elements are arranged by increasing atomic number  The periodic table is divided into sections based on the type of subshell (s, p, d, or f) which receives the last electron in the build up process  Different blocks on the periodic table correspond to the s, p, d, or f sublevels

54. Electron Configurations and the Periodic Table   The specific location of an element in the periodic table can be used to obtain information about its electron configuration   An electron configuration is a statement of how many electrons an atom has in each of its subshells   To write a complete electron configuration:   The order in which the various subshells are filled can be obtained by following a path of increasing atomic number through the table (also taking account of the various subshells along the path)   The periodic table can be used to determine the shell in which the last electron added is located   It is this last electron added that causes an element’s electron configuration to differ from the preceding element

55. Electron Configurations and the Periodic Table  s-block elements (Groups 1A and 2A) gain their last electron in an s-sublevel  p-block elements (Groups 3A to 8A) gain their last electron in a p-sublevel  d-block elements (transition metals) gain their last electron in a d-sublevel. First appear after calcium (element 20)  d-sublevel is (n-1) less than the period number  f-block elements are in the two bottom rows of the periodic table  f-sublevel is (n-2) less than the period number

56. 1 2 3 4 5 6 7 Subshell Filling Order ns np (n-1)d (n-2) f Principal quantum number (n)= number of the period

57. Writing Electron Configurations from the Periodic Table  Locate the element, the number of electrons is equal to the atomic number  Start at hydrogen and move from box to box, in order of increasing atomic number  The lowest energy sublevel fills first, then the next lowest following a path across each period  The configuration of each element builds on the previous element  The p, d, or f sublevels must completely fill with electrons before moving to the next higher sublevel

58. Electron Configuration Example #1  Write the complete electron configuration for chlorine  Chlorine is atomic number 17 (on the periodic table) so the neutral atom has 17 electrons  Writing sublevel blocks in order up to chlorine gives: 1s 2 2s 2 2p 6 3s 2 3p x

59. 1 2 3 4 5 6 7 Electron Configuration Example #1 ns np (n-1) d (n-2) f

60. Electron Configuration Example #1 1s2s2p3s3p Orbital diagramHund’s Rule

61. Electron Configuration Example #2  Write the complete electron configuration for calcium  Calcium is atomic number 20 (on the periodic table) so the neutral atom has 20 electrons  Writing sublevel blocks in order up to calcium gives: 1s 2 2s 2 2p 6 3s 2 3p 6 4s x

62. 1 2 3 4 5 6 7 Electron Configuration Example #2 ns np (n-1) d (n-2) f

63. Electron Configuration Example #2 1s2s2p3s3p4s Orbital diagramHund’s Rule

64. Electron Configurations Examples   May also use the condensed (inner) electron configuration   This shorthand notation uses the noble gas that precedes a particular element and places it inside square brackets abbrev. electron configuration [ ] Noble gas core

65. Electron Configurations and the Periodic Table  The periodic table graphically represents the behavior of the elements described by periodic law  Elements are arranged by increasing atomic number  In the periodic table, elements with similar properties occur at regular intervals (in the same vertical column)  The arrangement of electrons and not the mass determines chemical properties of the elements

66. Valence Electrons  Valence electrons are those electrons in the outermost (highest) energy level “n” (where n = 1, 2, 3 …)  Those electrons not in the outermost (highest) energy level are called core electrons  Valence electrons are the most important (chemically)  Always found in the outermost s or p sublevels in the representative elements  For elements in columns 1A-8A, group number equals the number of valence electrons

67. Valence Electrons  All elements within a column (group) have the same number of valence electrons and similar outer electron configurations  Group IA elements have one valence electron: ns  Group IA elements have one valence electron: ns 1  Group IIA elements have two valence electrons: ns 2  Group IIIA elements have three valence electrons: ns 2 np 1

68. Periodic Trends of the Elements/Valence Electrons  Write the electron configuration for lithium  Write the electron configuration for sodium  Each group 1A element has a single electron in an s-sublevel. This is the (one) valence electron Li: 1s 2 2s 1 Na: 1s 2 2s 2 2p 6 3s 1

69. Periodic Trends of the Elements/Valence Electrons  The periodic table list elements by increasing atomic number and arranges them in groups with similar chemical properties  Similar chemical properties arise in every eighth element due to the similarity in electronic configurations (every eighth element for main group elements)  Across a period, elements become less metallic and more nonmetallic  Metals tend to lose electrons in chemical reactions

70. Periodic Trends of the Elements/Valence Electrons  Alkali metals lose their one and only one valence electron in chemical reactions forming an ion with a single positive charge and a stable noble gas electronic configuration  Group IIA metals lose their two valence electrons in chemical reactions forming an ion with a 2 + charge and a stable noble gas electronic configuration  Group VIIA nonmetals readily gain one electron in chemical reactions forming an ion with a single negative charge and obtain the stable electron configuration of the next higher noble gas

71. Atomic Size  Atoms are considered spherical in shape and their size (atomic radius) is very dependent on the electronic configuration of the atom  The electronic configuration gives trends in atomic size within groups and across periods in the periodic table (representative elements)  Within groups, the atomic radius increases with the period number (increase from top to bottom)  Across periods, the atomic radius decreases from left to right with increasing atomic number (decrease from left to right)

72. Atomic Size  Within groups:  The period number increases downward in a group  Principal E level (n) increases  Valence electron is further from the nucleus  Across periods:  The atomic radius decreases from LEFT to RIGHT with increasing atomic number  As atomic number increases, so does the number of electrons  The increase in positive charge pulls the outermost electrons closer to the nucleus

73. Size of Atoms and Their Ions  The formation of a positive ion requires the loss of one or more valence electrons  Loss of the outermost (valence) electron causes a reduction in atomic size  Positive ions are always smaller than their parent ions

74. Size of Atoms and Their Ions  The formation of a negative ion requires the addition of one or more electrons to the valence shell of an atom  There is no increase in + nuclear charge to offset the added electron’s - charge  Increase in size due to repulsion between electrons

75. Ionization Energy  The minimum energy required to remove one electron from an atom of an element (physical state is a gas)  The more tightly an electron is held, the higher the ionization energy  The trend in ionization energy parallels the metallic to nonmetallic trend in the chemical properties of the elements in a period

76. Ionization Energy  In the same group (top to bottom) ionization Energy decreases  Energy required to remove an electron decreases  Due to larger principal energy level (larger n value)  This puts outer electron farther from nucleus  As n increases, ionization energy decreases  Across same period (left to right) ionization Energy increases  Metals (left end) have lower ionization E  Tend to lose electrons to form + ions  Nonmetals (right end) have higher ionization E  Tend to gain electrons in chemical reactions

77.  End