The Electronic Structures of Atoms Electromagnetic Radiation The wavelength of electromagnetic radiation has the symbol λ.

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1 CHAPTER 7 Atomic Structure Chapter 8 Atomic Electron Configurations and Periodicity 1 The Electronic Structures of Atoms Electromagnetic Radiation The wavelength of electromagnetic radiation has the symbol λ. Measured in units of distance such as m,cm, Å. 1 Å = 1 x 1-1 m = 1 x 1-8 cm The frequency of electromagnetic radiation has the symbol υ. Measured in units of 1/time - s -1 Electromagnetic Radiation For electromagnetic radiation the velocity is 3. x 1 8 m/s and has the symbol c. Thus c = λ υ for electromagnetic radiation. 3

2 Electromagnetic Radiation wavelength Visible light Amplitude wavelength Ultaviolet radiation Node 4 Electromagnetic Radiation Figure Wave motion: wave length and nodes 6

3 Electromagnetic Spectrum 7 Electromagnetic Radiation Example: What is the frequency of green light of wavelength 5 Å? c c = λν ν = λ -1 1 x 1 m (5 Å) = Å m/s ν = m 14-1 ν = s -7 m 8 Quantization of Energy An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA. Energy of radiation is proportional to frequency E = h ν h = Planck s s constant = 6.66 x 1-34 J s 9

4 Quantization of Energy E = h ν =hc/λ Light with large λ (small ν) ) has a small E. Light with a short λ (large ν) ) has a large E. 1 Electromagnetic Radiation Example: What is the energy of a photon of green light with wavelength 5 Å? What is the energy of 1. mol of these photons? E = ( ν = 5.77 x 1 E = E = h ν -34 s J s)( J per photon s -1 ) ( For 1. mol of photons)( photons : J per photon) = 31kJ/mol 11 Atomic Spectra and the Bohr Atom An emission spectrum is formed by an electric current passing through a gas in a vacuum tube (at very low pressure) which causes the gas to emit light. 1

5 Spectrum of White Light Figure Spectrum of Excited Hydrogen Gas Active Figure Atomic Spectra and the Bohr Atom Every element has a unique spectrum. Thus we can use spectra to identify elements. This can be done in the lab, stars, fireworks, etc. 15

6 Atomic Spectra and the Bohr Atom An absorption spectrum is formed by shining a beam of white light through a sample of gas. Absorption spectra indicate the wavelengths of light that have beenabsorbed. 16 Atomic Line Emission Spectra and Niels Bohr Bohr s s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the SHARP LINE EMISSION SPECTRA of excited atoms. Niels Bohr ( ) 196) 17 Atomic Spectra and Bohr One view of atomic structure in early th century was that an electron (e-) ) traveled about the nucleus in an orbit. + Electron orbit 1. Any orbit should be possible and so is any energy.. But a charged particle moving in an electric field should emit energy. End result should be destruction! 18

7 Atomic Spectra and the Bohr Atom In 1913 Neils Bohr incorporated Planck s quantum theory into the hydrogen spectrum explanation. Here are the postulates of Bohr s theory. 1. Atom has a number of definite and discrete energy levels (orbits) in which an electron may exist. (n principle quantum number) As the orbital radius increases so does the energy 1<<3<4< Atomic Spectra and the Bohr Atom. An electron may move from one discrete energy level (orbit) to another, but, in so doing, monochromatic radiation is emitted or absorbed. 3. An electron moves in a circular orbit about the nucleus Atomic Spectra and Bohr If e- s e s are in quantized energy states, then E E of states can have only certain values. This explain sharp line spectra. E = -C (1/ ) n = E = -C (1/1 ) n = 1 1

8 Energy Adsorption/Emission Active Figure 7.11 Atomic Spectra and the Bohr Atom Example: What is the wavelength of light emitted when the hydrogen atom s energy changes from n = 4 to n =? n = 4 and n = = R λ n1 n = m λ Atomic Spectra and the Bohr Atom The Rydberg equation is an empirical equation that relates the wavelengths of the lines in the hydrogen spectrum = R n1 n λ R is the Rydberg constant R = n < n 1 m n s refer to the numbers of the energy levels in the emission spectrum of hydrogen 7-1 4

9 Atomic Line Spectra and Niels Bohr Niels Bohr ( ) 196) Bohr s s theory was a great accomplishment. Rec d d Nobel Prize, 19 Problems with theory theory only successful for H. introduced quantum idea artificially. So, we go on to QUANTUM or WAVE MECHANICS 5 Quantum or Wave Mechanics L. de Broglie ( ) 1987) de Broglie (194) proposed that all moving objects have wave properties. For light: E = mc E = hν h = hc / λ Therefore, mc = h / λ and for particles (mass)(velocity) = h / λ 6 The Wave Nature of the Electron In 195 Louis de Broglie published his Ph.D. dissertation. -Electrons (in fact - all particles) have both a particle and a wave like character. This wave-particle duality is a fundamental property of submicroscopic particles. h λ = mv h = Planck s constant m = mass v = velocity of of particle particle 7

10 The Wave Nature of the Electron Example. Determine the wavelength, in m, of an electron, with mass 9.11 x 1-31 kg, having a velocity of 5.65 x 1 7 m/s. Remember Planck s constant is 6.66 x 1-34 Js which is also equal to 6.66 x 1-34 kg m /s. λ = λ = h mv ( kg)( m/s) λ = kg m s 11 m 8 Quantum or Wave Mechanics Schrodinger applied idea of e-e behaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION Solution gives set of math expressions called WAVE E. Schrodinger FUNCTIONS, Ψ Each describes an allowed energy state of an e-e Quantization introduced naturally. 9 Uncertainty Principle W. Heisenberg Problem of of defining nature of of electrons in in atoms solved by by W. Heisenberg. Cannot simultaneously define the position and momentum (= (= m v) m of of an an electron. We define e-e e-energy exactly but accept limitation that we do do not know exact position. 3

11 Quantum Numbers The principal quantum number has the symbol n. n = 1,, 3, 4,... shells n = K, L, M, N,... The electron s energy depends principally on n. 31 Quantum Numbers The angular momentum quantum number has the symbol l. l =, 1,, 3, 4, 5,...(n-1) l = s, p, d, f, g, h,...(n-1) l tells us the shape of the orbitals. 3 Quantum Numbers The symbol for the magnetic quantum number is m l. m l = - l, (- l + 1), (- l +),...,..., (l -), (l -1), l If l = (or an s orbital), then m l =. Notice that there is only 1 value of m l. This implies that there is one s orbital per n value. n 1 If l = 1 (or a p orbital), then m l = -1,,+1. There are 3 values of m l. Thus there are three p orbitals per n value. n 33

12 Atomic Orbitals Atomic orbitals are regions of space where the probability of finding an electron about an atom is highest. s orbital properties: There is one s orbital per n level. l = 1 value of m l 34 Atomic Orbitals s orbitals are spherically symmetric. 35 Atomic Orbitals p orbital properties: The first p orbitals appear in the n = shell. p orbitals are peanut or dumbbell shaped volumes. They are directed along the axes of a Cartesian coordinate system. There are 3 p orbitals per n level. The three orbitals are named p x, p y, p z. They have an l = 1. m l = -1,,+1 3 values of m l 36

13 p Orbitals When n =,, then l l = and 1 Therefore, in in n = shell there are types of of orbitals subshells For l l = m l l = this is is a s subshell For l l = 1 m l l = -1,,, this is is a p subshell with 3 orbitals Typical p orbital planar node When l = 1, there is a PLANAR NODE thru the nucleus. 37 Atomic Orbitals p orbitals are peanut or dumbbell shaped. 38 p Orbitals The three p orbitals lie 9 o apart in space 39

14 Atomic Orbitals d orbital properties: The first d orbitals appear in the n = 3 shell. The five d orbitals have two different shapes: 4 are clover leaf shaped. 1 is peanut shaped with a doughnut around it. The orbitals lie directly on the Cartesian axes or are rotated 45 o from the axes. There are 5 d orbitals per n level. The five orbitals are named dxy, dyz, dxz, d They have an l =. m = -,-1,,+1,+ 5 values of m l l x -y, d z 4 d Orbitals typical d orbital planar node s orbitals have no planar node (l = ) and so are spherical. p orbitals have l = 1, and have 1 planar node, and so are dumbbell shaped. This means d orbitals (with l = ) have planar nodes planar node Figure Atomic Orbitals d orbital shapes 4

15 Atomic Orbitals There are 7 f orbitals with l =3 ml = -3, -,-1,,+1,+, +3 7 values of m l These orbitals are hard to visualize 43 f Orbitals When n = 4, l =, 1,, 3 so there are 4 subshells in the shell. For l =, m l = ---> > s subshell with single orbital For l = 1, m l = -1,, > > p subshell with 3 orbitals For l =, m l = -, -1,, +1, + ---> > d subshell with 5 orbitals For l = 3, m l = -3, -, -1,, +1, +, > > f subshell with 7 orbitals 44 f Orbitals One of 7 possible f orbitals. All have 3 planar surfaces. Can you find the 3 surfaces here? 45

16 Quantum Numbers The last quantum number is the spin quantum number which has the symbol m s. The spin quantum number only has two possible values. m s = +1/ or -1/ m s = ± 1/ Wolfgang Pauli in 195 discovered the Exclusion Principle. No two electrons in an atom can have the same set of 4 quantum numbers. 46 Electron Spin Quantum Number, m s Can be proved experimentally that electron has a spin. Two spin directions are given by m s where m s = +1/ and -1/. 47 Electron Spin and Magnetism Diamagnetic:: NOT attracted to a magnetic field Paramagnetic: substance is attracted to a magnetic field. Substances with unpaired electrons are paramagnetic. 48

17 Atomic Orbitals Spin quantum number effects: Every orbital can hold up to two electrons. Consequence of the Pauli Exclusion Principle. The two electrons are designated as having one spin up and one spin down Spin describes the direction of the electron s magnetic fields. 49 The Periodic Table and Electron Configurations The principle that describes how the periodic chart is a function of electronic configurations is the Aufbau Principle. The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available. 5 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY 51

18 Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (m l ) 5 Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. That is, each electron has a unique address. 53 The Periodic Table and Electron Configurations The Aufbau Principle describes the electron filling order in atoms. 54

19 The Periodic Table and Electron Configurations To remember the correct filling order for electrons in atoms. 1. You can use this mnemonic. 55 Writing Atomic Electron Configurations Two ways of writing configs.. One is called the spdf notation. spdf notation for H, atomic number = s value of n no. of electrons value of l 56 Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. ORBITAL BOX NOTATION for He, atomic number = 1 s 1s Arrows depict electron spin One electron has n = 1, l =, m l =, m s = + 1/ Other electron has n = 1, l =, m l =, m s = - 1/ 57

20 The Periodic Table and Electron Configurations Now we will use the Aufbau Principle to determine the electronic configurations of the elements on the periodic chart. 1 st row elements. 1s Configuration 1 H He 1s 1s 1 58 The Periodic Table and Electron Configurations Hund s rule tells us that the electrons will fill the p orbitals by placing electrons in each orbital singly and with same spin until half-filled. Then the electrons will pair to finish the p orbitals Li Be B C N O F 1s Ne s p Configuration 1s 1s 1s 1s 1s 1s 1s 1s s s s s 1 s s s s p p 59 1 p p p p The Periodic Table and Electron Configurations Na Mg Al Si P S Cl Ar 3s 3p Configuration [ Ne] 1 [ Ne] 3s [ Ne] [ Ne] 3s [ Ne] 1 [ Ne] 3s 3p [ Ne] [ Ne] 3s 3p [ Ne] 3 [ Ne] 3s 3p [ Ne] 4 [ Ne] 3s 3p [ Ne] 5 [ Ne] 3s 3p [ Ne] 6 [ Ne] 3s 3p 6

21 The Periodic Table and Electron Configurations Now we can write a complete set of quantum numbers for all of the electrons in these three elements as examples. Na Ca Fe First for 11 Na. (remember Ne has 1 electrons) 3s 3p Configuration 1 [ Ne] [ Ne] 3 11Na s 61 The Periodic Table and Electron Configurations n l m l m s 1 e 1 st - e 1 nd - 3 e rd - 4 e th - 5 e th - 6 e th - 7 e th - 8 e th - 9 e th - 1 e th - 11 e 3 th / 1 s electrons 1/ + 1/ 1/ + 1/ + 1/ + 1/ 1/ 1/ 1/ + 1/ s electrons p electrons } 3 s electron 6 The Periodic Table and Electron Configurations Next we will do the same exercise for Ca. Again, when finished we must have one set of 4 quantum numbers for each of the electrons in Ca. We represent the first 18 electrons in Ca with the symbol [Ar]. 3d 4s 4p Configuration [ Ar] Ca [Ar] 4s 63

22 The Periodic Table and Electron Configurations n l m l m s [Ar ] 19 th e / 64 The Periodic Table and Electron Configurations n l m l m s [Ar ]19 th th e - e / 4 s electrons 1/ 65 66

23 67 Ion Configurations To form cations from elements remove 1 or more e-e from subshell of highest n [or highest (n + l)]. P [Ne[ Ne] ] 3s 3p 3-3e- ---> > P 3+ [Ne]] 3s 3p 3s 3p 3s 3p s p s p 1s 1s 68 Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar[ Ar] ] 4s 3d 6 loses electrons ---> > Fe + [Ar]] 4s 3d 6 Fe Fe + 4s 3d 4s 3d To form cations, always remove electrons of highest n value first! 4s Fe 3+ 3d 69

24 The Periodic Table and Electron Configurations Finally, we do the same exercise for 6 Fe. We should have one set of 4 quantum numbers for each of the 6 electrons in Fe. To save time and space, we use the symbol [Ar] to represent the first 18 electrons in Fe 3d Configurat ion 6 [ Ar] [ Ar] 4s 6 Fe 3d 4s 4p 7 More About the Periodic Table Noble Gases filled electron shells. similar electronic structures similar chemical reactions He 1s Ne [He] s p 6 Ar [Ne] 3s 3p 6 Kr [Ar] 4s 4p 6 Xe [Kr] 5s 5p 6 Rn [Xe] 6s 6p 6 71 More About the Periodic Table Representative Elements Groups 1,,13-18 These elements will have their outermost electron in an outer s or p orbital. regular variations in their properties - periodic 7

25 More About the Periodic Table d-transition Elements Each metal has d electrons. ns (n-1)d configurations Smaller variations from row-to-row than the representative elements. 73 More About the Periodic Table f - transition metals Sometimes called inner transition metals. Electrons are being added to f orbitals. Consequently, very slight variations of properties from one element to another. 74 Periodic Properties of the Elements Atomic Radii Atomic radii describes the relative sizes of atoms. Atomic radii increase within a column going from the top to the bottom of the periodic table. Atomic radii decrease within a row going from left to right on the periodic table. 75

26 Atomic Radii The reason the atomic radii decrease across a period is due to shielding or screening effect. Effective nuclear charge, Z eff, experienced by an electron is less than the actual nuclear charge, Z. The inner electrons block the nuclear charge s effect on the outer electrons. 76 Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. Explains why E(s) < E(p) Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* by --> > [ Z - (no. inner electrons) ] Charge felt by s e-e in Li Z* = 3 - = 1 Be Z* = 4 - = B Z* = 5 - = 3 and so on! 77 Effective Nuclear Charge Electron cloud for 1s electrons Figure 8.6 Z* is the nuclear charge experienced by the outermost electrons. 78

27 Atomic Radii on their atomic radii. Se, S, O, Te 79 Atomic Radii on their atomic radii. Se, S, O, Te O < S < Se < Te 8 Atomic Radii on their atomic radii. P, Cl, S, Si 81

28 Atomic Radii on their atomic radii. P, Cl, S, Si Cl < S < P < Si 8 Atomic Radii on their atomic radii. Ga, F, S, As 83 Atomic Radii on their atomic radii. Ga, F, S, As F < S < As < Ga 84

29 Ionization Energy First ionization energy (IE 1 ) Energy required to remove the first electron from an isolated gaseous atom to form a 1+ ion. Symbolically: Atom (g) + energy ion + (g) + e - Mg (g) + 738kJ/mol Mg + + e - 85 Ionization Energy Second ionization energy (IE ) The amount of energy required to remove the second electron from a gaseous 1+ ion. Symbolically: ion + + energy ion + + e - Mg kj/mol Mg + + e - Atoms can have 3 rd (IE 3 ), 4 th (IE 4 ), etc. ionization energies. 86 Ionization Energy Periodic trends for Ionization Energy: 1. IE > IE 1 It always takes more energy to remove a second electron from an ion than from a neutral atom.. IE 1 generally increases to the right Important exceptions at Be & Mg, N & P, etc. due to filled and halffilled subshells. 3. IE 1 generally decreases down 87

30 First Ionization Energies of Some Elements Ionization Energy (kj/mol) He F N C H Be O B Li Ne Ar Cl P Mg S Ca Si Na Al K Atomic Number 88 Ionization Energy on their first ionization energies. Sr, Be, Ca, Mg 89 Ionization Energy on their first ionization energies. Sr, Be, Ca, Mg Sr < Ca < Mg < Be 9

31 Ionization Energy on their first ionization energies. Al, Cl, Na, P 91 Ionization Energy on their first ionization energies. Al, Cl, Na, P Na < Al < P < Cl 9 Ionization Energy on their first ionization energies. B, O, Be, N 93

32 Ionization Energy on their first ionization energies. B, O, Be, N B < Be < O < N 94 Electron Affinity Energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge. Sign conventions for electron affinity. If electron affinity > energy is absorbed. If electron affinity < energy is released. Electron affinity is a measure of an atom s ability to form negative ions. Symbolically: atom(g) + e - + EA ion - (g) 95 Electron Affinity Two examples of electron affinity values: Mg (g) + e kj/mol Mg - (g) EA = +31 kj/mol Br (g) + e - Br - (g) + 33 kj/mol EA = -33 kj/mol 96

33 Electron Affinity General periodic trend for electron affinity is more negative from left to right more negative from bottom to top 97 Electron Affinity Electron Affinities of Some Elements Electron Affinity (kj/mol) He Be B N Ne Mg Al Na H Li O C Si F Ar Ca P K S Cl Atomic Number 98 Electron Affinity on their electron affinities. Al, Mg, Si, Na 99

34 Electron Affinity on their electron affinities. Al, Mg, Si, Na Si < Al < Na < Mg 1 Ionic Radii Cations are always smaller Element Li Be Atomic Radius (Å) Ion Li + Be + Ionic Radius (Å) Ionic Radii Anions are always larger Element N O F Atomic Radius(Å) Ion N 3- O - F 1- Ionic Radius(Å)

35 Ionic Radii Cation radii decrease from left to right across a period. Increasing nuclear charge attracts the electrons and decreases the radius. Ion Rb + Sr + In 3+ Ionic Radii(Å) Ionic Radii Anion radii decrease from left to right across a period. Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius. Ion N 3- O - F 1- Ionic Radii(Å) Trends in Ion Sizes Active Figure

36 Ionic Radii on their ionic radii. Ga 3+, K +, Ca + 16 Ionic Radii on their ionic radii. K 1+ < Ca + < Ga Ionic Radii on their ionic radii. Cl -1, Se -, Br -1, S - 18

37 Ionic Radii on their ionic radii. Cl -1, Se -, Br -1, S - Cl 1- < S - < Br 1- < Se - 19 Electronegativity measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element. Fluorine is the most electronegative element. Cesium and francium are the least electronegative elements. Increase from left to right across periods and decrease from top to bottom within groups. 11 Electronegativity on their electronegativity. Se, Ge, Br, As 111

38 Electronegativity on their electronegativity. Se, Ge, Br, As Ge < As < Se < Br 11 Electronegativity on their electronegativity. Be, Mg, Ca, Ba 113 Electronegativity on their electronegativity. Be, Mg, Ca, Ba Ba < Ca < Mg < Be 114

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