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4 Chapter 6. Electronic Structure of Atoms Chemical behavior of atoms is primarily determined by arrangement of electrons outside atomic nucleus. Introduce a model of the atom help us understand this arrangement, to understand why elements exhibit their characteristic kinds of chemical behavior Copyright Houghton Mifflin Company.All rights reserved. 1 4

5 Nature of Light Most of our understanding of electronic structure comes from analysis of light emitted or absorbed by substances Need to look at nature of visible light & other forms of radiant energy. Includes radio waves, X-rays, IR, microwaves, visible light, etc. & are different kinds of electromagnetic radiation exhibits wavelength-like behavior and travels through space at speed of light in vacuum. Copyright Houghton Mifflin Company.All rights reserved. 1 5

6 Waves To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation. The distance between corresponding points on adjacent waves is the wavelength ( ). Copyright Houghton Mifflin Company.All rights reserved. 1 6

7 Waves The number of waves passing a given point per unit of time is the frequency ( ). For waves travelling at the same velocity, the longer the wavelength, the smaller the frequency. Copyright Houghton Mifflin Company.All rights reserved. 1 7

9 Electromagnetic Radiation All waves are characterized by: 1. Wavelength ( ): distance between 2 peaks in wave. 2. Frequency (v): number of waves per second that pass given point in space (s -1 or hertz, Hz) 3. Speed (c): speed of light is x 10 8 m/s (same rate for all electromagnetic radiation in vacuum). Copyright Houghton Mifflin Company.All rights reserved. 1 9

10 FREQUENCY AND WAVELENGTH RELATED The wavelength, frequency and speed of electromagnetic radiation are all related by: = c displacement direction of This means we can.. propagat Calculate the wavelength ion From frequency and velocity. Copyright Houghton Mifflin Company.All rights reserved. 1 10

11 EXAMPLES Calculating the wavelength of electromagnetic radiation from its frequency and velocity. The FOX broadcasts at 99.3 MHz. What is the wavelength of the radiation? The frequency is 99.3 x 10 6 Hz = 99.3 x10 6 s -1 Radio waves are a form of electromagnetic radiation c c c = x 10 8 m/s We can also. ms 3.019m Copyright Houghton Mifflin Company.All rights reserved s 1 1

12 EXAMPLES Calculate the frequency of electromagnetic radiation from its wavelength and velocity. The wavelength of the yellow light from a sodium lamp is 589 nm. What is the frequency of the radiation? c 589nm nm 9 m m c ms s m Copyright Houghton Mifflin Company.All rights reserved. 1 12

14 The Particle Nature of Light The wave model of light explains many aspects of its behavior but not all It could not explain the emission of light from electronically excited gas atoms This is known as emission or atomic spectra Copyright Houghton Mifflin Company.All rights reserved. 1 14

15 Quantised Energy and Photons The wave nature of light does not explain how an object can glow when its temperature increases. Max Planck explained it by assuming that energy comes in packets called quanta. Copyright Houghton Mifflin Company.All rights reserved. 1 15

16 Quantum Theory Max Planck (1900): energy can be released or absorbed by atoms only in discrete amounts Quantum (fixed amount): smallest quantity of energy that can be emitted or absorbed E = nhv E = energy, n = + integer, v = frequency, h (Planck s constant) = x Js Copyright Houghton Mifflin Company.All rights reserved. 1 16

17 Quantised Energy and Photons Einstein used this assumption to explain the photoelectric effect. He concluded that energy is proportional to frequency: E = h where h is Planck s constant, Js. Copyright Houghton Mifflin Company.All rights reserved. 1 17

18 Quantised Energy and Photons Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c = E = h Copyright Houghton Mifflin Company.All rights reserved. 1 18

19 Photoelectric Effect E = hv = hc/ Example: Calculate the energy of one photon of yellow light whose wavelength is 589 nm. Which has more energy: X-rays or microwaves? Copyright Houghton Mifflin Company.All rights reserved. 1 19

21 Quantised Energy and Photons One does not observe a continuous spectrum as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed. Copyright Houghton Mifflin Company.All rights reserved. 1 21

23 Rydberg Equation Line spectrum of excited hydrogen atoms Rydberg equation - 1 = R H 1 n i 2 1 n f 2 R is the Rydberg constant = x 10 7 m -1 Copyright Houghton Mifflin Company.All rights reserved. 1 23

24 Emission Spectra Atoms give off light when heated or otherwise excited energetically Light given off by energetically excited atom is not continuous distribution of s Continuous spectrum: Contains all the wavelengths (& all energies) of light Line (discrete) spectrum: Contains only some of the wavelengths of light. Copyright Houghton Mifflin Company.All rights reserved. 1 24

25 Nuclear Model of the Atom Rutherford showed: Atomic nucleus is composed of protons (+) & neutrons (0). Nucleus is very small compared to size of entire atom. Questions left unanswered: How are electrons arranged & how do they move? Electrons are moving charged particles Moving charged particles give off energy Atom should constantly be giving off energy Electrons should crash into nucleus and atom collapse!! This led to Neil Bohr Model of the Atom Copyright Houghton Mifflin Company.All rights reserved. 1 25

26 Bohr Model Explained spectra of hydrogen: H atom has only certain allowable energy levels (stationary states) fixed circular orbits. Atom does not radiate energy while in one of its stationary states Atom changes stationary states by absorbing or emitting a photon whose energy equals difference in energy between two states (E = hv) Spectrum is not continuous energy has only certain states! Copyright Houghton Mifflin Company.All rights reserved. 1 26

27 The Bohr explanation of the series of spectral lines When excited hydrogen atoms return to lower energy states, they emit photons of certain energies & certain colors. Copyright Houghton Mifflin Company.All rights reserved. 1 27

28 Bohr s Model Atom s energy has only certain levels or states Atoms have minimum energy: ground state Higher energy levels: excited states Farther from nucleus, higher its energy Putting e- in excited state requires addition of energy Bringing e- back to ground state releases energy Only specific frequencies can be absorbed/emitted Copyright Houghton Mifflin Company.All rights reserved. 1 28

29 Quantised Energy and Photons Niels Bohr adopted Planck s assumption and explained these phenomena in this way: 1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies). 2. Electrons in permitted orbits have specific, allowed energies; these energies will not be radiated from the atom. 3. Energy is only absorbed or emitted in such a way as to move an electron from one allowed energy state to another; the energy is defined by: E = h Copyright Houghton Mifflin Company.All rights reserved. 1 29

30 Quantised Energy and Photons The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation: 1 1 E = R ( ) H - n f 2 n i 2 Note E = hv = hc/ where R H is the Rydberg constant, J, and n i and n f are the initial and final energy levels of the electron. Copyright Houghton Mifflin Company.All rights reserved. 1 30

31 Emission Spectrum of the Hydrogen Atom Bohr showed that the energies that the electron in the H atom can possess is given by the equation: E n R 2 n where R = Rydberg s constant (2.18 x J or cm -1 ) n = principal quantum number (1,2,3,..) For n = 1, E n becomes most negative and this corresponds to the most stable energy state called the ground state. n = 2,3, are called excited states and the corresponding energies are higher in magnitude. H Copyright Houghton Mifflin Company.All rights reserved. 1 31

32 Interaction of a photon and an atom results in the atom moving to a higher energy state; atom is excited. State of lowest energy called the ground state. E E 2 (excited state) E 1 (ground state) E 2 E 1 = E 12 = hν Copyright Houghton Mifflin Company.All rights reserved. 1 32

33 Absorption Spectroscopy Atoms selectively absorb photons of light. Spectra that results are termed line spectra as only specific energies are absorbed due to selective absorption. Consider the absorption spectrum of an atom that was measured to give the following spectrum: E E 2 (excited state) A E 1 (ground state) wavenumber(cm -1 ) Let us now calculate the difference in energy in joules. Copyright Houghton Mifflin Company.All rights reserved. 1 33

34 Use the equation: E = hcv Note: the unit for wavenumber is cm -1, but c is given in m s x 10 8 m s -1 c = x 100 cm 1 s 1 m = x cm s -1 Therefore E 12 = (6.626 x J s)(2.998 x cm s -1 )( cm -1 ) = x J Now, consider an atom with one electron and four allowed energy states Copyright Houghton Mifflin Company.All rights reserved. 1 34

35 E 4 E 3 E 2 E 1 How many lines do we see in the absorption? The answer is 3, at E 12, E 13 and E 14 Emission Spectroscopy Atom emits a photon of energy. Note: electron does not have to return directly to the ground state; instead can return stepwise down the energy ladder. Copyright Houghton Mifflin Company.All rights reserved. 1 35

36 1 E 4 E E E 2 In the above case, the emission spectrum would have 6 lines, as 6 different transitions can occur at 6 different energies. E 34 E 23 E 24 E 12 E 13 E 14 Intensity of Energy E 1 Energy Copyright Houghton Mifflin Company.All rights reserved. 1 36

37 The electron is initially in the excited state, n i ; during emission it drops to a lower energy state, n f.. The change in energy is given by: Therefore, OR E = E = E f - E i -RH n 2 f - -RH n i 2 OR 1 E = 1 -RH - n f 2 n i 2 OR hv = E = 1 1 RH - When a photon is emitted, n i > n f, E < 0 and when energy is absorbed n f > n i and E > 0. n i 2 n f 2 Copyright Houghton Mifflin Company.All rights reserved. 1 37

38 The various series in atomic H emission spectrum Series n f n i Region Lyman 1 2,3,4,. UV Balmer 2 3,4,5,. Visible Paschen 3 4,5,6,. IR Brackett 4 5,6,7,. IR Pfund 5 6,7,8,. IR Copyright Houghton Mifflin Company.All rights reserved. 1 38

39 Important Points of Bohr s Model 1. The energies of electrons (energy levels) in an atom are quantized. 2. Quantum numbers are necessary to describe certain properties of electrons in an atoms (such as energy & location). 3. An electron s energy increases with increasing distance from the nucleus 4. The discrete energies (lines) in the spectra of the elements result from quantized electronic energies. Neil Bohr won the the Nobel Prize (1922). Copyright Houghton Mifflin Company.All rights reserved. 1 39

40 Problems with the Bohr Model Only explains hydrogen atom spectrum and other 1 electron systems Neglects interactions between electrons small particle circling the nucleus Assumes circular or elliptical orbits for electrons - which is not true Can electrons also exhibit wave-like character? (and not just a small particle circling about the nucleus) Copyright Houghton Mifflin Company.All rights reserved. 1 40

41 Wave Behavior of Matter Radiation appears to have either a wavelike or a particle-like (photon) character. Louis de Broglie: Can the electron orbiting the nucleus of a H atom be thought of not as a particle but rather a wave with a characteristic wavelength? He posited that if light can have material properties, matter should exhibit wave properties. He demonstrated that the relationship between mass and wavelength is: = h mv Copyright Houghton Mifflin Company.All rights reserved. 1 41

42 The Wave Nature of Matter Heisenberg (uncertainty principle): impossible to know simultaneously both the exact momentum (mass times speed) of the electron & its exact location in space. x * m u h 4p Copyright Houghton Mifflin Company.All rights reserved. 1 42

43 The Uncertainty Principle Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely its position is known. In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself! Copyright Houghton Mifflin Company.All rights reserved. 1 43

44 Quantum Mechanics Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics. Copyright Houghton Mifflin Company.All rights reserved. 1 44

45 Quantum Mechanics The wave equation is designated with a lower case Greek psi ( ). The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time. Copyright Houghton Mifflin Company.All rights reserved. 1 45

46 Wave Mechanics Treats electrons as waves and uses wave equations to calculate probability densities of finding electron in particular region in atom solutions are wave functions or atomic orbitals cannot know precisely location of electron at any moment but can describe its probability for given energy level, we can depict probability with electron density diagram or electron cloud. Copyright Houghton Mifflin Company.All rights reserved. 1 46

47 Quantum-Mechanical Model of an Atom Electron density: describes distribution of electron in orbital High in those regions of orbital where probability of finding electron is high and low in regions where probability is low H atom (ground state): e - is almost always found within sphere with certain radius, (0.529 Å) centered about nucleus Copyright Houghton Mifflin Company.All rights reserved. 1 47

49 Orbitals and Quantum Numbers Each orbital describes a specific distribution of electron density in space (We called them Shells ) as given by the orbital s probability density. Each orbital has a characteristic energy & shape. Electrons occupy orbitals characterized by 4 Quantum Numbers (QN) Copyright Houghton Mifflin Company.All rights reserved. 1 49

50 1. Principal Quantum Number (n) Principal QN (n = 1, 2, 3,...): related to size and distance/energy of orbital Identifies how much energy electrons in orbital have Higher values mean higher energy & further distance from nucleus Less tightly bound to nucleus Copyright Houghton Mifflin Company.All rights reserved. 1 50

51 2. Angular Momentum QN (l) Angular Momentum QN (l = 0 to n - 1) relates to shape of the orbital. Each principal energy level contains one or more energy sublevels there are n sublevels in each principal energy level each type of sublevel has a different shape and energy Copyright Houghton Mifflin Company.All rights reserved. 1 51

52 Quantum Number, l Value of l Type of orbital s p d f

53 3. Magnetic Quantum Number, m l Describes the three-dimensional orientation of the orbital. Values are integers ranging from -l to l: l m l l. Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc. Copyright Houghton Mifflin Company.All rights reserved. 1 53

55 Observing the Effect of Electron Spin When a beam of H atoms is shot into a powerful magnetic field, the beam was split by the field into two beams - (opposing directions: half are attracted, half are repelled) Copyright Houghton Mifflin Company.All rights reserved. 1 55

56 4. Spin Quantum Number, m s In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. The spin of an electron describes its magnetic field, which affects its energy. This led to a fourth quantum number, the spin quantum number, m s. The spin quantum number has only 2 allowed values: +½ and ½. Copyright Houghton Mifflin Company.All rights reserved. 1 56

57 Electron Spin Electron Spin QN (m s = + 1 /2, - 1 /2): relates to spin states of electrons. Important when more than 1 electron is present Two e - in same orbital have opposite spins = lower energy Write 4 QNs for any electron Pauli Exclusion Principle: no two electrons in the same atom can have the same four QNs. Atomic orbital can hold a maximum of 2 electrons and they must have opposing spins. Copyright Houghton Mifflin Company.All rights reserved. 1 57

58 Quantum Numbers Name Symbol Values Meaning Principal n 1,2,3,4. Energy & distance Angular Momentum l 0,1,.,n-1 Shape l = s p d f Magnetic m l +l.-l Orientation Spinmagnetic m s +1/2, -1/2 Spin Copyright Houghton Mifflin Company.All rights reserved. 1 58

61 d Orbitals Value of l is 2 Figure 5.21 Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the centre Copyright Houghton Mifflin Company.All rights reserved. 1 61

62 Atomic Orbitals Relationship between quantum numbers and atomic orbitals. n l m l No. of orbitals AO designations s s 1-1,0,1 3 2p x, 2p y, 2p z s 1-1,0,1 3 3p x, 3p y, 3p z 2-2,-1,0,1,2 5 3d xy, 3d yz, 3d zx, 3d 2 x - y2, 3d 2 z Copyright Houghton Mifflin Company.All rights reserved. 1 62

67 What is maximum number of orbitals in n = 3? a) 1 b) 3 c) 4 d) 7 e) 9 Each shell has a total of n 2 orbitals. n=3 : l= 0 : 1 x 3s l=1 : 3 x 3p l = 2 : 5 x 3d Copyright Houghton Mifflin Company.All rights reserved. 1 67

68 All of the following sets of quantum numbers are allowed EXCEPT a) n = 3, l = 1, m l = -1 b) n = 2, l = 1, m l = 0 c) n = 5, l = 0, m l = -1 d) n = 4, l = 2, m l = +2 e) n = 1, l = 0, m l = 0 Copyright Houghton Mifflin Company.All rights reserved. 1 68

71 Which of the following sets of quantum numbers is not allowed? a) n = 4, l = 0, m l = 0, m s = +1/2 b) n = 5, l = 3, m l = 2, m s = -1/2 c) n = 2, l = 1, m l = -1, m s = -1/2 d) n = 3, l = 2, m l = 0, m s = +1/2 e) n = 1, l = 0, m l = 0, m s = 0 Copyright Houghton Mifflin Company.All rights reserved. 1 71

73 Energies of Orbitals As the number of electrons increases, so does the repulsion between them. Therefore, in manyelectron atoms, orbitals on the same energy level are no longer degenerate. Copyright Houghton Mifflin Company.All rights reserved. 1 73

74 Factors Affecting Atomic Orbital Energies Electrostatic effects play major role in determining energy states of many e - atoms caused by nucleus-electron attractions and electron-electron repulsions. reasons for differences in energy: nuclear charge, shielding and shape Nuclear charge (Z): Higher nuclear charge lowers orbital energy (stabilizes system) by increasing nucleus-electron attractions. Copyright Houghton Mifflin Company.All rights reserved. 1 74

75 Shielding Effect of Electron Repulsions (shielding): electron feels repulsion from other e - s All electrons located between given electron & nucleus screen, or shield, that electron from full attractive force of nucleus. Effective nuclear charge (Z eff ): nuclear charge electron actually experiences, thus making it easier to remove. Inner electrons shield outer electrons very effectively. Copyright Houghton Mifflin Company.All rights reserved. 1 75

76 Electron Configurations Distribution of all electrons in an atom. Consist of: Number denoting the energy level. Letter denoting the type of orbital. Superscript denoting the number of electrons in those orbitals. Copyright Houghton Mifflin Company.All rights reserved. 1 76

77 Electron Configuration States how many e - an atom has in each of its orbitals (or how many e - of various energies) A shorthand system of symbols 1s 2 2s 2 2p 4 : indicates n, l, For a many electron atom, build-up the energy levels, filling each orbital in succession by energy Copyright Houghton Mifflin Company.All rights reserved. 1 77

78 Orbital Diagrams Each box represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the spin of the electron. Copyright Houghton Mifflin Company.All rights reserved. 1 78

79 Orbital Diagrams Arrangement of electrons can also be specified in terms of orbital occupancy Orbital diagram: shows e - occupancy of each orbital about nucleus of atom Hund s rule: every orbital in sublevel is singly occupied with one electron before any orbital is doubly occupied And all electrons in singly occupied orbitals have same spin Copyright Houghton Mifflin Company.All rights reserved. 1 79

81 Hund s Rule The electron configuration of carbon is 1s 2 2s 2 2p 2. The different ways of placing 2 electrons in the three p orbitals are as follows: 2p x 2 2p y 2p z 2p x 1 2p y 1 2p z 2p x 1 2p y 1 2p z None of these arrangements violate Pauli Exclusion Principle, but we must determine which one will afford the greatest stability. The answer is provided by Hund s Rule which states that the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. Thus the orbital diagram for carbon is: 1s 2 2s 2 2p x 1 2p y 1 2p z Copyright Houghton Mifflin Company.All rights reserved. 1 81

82 The box below represents the 1s atomic orbital: 1s 1 The Pauli Exclusion Principle Statement: no two electrons in an atom can have the same four quantum numbers. Only two electrons may occupy the same atomic orbital, but they must have opposite spins. Consider the He atom which has two electrons. There are 3 different ways of placing the two electrons in the 1s or 1s 2 1s 2 1s 2 (a) (b) (c) Copyright Houghton Mifflin Company.All rights reserved. 1 82

83 (a) and (b) do not conform to Pauli Exclusion Principle; only (c) acceptable. Diamagnetism and Paramagnetism Paramagnetic substances are attracted by a magnet; an odd number of electrons must be present. On the other hand, if the electron spins are paired, magnetic effects cancel out and the atom is referred to as diamagnetic. He Li 1s 2 diamagnetic 1s 2 2s 1 paramagnetic Copyright Houghton Mifflin Company.All rights reserved. 1 83

84 Electronic Configurations and the Periodic Table Properties of the elements repeat themselves in a regular manner Groups have similar chemical properties due to similar e - configurations (outer e - configs) Chemical properties repeat in a regular manner because e - configs repeat. Let s look at the electronic configurations for Group 1A elements Li, Na, K Copyright Houghton Mifflin Company.All rights reserved. 1 84

85 Electronic Configurations and the Periodic Table Elements in Group 1A have one outer s electron causing similar chemical properties No. of outer electrons = group number Much of the chemical reactivity of elements is due to the No. of outer electrons (valence e-) Electron configuration can also be determined from the periodic table Copyright Houghton Mifflin Company.All rights reserved. 1 85

86 Highest energy level called the valence orbital or shell electrons in valence shell are valence electrons involved in forming compounds outer electrons: highest quantum number (n) valence and outer e - are same for main group elements but not transition metals. electrons not in valence shell: core electrons Copyright Houghton Mifflin Company.All rights reserved. 1 86

87 Increasing energy The Effect of Orbital Shape on Orbital Energy Shielding and Penetration causes an energy level to split into sublevels of different energy Order of sublevel energies: s < p < d < f Penetration is greater for 2s than 2p reduces electrostatic attraction 05_19 2 7s 6s 5s 4s 3s 2s 1s Subshell electron capacity d 6p 5d 5p 4d 4p 3d 3p 2p 5f 4f Copyright Houghton Mifflin Company.All rights reserved. 1 87

88 Aufbau Aufbau principle: e - normally occupy e - sublevels in an atom in order of increasing energy Energies of orbitals in different levels often overlap (4s is lower than 3d, 5s lower than 4d) 05_20 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Copyright Houghton Mifflin Company.All rights reserved. 1 88

90 Electronic Configurations and the Periodic Table 1. Begin with H and He and continue in order of increasing atomic number 2. As you move across each period a. add e - to the ns sublevel b. add e - to the np sublevel c. add e - to the (n-1)d sublevel d. add e - to the (n-2)f sublevel 3. Continue until you reach your element Copyright Houghton Mifflin Company.All rights reserved. 1 90

92 Noble Gas Configuration One can often use symbol of previous noble gas to represent core electrons 1s 2 2s 2 2p 6 = [Ne] Also called a condensed electron configuration Copyright Houghton Mifflin Company.All rights reserved. 1 92

94 What ( 1) ion (X - ) has following electron configuration? a) Na - b) Ar - c) Br - [Ne] = 1s 2 2s 2 2p 6 d) K - e) F - Fluorine is normally 1s 2 2s 2 2p 5,but with the addition of an electron to make F -, the config. is: 1s 2 2s 2 2p 6 Write the electron config. for Li, Na, K when finished Copyright Houghton Mifflin Company.All rights reserved. 1 94

95 Select the correct set of quantum numbers (n, l, m l, m s ) for the highest energy electron in the ground state of potassium, K. Write out the electron configuration: K = [Ar] 4s 1 A. 4, 1, -1, ½ B. 4, 1, 0, ½ C. 4, 0, 1, ½ D. 4, 0, 0, ½ E. 4, 1, 1, ½ Copyright Houghton Mifflin Company.All rights reserved. 1 95

96 Practice Determine the electron configuration and orbital diagram for the following: He, C, S, Co, Kr What are 4 quantum numbers for the last electron added? Copyright Houghton Mifflin Company.All rights reserved. 1 96

98 Some Anomalies For instance, the electron configuration **for chromium is: [Ar] 4s 1 3d 5 rather than the expected: [Ar] 4s 2 3d 4. This occurs because the 4s and 3d orbitals are very close in energy. These anomalies occur in f-block atoms, as well. Copyright Houghton Mifflin Company.All rights reserved. 1 98

99 Some Anomalies For instance, the electron configuration **for chromium is: [Ar] 4s 1 3d 5 rather than the expected: [Ar] 4s 2 3d 4. This occurs because the 4s and 3d orbitals are very close in energy. These anomalies occur in f-block atoms, as well. Copyright Houghton Mifflin Company.All rights reserved. 1 99

100 Exceptions- Cr and Cu One electron in the s -subshell Cr (Z=24) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cr [Ar]4s 1 3d 5 <=All s and d subshells are half full Cu ( Z= 29) Cu [Ar]4s 1 3d 10 <=Prefers a filled d subshell, leaving s with one electron Copyright Houghton Mifflin Company.All rights reserved

101 Unusual Configurations Unusual configuration for some transition metals: Cr, Cu, Ag, Au, Mo, Nb Half-filled and filled sublevels (Noble gases) are unexpectedly stable Copyright Houghton Mifflin Company.All rights reserved

102 Which of the following two electronic configuration is more stable? a [Ar]4s 1 3d 5 b [Ar]4s 2 3d 4 This is the electronic configuration for chromium. Copyright Houghton Mifflin Company.All rights reserved

103 Which of the following electron configurations represents the ground state for an element? A. [Ne]3s 1 3p 1 B. [He]2s 1 2p 3 C. [Ne]3s 2 3p 2 3d 1 D. [Ne]3s 2 3p 3 3d 1 E. [Ne]3s 2 3p 3 Copyright Houghton Mifflin Company.All rights reserved

106 Valence electron configuration Consider the electron configuration of Br: Br =[Ar] 4s 2 3d 10 4p 5 Although the 3d electrons are outer- shell electrons, they are not involved in bonding and are not considered valence electrons. We do not consider completely filled d and f subshells to be valence electrons. Copyright Houghton Mifflin Company.All rights reserved

107 An element with the electron configuration [Ar]ns 2 (n - 1)d 10 np 4 has valence (or outer) electrons. A. 2 B. 4 C. 6 D. 8 E. 16 X =[Ar] 4s 2 3d 10 4p 4 X =[Ar] 3d 10 4s 2 4p 4 Copyright Houghton Mifflin Company.All rights reserved