2. Chapter 7 Atomic Structure and Periodicity

3. 7.1 Electromagnetic Radiation electromagnetic radiation: electromagnetic radiation: electromagnetic radiation electromagnetic radiation form of energy that acts as a wave as it travels form of energy that acts as a wave as it travels includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves travel at a speed of 2.9979 x 10 8 m/s in a vacuum travel at a speed of 2.9979 x 10 8 m/s in a vacuum All forms are combined to form electromagnetic spectrum All forms are combined to form electromagnetic spectrum

4. Electromagnetic Spectrum

5. The Wave-like Electron Louis deBroglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.

6. Electromagnic radiation propagates through space as a wave moving at the speed of light.

7. Wave nature of electromagnetic Radiation wavelength: wavelength: λ = Greek letter lambda λ = Greek letter lambda distance between points on adjacent waves (consicutive peaks or troughs) distance between points on adjacent waves (consicutive peaks or troughs) in nm (10 9 nm = 1m) in nm (10 9 nm = 1m) frequency: frequency:  = Greek letter nu  = Greek letter nu number of wave cycles that passes a point in a second. 10 8 cycles/s= 10 8 s -1 number of wave cycles that passes a point in a second. 10 8 cycles/s= 10 8 s -1 =10 8 Hertz = 10 8 Hz =10 8 Hertz = 10 8 Hz in 1/second (Hertz = Hz) in 1/second (Hertz = Hz)

8. C = speed of light, a constant (3.00 x 10 8 m/s) = frequency, in units of hertz (hz, sec -1 ) = wavelength, in meters

9. Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

10. Calculate the energy of red light vs. blue light. red light: 700 nmblue light: 400 nm red: blue: E = 2.85 x 10 -19 J E = 4.96 x 10 -19 J sunburn????? uv

11. 7.2 Nature of Matter Before 1900, scientists thought that matter and energy were totally different Before 1900, scientists thought that matter and energy were totally different matterenergy particlesmasspositionwavemasslessdelocalized

12. In 1900 Matter and energy were seen as different from each other in fundamental ways. Matter and energy were seen as different from each other in fundamental ways. Matter was particles. Matter was particles. Energy could come in waves, with any frequency. Energy could come in waves, with any frequency. Max Planck found that as the cooling of hot objects couldn ’ t be explained by viewing energy as a wave. Max Planck found that as the cooling of hot objects couldn ’ t be explained by viewing energy as a wave.

13. e-e- Photoelectric effect: Light of the right frequency (energy) can strike a metal and cause an electron to be ejected (n = infinity). Metal surface

14. Nature of Matter Max Planck: a German physicist Max Planck: a German physicist suggested that an object emits energy in the form of small packets of energy called suggested that an object emits energy in the form of small packets of energy called quanta Quantum- the minimum amount of energy that can be gained or lost by an atom Quantum- the minimum amount of energy that can be gained or lost by an atom Planck ’ s constant (h): 6.626 x 10 -34 J*s

15. Nature of Matter Einstein proposed that radiation itself is really a stream of particles called photons Einstein proposed that radiation itself is really a stream of particles called photons Energy of each photon is : Energy of each photon is : also showed that energy also showed that energy has mass

16. Nature of Matter shows that anything with both mass and velocity has a corresponding wavelength

17. Nature of Matter In 1924, Louis de Broglie In 1924, Louis de Broglie (French scientist) suggested that matter has both particle-like and wave-like characteristics suggested that matter has both particle-like and wave-like characteristics

18. Main Ideas: matter and energy are not distinct matter and energy are not distinct energy is a form of matter energy is a form of matter larger objects are mostly particle-like larger objects are mostly particle-like smaller objects are mostly wave-like smaller objects are mostly wave-like

19. …produces all of the colors in a continuous spectrum Spectroscopic analysis of the visible spectrum… 7.3 The Atomic Spectrum of Hydrogen White light

20. Continuous Spectra White light passed through a prism produces a spectrum – colors in continuous form.

21. The Continuous Spectrum The different colors of light correspond to different wavelengths and frequencies ~ 650 nm ~ 575 nm ~ 500 nm ~ 480 nm ~ 450 nm

22. Continuous Emission Spectrum line-emission spectrum- series of wavelengths of light created when visible portion of light from excited atoms is shined through a prism line-emission spectrum- series of wavelengths of light created when visible portion of light from excited atoms is shined through a prism scientists using classical theory expected atoms to be excited by whatever energy they absorbed scientists using classical theory expected atoms to be excited by whatever energy they absorbed continuous spectrum- continuous spectrum- continuous spectrum- continuous spectrum- emission of continuous range of frequencies of EM radiation emission of continuous range of frequencies of EM radiation contains all wavelengths of visible light contains all wavelengths of visible light

23. …produces a “bright line” spectrum Spectroscopic analysis of the hydrogen spectrum… H receives a high energy spark H-H bonds Are broken and H atoms are excited

24. Line Spectra Light passed through a prism from an element produces a discontinuous spectrum of specific colors

25. Hydrogen only four lines are observed

26. Line Spectra The pattern of lines emitted by excited atoms of an element is unique = atomic emission spectrum

27. These are called the atomic emission spectrum Unique to each element, like fingerprints! Very useful for identifying elements

28. H Line-Emission Spectrum light is emitted by excited H atoms when bond is broken in the light is emitted by excited H atoms when bond is broken in the diatomic molecule diatomic molecule ground state- lowest energy state of an atom ground state- lowest energy state of an atom excited state- when an atom has higher potential energy than it has at ground state excited state- when an atom has higher potential energy than it has at ground state

29. H Line-Emission Spectrum when an excited electron falls back to ground state, it emits photon of radiation when an excited electron falls back to ground state, it emits photon of radiation the photon is equal to the difference in energy of the original and final states of electron the photon is equal to the difference in energy of the original and final states of electron since only certain frequencies are emitted, only certain energies are allowed for electrons in H atom since only certain frequencies are emitted, only certain energies are allowed for electrons in H atom

30. This produces bands of light with definite wavelengths. Electron transitions involve jumps of definite amounts of energy.

31. 7.4 The Bohr Model Niels Bohr (Danish physicist) in 1913 Niels Bohr (Danish physicist) in 1913 Developed a quantum model Developed a quantum model for H atom that explained the emission line spectrum Electron moves around the nucleus only in certain allowed circular orbits, in which it has a certain amount of energy Electron moves around the nucleus only in certain allowed circular orbits, in which it has a certain amount of energy

32. The Bohr model Energy level of an electron analogous to the steps of a ladder Energy level of an electron analogous to the steps of a ladder The electron cannot exist between energy levels, just like you can’t stand between steps on a ladder The electron cannot exist between energy levels, just like you can’t stand between steps on a ladder A quantum of energy is the amount of energy required to move an electron from one energy level to another A quantum of energy is the amount of energy required to move an electron from one energy level to another

33. Niels Bohr Developed the quantum model of the hydrogen atom. Developed the quantum model of the hydrogen atom. He said the atom was like a solar system. He said the atom was like a solar system. The electrons were attracted to the nucleus because of opposite charges. The electrons were attracted to the nucleus because of opposite charges. Didn’t fall in to the nucleus because it was moving around. Didn’t fall in to the nucleus because it was moving around.

34. The Bohr Atom He didn’t know why but only certain energies were allowed. He didn’t know why but only certain energies were allowed. He called these allowed energies energy levels. He called these allowed energies energy levels. Putting Energy into the atom moved the electron away from the nucleus. Putting Energy into the atom moved the electron away from the nucleus. From ground state to excited state. From ground state to excited state. When it returns to ground state it gives off light of a certain energy. When it returns to ground state it gives off light of a certain energy.

35. The Model: Summary Space around nucleus is divided into spherical (circualr) paths (orbits) each has a number called “ Principal Quantum number ” Space around nucleus is divided into spherical (circualr) paths (orbits) each has a number called “ Principal Quantum number ” The electron can exist only in one of these orbitals but not in between The electron can exist only in one of these orbitals but not in between Orbits possess fixed size and energy, therefore electron has a definite energy characteristic of its orbit Orbits possess fixed size and energy, therefore electron has a definite energy characteristic of its orbit

36. Orbits allowed for electron are those in which electron has an angular momentum= Orbits allowed for electron are those in which electron has an angular momentum= An electron can pass only from one bit to another. Absorption or emission will occur An electron can pass only from one bit to another. Absorption or emission will occur Energy of the outermost orbit is zero Energy of the outermost orbit is zero

37. The Bohr Atom n = 3 n = 4 n = 2 n = 1

38. Bohr Model To create an accurate model, he had to use quantum theory instead of classical To create an accurate model, he had to use quantum theory instead of classical created an equation used to calculate the energy levels available to electrons in a certain atom: created an equation used to calculate the energy levels available to electrons in a certain atom: where n= integer and Z=atomic number negative sign makes the energy more negative the closer it is to the nucleus

39. Bohr Model Can gain energy by moving to a higher energy level Can gain energy by moving to a higher energy level Can lose energy by moving to lower energy level Can lose energy by moving to lower energy level

40. Bohr Model a photon is released that has an energy equal to the difference between the initial and final energy orbits

41. Bohr Model equation can be used twice to find the ∆E when an electron moves energy levels equation can be used twice to find the ∆E when an electron moves energy levels

42. Bohr Model can wavelength of photon released by using can wavelength of photon released by using E=0 is set at an distance of ∞ away from the nucleus and becomes more negative as the electron comes closer to the nucleus E=0 is set at an distance of ∞ away from the nucleus and becomes more negative as the electron comes closer to the nucleus

43. Example 1 Calculate the energy required to move the hydrogen electron from n=1 to n=2. Find the wavelength of radiation that had to be absorbed by the electron.

44. Example 2 Calculate the energy required to remove the electron from the hydrogen atom in its ground state. Energy was absorbed by the electron so the value of ∆E value is positive.

45. The Bohr Model Doesn’t work. Doesn’t work. Only works for hydrogen atoms. Only works for hydrogen atoms. Electrons don’t move in circles. Electrons don’t move in circles. The quantization of energy is right, but not because they are circling like planets. The quantization of energy is right, but not because they are circling like planets.

46. Bohr Model problems: problems: did not work for other atoms did not work for other atoms did not explain chemical behavior of atoms did not explain chemical behavior of atoms

47. Heisenberg ’ s Uncertainty Principle According to de Broglie: Electron behaves like a wave According to de Broglie: Electron behaves like a wave It is possible to specify the position of a wave at a particular instant? It is possible to specify the position of a wave at a particular instant? Energy, wavelength and amplitude can be determined Energy, wavelength and amplitude can be determined But exact position is impossible to be determined But exact position is impossible to be determined The electron cannot be imagined as : The electron cannot be imagined as : moving particle moving particle In a path of the same radius (well defined orbits) In a path of the same radius (well defined orbits) Thus, location, direction and speed of motion of a particle cannot be determined Thus, location, direction and speed of motion of a particle cannot be determined Then Bohr Model had to be “ Abandoned Then Bohr Model had to be “ Abandoned

48. Heissenberg Uncertainty Principle “ It is impossible to determine both the position and momentum of a subatomic particle (such as the electron) with arbitrarily high accuracy ” The effect of this principle is to convert the laws of physics into statements about relative, instead of absolute, certainties. The effect of this principle is to convert the laws of physics into statements about relative, instead of absolute, certainties.

49. Heisenberg Uncertainty Principle we cannot know the exact position and momentum (motion) of the electron we cannot know the exact position and momentum (motion) of the electron as more is known about position, less is known about momentum as more is known about position, less is known about momentum uncertainties are inversely proportional uncertainties are inversely proportional where ∆x: uncertainty in position ∆m  : uncertainty in mometum  minimum uncertainty is h/4 

50. 7.5 The Quantum Mechanical Model Exactt position of electron can not be defined Exactt position of electron can not be defined Exact bath of electron about nucleus can not be defined Exact bath of electron about nucleus can not be defined Werner Heisenberg, Louis de Broglie and Erwin Schrodinger made the approach called “ Quantum Mechanics ” Werner Heisenberg, Louis de Broglie and Erwin Schrodinger made the approach called “ Quantum Mechanics ” They assumed that the electron is a standing wave They assumed that the electron is a standing wave

51. The Quantum Mechanical Model Waves are associated with electrons Waves are associated with electrons Information about energies of electrons and their positions are obtained from studying the associated waves Information about energies of electrons and their positions are obtained from studying the associated waves Description of electron is based upon Description of electron is based upon “ Probability of finding a particle within a given region of space ” “ but not on the exact position ” “ Probability of finding a particle within a given region of space ” “ but not on the exact position ”

52. Schrödinger Equation Wave equation describing electron as being a wave Wave equation describing electron as being a wave The amplitudes (height), , of electron wave at various points of space are calculated The amplitudes (height), , of electron wave at various points of space are calculated  commonly called “ wave function ”  commonly called “ wave function ”  provides information about the allowable energies for an electron in H atom.  provides information about the allowable energies for an electron in H atom.  corresponds to a certain energy and describes a region around nucleus “ Orbital ” where the electron having that energy may be found  corresponds to a certain energy and describes a region around nucleus “ Orbital ” where the electron having that energy may be found

53. Orbital: Region around the nucleus where the electron can be expected to be found Orbital: Region around the nucleus where the electron can be expected to be found The Function  2 The Function  2  2 describes the probability of the position of the electron at a particular point  2 describes the probability of the position of the electron at a particular point  2  Probablity of finding a particle in a given region of space  2  Probablity of finding a particle in a given region of space  2  Electric charge density at a given region of space  2  Electric charge density at a given region of space

54. Thus, The charge can be assumed to be spread out as a charge cloud by rapid motion of electron The charge can be assumed to be spread out as a charge cloud by rapid motion of electron The cloud is denser in some regions than others The cloud is denser in some regions than others The probability of finding electron in a given region in space is proportional to the density of the cloud The probability of finding electron in a given region in space is proportional to the density of the cloud

55. Meaning of Wave Function the wave function itself does not have concrete meaning the wave function itself does not have concrete meaning the square of the wave function represents the probability of finding an electron at a certain point the square of the wave function represents the probability of finding an electron at a certain point easily represented as probability distribution where the deepness of color indicates the probability easily represented as probability distribution where the deepness of color indicates the probability

56. Meaning of Wave Function (a) electron density map (a) electron density map probability of finding an electron is highest at short distances from nucleus probability of finding an electron is highest at short distances from nucleus (b) calculated probability of finding an electron at certain distances from nucleus in the 1s orbital (b) calculated probability of finding an electron at certain distances from nucleus in the 1s orbital

58. 7.6 Quantum Numbers There are many solutions to Schroedinger ’ s equation for H atom There are many solutions to Schroedinger ’ s equation for H atom Each solution is a wave function called. Each solution is a wave function called Orbital. Each solution can be described with quantum numbers that describe some aspect of the solution. Each solution can be described with quantum numbers that describe some aspect of the solution. Schrödinger’s equation requires 3 quantum numbers

59. 7.6 Quantum Numbers Quantum numbers specify the properties of atomic orbitals and of electrons in orbitals Quantum numbers specify the properties of atomic orbitals and of electrons in orbitals the first three numbers come from the Schr ö dinger equation and describe: the first three numbers come from the Schr ö dinger equation and describe: main energy level main energy level shape shape orientation orientation 4 th describes state of electron 4 th describes state of electron

60. 1 st Quantum Number Principal Quantum Number: n Main energy level (or shell) occupied by electron. They are called atomic orbitals Main energy level (or shell) occupied by electron. They are called atomic orbitals regions where there is a high probability of finding an electron. values are all positive integers >0 (1,2,3, … ) values are all positive integers >0 (1,2,3, … ) As n increases size of orbital is larger size of orbital is larger electron has higher energy electron has higher energy the electron’s average distance from the nucleus increases the electron’s average distance from the nucleus increases

61. Principal Quantum Number Maximum number of electrons that can fit in an energy level: 2n 2

62. 1 st Quantum Number Energy

63. 2 nd Quantum Number Angular Momentum Quantum Number: l indicates the shape of the orbital (sublevel or subshell) indicates the shape of the orbital (sublevel or subshell) the number of possible shapes (or l values) for an energy level is equal to n the number of possible shapes (or l values) for an energy level is equal to n the possible values of l are 0 and all positive integers less than or equal to n - 1 the possible values of l are 0 and all positive integers less than or equal to n - 1

64. n l has integer values from 0 to n-1 n l = 0 is called s n l = 1 is called p n l =2 is called d n l =3 is called f n l =4 is called g

65. 2 nd Quantum Number s orbitalss orbitals: 1: s spherical l value of 0 1 st occur at n=1

66. 2 nd Quantum Number p orbitals: 3 2p x 2p x, 2p y, 2p z2p y2p z dumbbell- shaped l value of 1 1 st occur at n=2 for n>2, shape is same but size increases

67. 2 nd Quantum Number d orbitals: 5: 3d xz, 3d yz, 3d xy, 3d x2-y2, d z23d xz3d yz3d xy3d x2-y2d z2 mostly cloverleaf l value of 2 1 st occur at n=3 for n>3, same shape but larger size

68. 2 nd Quantum Number f orbitals: 7 types various shapes l value of 3 begin in n=4

69. 2 nd Quantum Number Other shapes can exist in energy levels as long as they follow the rules Other shapes can exist in energy levels as long as they follow the rules g (l=4) starts in 5 with 9 orbitals g (l=4) starts in 5 with 9 orbitals h (l=5) starts in 6 with 11 orbitals, etc h (l=5) starts in 6 with 11 orbitals, etc but no known elements have electrons in them at ground state but no known elements have electrons in them at ground state

70. 2 nd Quantum Number LevelSublevels 0 1 2 3 0 1 2 0 1 0

71. 3 rd Quantum Number Magnetic Quantum Number: m l indicates the orientation of an orbital around the nucleus indicates the orientation of an orbital around the nucleus has values from +l  -l has values from +l  -l specifies the exact orbital that the electron is contained in specifies the exact orbital that the electron is contained in each orbital holds maximum of each orbital holds maximum of 2 electrons total number of orbitals is equal to n 2 for an energy level total number of orbitals is equal to n 2 for an energy level number of possible m l values for a certain subshell is equal to 2l + 1 number of possible m l values for a certain subshell is equal to 2l + 1

72. 3 rd Quantum Number

73. Energy Level (n) Sublevels in Level # Orbitals in Sublevel Total # of Orbitals in Level 1s11 2s14 p3 3s19 p3 d5 4s116 p3 d5 f7

74. 4 th Quantum Number Spin Quantum Number: m s indicates the spin state of the electron indicates the spin state of the electron only 2 possible directions only 2 possible directions only 2 possible values: + ½ and - ½ only 2 possible values: + ½ and - ½ paired electrons must paired electrons must have opposite spins maximum number of maximum number of electrons in an energy level is 2n 2

75. 7.9 Polyelectronic Atoms Kinetic energy - as the electrons move around the nucleus Kinetic energy - as the electrons move around the nucleus Potential energy - from their attraction to nucleus Potential energy - from their attraction to nucleus Potential energy - from their repulsion to each other Potential energy - from their repulsion to each other

76. Electron Correlation Problem can ’ t find the exact location of electrons can ’ t find the exact location of electrons can ’ t find the specific repulsions between electrons can ’ t find the specific repulsions between electrons so we must treat each electron as if it has an average amount of attraction to nucleus and repulsion to other electrons so we must treat each electron as if it has an average amount of attraction to nucleus and repulsion to other electrons

77. Electron Shielding occurs when an electron is not attracted to the nucleus occurs when an electron is not attracted to the nucleus because of electrons in lower energy levels repelling it. because of electrons in lower energy levels repelling it.

78. Penetration Effect all orbitals in the same energy level do NOT have the same amount of energy ( are not degenerate) all orbitals in the same energy level do NOT have the same amount of energy ( are not degenerate) E s < E p < E d < E f the amount of energy in each sublevel is determined by its average distance from the nucleus the amount of energy in each sublevel is determined by its average distance from the nucleus

79. Section 5.2 Electron Arrangement in Atoms OBJECTIVES: OBJECTIVES: Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle. Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

81. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f aufbau diagram - page 133

82. Electron Configurations… l …are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: 1) Aufbau principle - electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! 2) Pauli Exclusion Principle - at most 2 electrons per orbital - different spins

83. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli To show the different direction of spin, a pair in the same orbital is written as:

84. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. 1) Principal quantum number 2) Angular momentum quantum number 3) Magnetic quantum number 4) Spin quantum number

85. Electron Configurations 3) Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. l Let’s write the electron configuration for Phosphorus  We need to account for all 15 electrons in phosphorus

86. l The first two electrons go into the 1s orbital Notice the opposite direction of the spins l only 13 more to go... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

87. l The next electrons go into the 2s orbital l only 11 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

88. The next electrons go into the 2p orbital only 5 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

89. The next electrons go into the 3s orbital only 3 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

90. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes (Hund’s) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3

91. Orbitals fill in an order l Lowest energy to higher energy. l Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. l However, half filled orbitals have a lower energy, and are next best Makes them more stable. Changes the filling order

92. Write the electron configurations for these elements: lTlTitanium - 22 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 lVlVanadium - 23 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 lClChromium - 24 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 (expected) BBut this is not what happens!!

93. Chromium is actually: l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 l Why? l This gives us two half filled orbitals (the others are all still full) l Half full is slightly lower in energy. l The same principal applies to copper.

94. Copper’s electron configuration l Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 l But the actual configuration is: l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 l This change gives one more filled orbital and one that is half filled. l Remember these exceptions: d 4, d 9

95. Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

96. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f aufbau diagram - page 133

97. Electron Configurations … … are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: … are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: 1) Aufbau principle - electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! 2) Pauli Exclusion Principle - at most 2 electrons per orbital - different spins

98. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli To show the different direction of spin, a pair in the same orbital is written as:

99. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. 1) Principal quantum number 2) Angular momentum quantum number 3) Magnetic quantum number 4) Spin quantum number

100. Electron Configurations 3) Hund ’ s Rule- When electrons occupy orbitals of equal energy, they don ’ t pair up until they have to. Let ’ s write the electron configuration for Phosphorus Let ’ s write the electron configuration for Phosphorus  We need to account for all 15 electrons in phosphorus

101. The first two electrons go into the 1s orbital Notice the opposite direction of the spins only 13 more to go... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

102. The next electrons go into the 2s orbital only 11 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

103. The next electrons go into the 2p orbital only 5 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

104. The next electrons go into the 3s orbital only 3 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

105. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes (Hund’s) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3

106. Orbitals fill in an order Lowest energy to higher energy. Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. However, half filled orbitals have a lower energy, and are next best However, half filled orbitals have a lower energy, and are next best Makes them more stable. Makes them more stable. Changes the filling order Changes the filling order

107. Write the electron configurations for these elements: Titanium - 22 electrons 1111s22s22p63s23p64s23d2 Vanadium - 23 electrons 1111s22s22p63s23p64s23d3 Chromium - 24 electrons 1111s22s22p63s23p64s23d4 (expected) BBBBut this is not what happens!!

108. Chromium is actually: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Why? Why? This gives us two half filled orbitals (the others are all still full) This gives us two half filled orbitals (the others are all still full) Half full is slightly lower in energy. Half full is slightly lower in energy. The same principal applies to copper. The same principal applies to copper.

109. Copper ’ s electron configuration Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 But the actual configuration is: But the actual configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This change gives one more filled orbital and one that is half filled. This change gives one more filled orbital and one that is half filled. Remember these exceptions: d 4, d 9 Remember these exceptions: d 4, d 9

110. Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

111. 7.12 Periodic Trends

112. Ionization Energy An electron can be removed from an atom if enough energy is absorbed (+) An electron can be removed from an atom if enough energy is absorbed (+) Ionization energy – the energy required to remove one electron from a gaseous neutral atom Ionization energy – the energy required to remove one electron from a gaseous neutral atom A (g) + energy  A + (g) + e - measurements of this are made on individual atoms in gas phase to avoid interactions with nearby atoms measurements of this are made on individual atoms in gas phase to avoid interactions with nearby atoms

113. Ionization Energy

115. if one electron is removed, the positive charge binds the electrons more tightly so 2 nd ionization energy must be higher if one electron is removed, the positive charge binds the electrons more tightly so 2 nd ionization energy must be higher the largest jump in energy is when you remove a core electron instead of valence the largest jump in energy is when you remove a core electron instead of valence

116. Ionization Energy

117. Across Period: Across Period: requires more energy to remove an electron so increases requires more energy to remove an electron so increases because electrons added in the same energy level do not shield electrons from nuclear charge because electrons added in the same energy level do not shield electrons from nuclear charge Down Group: requires less energy to remove electron so decreases because the valence electrons are farther away from protons attracting them

118. Ionization Energy

120. Electron Affinity Electron Affinity – the energy change when an electron is added to a gaseous neutral atom exothermic (-) exothermic (-) A + e -  A - + energy

121. Electron Affinity Across Period: Across Period: releases more energy so number increases (gets more negative) releases more energy so number increases (gets more negative) because electrons added in the same energy level do not shield electrons from nuclear charge because electrons added in the same energy level do not shield electrons from nuclear charge Down Group: releases less energy so number decreases (gets less negative) because the electrons being added are farther away from the attracting protons

123. Electron Affinity

124. Atomic Radii Defined by the edge of its orbital but since the edges are fuzzy, difficult to determine Defined by the edge of its orbital but since the edges are fuzzy, difficult to determine Atomic Radii – half the distance between the nuclei of identical atoms that are bonded together Atomic Radii – half the distance between the nuclei of identical atoms that are bonded together

125. Atomic Radii Across Period: Across Period: atoms get smaller atoms get smaller because of the increased number of protons attracting the electrons because of the increased number of protons attracting the electrons the electrons added in the same energy level do not shield electrons from nuclear charge the electrons added in the same energy level do not shield electrons from nuclear charge Down Group: atoms get larger increases because the energy levels being added to the atom

128. Atomic/Ionic Radii

129. Why is the periodic table shaped like it is and how are the elements arranged?

131. Elements are arranged according to atomic # and e - configuration. Li: 3 e - ’s 1s 2 2s 1 Na: 11 e - ’s 1s 2 2s 2 2p 6 3s 1 K: 19 e - ’s 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Paramagnetic or diamagnetic?

132. Valence orbitals: outer shell orbitals beyond the closest noble-gas configuration Valence electrons: “the ones that can react” (located in the valence orbitals). Elements in a vertical row have the same number of valence electrons. 2s 2 3s 2 4s 2 5s 2 6s 2 7s 2 The other e - ’s are called core electrons and don’t react.

135. Atomic sizes: Atomic size affects many properties, both physical and chemical

137. Li Be B C N O F Na Smaller SmallerSmaller WHY? K

139. Ionization Energy: The energy required to completely remove an e - from an atom in its gaseous state. Mg(g)  Mg 1+ + e - 1st ionization energy Mg 1+ (g)  Mg 2+ + e - 2nd ionization energy Question: Which of the above ionizations would have the highest ionization energy and why?

140. electron being lost: 1st 2nd 3rd 4th 5th 6th 7th

141. Increases

143. I.E. Overhead

144. Electron Affinites: The energy change that occurs when an electron is added to a gaseous atom. Cl(g) + e -  Cl - (g)  E = - 349 kJ/mol What does the negative value mean?

145. Electron affinity values

146. What is meant by metallic character?

147. Common Oxidation states: note the vertical similarities.

148. I 2 (s) Cl 2 (g) Br 2 (l) The Halogen Family:

149. LiK Na Alkali Metal Family