Lecture outline: Chapter 6 Electronic structure of atoms 1. Radiant energy 2. Quantum effects 3. The Bohr atom, orbitals 4. Many electron systems 1
Electronic structure of atoms Understanding the arrangement of electrons in atoms is the key to understanding ng the reactivity ty of atoms and molecules Total number of electrons in atom Locations of electrons in space Energy states of electrons 2
The atom is mostly empty space! Nucleus (proton(s) and neutron(s) ~10-18 cm electron: orbits the nucleus multielectron atoms have more than one e- orbiting the nucleus ~10-13 cm Diameter of atom ~ 10-8 cm (~1-5 Å) 3
The superstars of chemical physics Max Planck Albert Einstein Niels Bohr Werner Heisenberg Erwin Schrodinger 4
Electromagnetic radiation Carries energy Many types Moves Wavelike character 5
Properties of waves Regular rise and fall pattern Repeating periodicity Peak maxima and peak minima Wavelength (λ, lambda): the distance from one peak maximum to the next Frequency (ν, nu): the number of peak maxima that are passed per unit time Amplitude: the height of the peak maxima and minima from the central axis λ Amplitude 6
Some waves with different wavelengths λ 1 and amplitudes λ 2 λ 3 7
Electromagnetic waves Carry energy Electrical and magnetic components Classified based on wavelength Move at a constant speed (c) No propagating medium needed 8
1 7/8 1/8 3/4 1/4 1.0 sec/turn 5/8 3/8 1/2 Frequency (ν, nu): the number of peak maxima that are passed per unit time as the wave propagates ν = 16 maxima /second = 16s -1 = 16 hz 3.0 x 10 8 m ν = 8 maxima /second = 8s -1 =8hz 3.0 x 10 8 m ν = 4 maxima /second = 4s -1 = 4 hz 9
Proportionality Two variables are proportional if a change in the value of one results in a change in the value of the other, and if the two values are related mathematically by a constant t ( k ) Directly proportional: y x y = k x Inversely proportional: x y 1 x y = k x 1 y x = k 1 y 10
What is the difference between AM and FM radio? 11
Electromagnetic Spectrum http://commons.wikimedia.org/wiki/file:em_spectrumrevised.png This file is licensed under the Creative Commons Attribution-Share Alike 3.0 Unported license. 12
http://commons.wikimedia.org/wiki/file:em_spectrum_properties_edit.svg Author Inductiveload, NASA GNU Free Documentation License, Version 1.2 13
Unit commonly used for wavelengths Unit Symbol Meaning Radiation (10 n ) type meter m 1 TV, radio centimeter cm 10-2 microwave millimeter mm 10-3 infrared micrometer μm 10-6 infrared nanometer nm 10-9 UV, visible Angstrom Å 10-10 X-ray, gamma ray http://commons.wikimedia.org/wiki/file:electromagnetic-spectrum.svg Author: Victor Blacus, licensed under the Creative Commons Attribution-Share Alike 3.0 Unported license 14
How many peak maxima does red light with a wavelength of 649 nm pass in one second? In other words, what is the frequency of 649 nm light? c = λν 15
What is the wavelength of Q92.1 FM radiowaves? c=λν 16
Quantum effects: Planck 1900: Max Planck Energy is released/absorbed only in discrete units, packets, or quanta A quantum is the smallest quantity or packet of a given form of electromagnetic radiation Energy is quantized E = hν 17
Quantum analogies 18
What is the energy associated with 1 quantum of: (1) () 649 nm red light (ν = 4.62 x 10 14 Hz) (2) a Q92.1 radiowave (ν = 92.1 MHz) (3) a medical x-ray (ν = 9.55 x 10 17 Hz) E = hν 19
Photons: Einstein An explanation for a phenomenon known as the photoelectric effect e - Red light e - e - Blue light e - Light energy absorbed e - Sodium metal Sodium metal 20
Photons: Einstein An explanation for a phenomenon known as the photoelectron effect Electromagnetic radiation behaves as if composed of a stream of particles, or photons The energy of one photon = the energy of one quantum 21
The dual nature of electromagnetic radiation Light behaving as a wave: Light behaving as a particle: Light behaving as both a wave and a particle: 3.0 x 10 8 m in one second 22
The dual nature of light for three wavelengths λ 1 ν = 16 maxima /second = 16s -1 = 16 hz λ 2 ν = 8 maxima /second = 8s -1 = 8h hz λ 3 ν = 4 maxima /second = 4s -1 =4hz 1 second 23
Both matter and energy are quantized (exist only in discrete units) Matter: 1 H atom 1.5 H atoms 2 H atoms 2.3 H atoms 3 H atoms 3.7 H atoms Energy: 1 photon 1.5 photons 2 photons 3 photons 2.3 photons 3.7 photons 24
What is the energy associated with 1 mol of photons of: (1) 649 nm red light (ν = 4.62 x 10 14 Hz) E = nhν (2) a Q92.1 radiowave (ν = 92.1 MHz) (3) a medical x-ray (ν = 9.55 x 10 17 Hz) 25
Niels Bohr: a model for the hydrogen atom Based on a phenomenon of atoms and light known as line spectra 26
Monochromatic vs. polychromatic light Monochromatic light: light (radiation) of a single wavelength (ex.: laser light) Visible light: composed of light (radiation) of many different wavelengths 27
Scattering of visible light through a prism commons.wikimedia.org/wiki/file:arcoiris_high_contrasted_and_filtere d.jpg This work has been released into the public domain by its author, I, Alfredo55. http://missionscience.nasa.gov/ems/03_behaviors.html Credit: D-Kuru/Wikimedia Commons, licensed under the Creative Commons Attribution-Share Alike 3.0 Austria license. 28
Attribution: User Kieff, http://en.wikipedia.org/wiki/user:kieff/gallery http://commons.wikimedia.org/wiki/file:light_dispersion_conceptual_waves.gif Released into public domain 29
What would happen to the light of a laser if you shined it through a prism? http://commons.wikimedia.org/wiki/file:laser_pointers.jpg Author Netweb, Creative Commons Attribution-Share Alike 3.0 Unported license. 30
Electron excitation and relaxation gives rise to all of the color that we observe! Source light Reflected light How does it work???? 31
The heat of the flame excites an electron to a higher energy state. When the electron relaxes back to the ground state, energy is released as visible light. Li Na K Cu Pb 32
Scattering of the light emitted by an + excited atom through a prism H 2 Emitted light E = hν = hc λ - Prism Line spectra: 410 nm 486 nm 656 nm 434 nm 33
Bohr s postulates 1. Only electron orbits of certain energies are allowed (the energy of an e - is quantized) 2. An electron in a permitted orbit has a specific energy (an allowed energy state ) 3. An electron in an allowed state is stable and will not radiate energy 34
Some simple models to illustrate the idea of quantized orbits and electron energies n = 3 hνν n=4 E = = hc λ electron absorbs radiant energy nucleus n = 1 n=2 E in n = 1 is the ground state allowed orbit n > 1 is an excited state allowed orbit n = 1 n = 2 electron absorbs radiant energy n = 3 n=4 35
Some simple models to illustrate the idea of quantized orbits and electron energies n = 3 E hνν n=4 E = = hc λ electron emits radiant energy nucleus n = 1 n=2 n = 1 is the ground state allowed orbit n > 1 is an excited state allowed orbit n = 3 E out n = 2 electron emits radiant n = 1 energy n=4 36
The energy of an orbit is referenced relative to the electrons zero point energy, the point where the electron has been completely removed from the atom ΔE = E final E initial Increasing E ΔE 2 3 Zero point E = 0 n = 6 n = 5 n = 4 ΔE 3 4 n = 3 n = 2 First excited state n = n = 8 n = 7 ΔE (+) for removing an electron from the atom s ground state orbit ΔE 1 2 n = 1 Ground state E = (-) 37
An electron can jump (or relax) from any one orbit to another ΔE = E final E initial Increasing E ΔE 1 4 Zero point E = 0 n = 2 n = 6 n = 5 n = 4 n = 3 First excited state n = n = 8 n = 7 ΔE (+) for removing an electron from the atom s ground state orbit n = 1 Ground state E = (-) 38
The energy of an orbit is referenced relative to the electrons zero point energy, the point where the electron has been completely removed from the atom ΔE = E final E initial Increasing E ΔE 2 3 Zero point E = 0 n = 6 n = 5 n = 4 ΔE 3 4 n = 3 n = 2 First excited state n = n = 8 n = 7 ΔE (-) for placing an external electron in the atom s ground state orbit ΔE 1 2 n = 1 Ground state E = (-) 39
Energy levels in the Bohr atom nucleus n = 1 n = 2 n = 3 n = 4 E = (-R )( 1 n H 2 n R = 218x10-18 H 2.18 10 J ) ΔE =E final E initial n = 4 n = 2 n = 3 1 1 (R )( ) = H 2 n n ΔE E = 2 i f hν n = 1 40
When an excited electron in the n=4 orbit relaxes directly to the ground state orbit (n = 1), what wavelength of energy is released? Zero point E = 0 n = 6 Increasing E n = 5 n = 4 ΔE 1 4 n = 3 n = 2 n = 8 n = 7 λ = c ν 1 1 ΔE = (RH )( ) = 2 2 n n i f n = hν n = 1 Ground state E = (-) 41
When an excited electron in the n=4 orbit relaxes directly to the ground state orbit (n = 1), what wavelength of energy is released? 1 1 ΔE = (RH )( ) = hν c 1 2 2 λ = En = (-RH )( ) 2 n n n i f ν 42
Electron excitation and relaxation gives rise to all of the color that we observe! Source light Reflected light What is the difference between a normal incandescent lightbulb and a halogen lightbulb? 43
Summary of energy and matter Radiant energy has wavelike properties Radiant energy is quantized (can only exist in discrete packets) Radiant energy has particle like properties Matter (sp., the e-) has particle like properties The energies/orbits of matter are quantized Does matter have wavelike properties? 44
The debroglie wavelength of matter h λ = mv velocity Note: lower case v is velocity The greek letter nu is ν, which is frequency Don t confuse the two! 45
What is the wavelength of a baseball (120 g) travelling at a speed of 100 mph (44.7 m/s)? λ = h mv 46
Heisenberg The dual nature of matter (a particle and a wave) places limitations on the preciseness with which we can know both the location and the momentum (mass x velocity) of an object Know location precisely, then momentum is uncertain Know momentum precisely, then location is uncertain How does this apply to the electron? 47
Compare the sizes and wavelengths of the following moving obects: λ = h mv object mass velocity diameter λ ratio λ to diameter baseball 0.12 kg 45 m/s 0.08 m 1.24 x 10-34 m 1.5 x 10-33 earth 6.0 x 10 24 kg 2.98 x 10 4 m/s 1.276 x 10 7 m 3.71 x 10-63 m 2.9 x 10-70 electron 9.11 x 10-31 kg 1 x 10 6 m/s 10-20 m 1.22 x 10-10 m 1.2 x 10 10 48
What does all of this mean for electronic structure??!? Schrodinger: The behavior of the electron is better described by focusing on it s wavelike properties p An orbit: a defined, known pathway An orbital: a probability function; the probability that an electron will be found in a given location A description of the distribution of electron density in space Orbitals have a characteristic shape and energy 49
The n = 1 orbital Each dot represents a position where an electron may be found at any given moment with respect to the nucleus, which is at the center of the axes 50
An overview of the quantum mechanical model of the atom Electrons reside in areas of space called orbitals Orbitals have defined energies, shapes and sizes n = principal p quantum number = shell Subshells are energy levels within shells that have defined shapes (s, p, d, f) Orbitals of a given shape (within a subshell) have a specific orientation in space 51
Orbital quantum numbers n (principal) describes the energy of the orbital n = 1,2,3,4 shells l (azimuthal) describes the 3-dimensional i shape of fthe orbital (subshells) l values for a given n n are integers from 0 to n-1 m l (magnetic) describes the orientation of an orbital in space m l values for a given subshell are integers from -l to +l (2l + 1 possible orbitals for each subshell) A specific orbital is defined by specific and unique values for n, l, and m l 52
The n = 1 orbital n Allowed values for l: integers from 0 to n-11 Allowed values for m l : integers from -l to +l n = 2 n = 3 = 4 Shell Subshell Orbital Orbital Orbital Orbital # (n) # (l) # (m l ) name shape orientation n=1 1 0 0 1s spherical symmetric z y x s n = 1 l = 0 m l = 0 There is only one orbital in n = 1 (1,0,0) 53
The n = 2 orbitals n Allowed values for l: integers from 0 to n-11 Allowed values for m l : integers from -l to +l n = 2 n = 3 = 4 Shell # (n) Subshell # (l) Orbital # (m l ) Orbital name Orbital shape p n=1 Orbital orientation 2 0 0 2s spherical Symmetric with a wave node 2 1-1 2p y Dumbbell/ elipsoid about node centered on y axis 2 1 0 2p z Dumbbell/ about node centered on z axis elipsoid 2 1 1 2p x Dumbbell/ elipsoid about node centered on x axis There are four orbitals in n = 2 (2,0,0; 2,1,-1; 2,1,0; 2,1,1 ) 54
The four n = 2 orbitals n = 4 n = 3 2s 2p y 2p z 2p x n = 2 1s n = 1 55
The four n = 2 orbitals 2s 2p y 2p z 2p x z z y y z y z y x x x x 2s spherical n = 2 l = 0 m l = 0 2p y 2p z 2p x Dumbbell (elipsoid) n = 2 n = 2 n = 2 l = 1 l = 1 l = 1 m l = -1 m l = 0 m l = 1 56
2p orbitals n energy shell l shape subshell m l orientation orbital n = 2 n = 2 n = 2 l = 1 l = 1 l = 1 m l = 1 z y m l = 0 z y m l = -1 z y x x x p x p z p y 57
Compare the 1s and 2s orbitals z z y node y x x 1s 2s node 58
The n = 3 orbitals n Allowed values for l: integers from 0 to n-11 Allowed values for m l : integers from -l to +l n = 2 n = 3 = 4 Shell Subshell Orbital Orbital Orbital Orbital orientation # (n) # (l) # (m l ) name shape 3 0 0 3s spherical Symmetric with 2 wave nodes 3 1-1 3p y Dumbbell about node centered on y axis p y 3 1 0 3p z Dumbbell about node centered on z axis 3 1 1 3p x Dumbbell about node centered on x axis n=1 3 2-2 3d xy 4 pears four quadrants of xy plane 3 2-1 3d yz 4 pears four quadrants of yz plane 3 2 0 3d z2 2 pears to one torus 2 pears in z axis, torus to pears 3 2 1 3d xz 4 pears Centered about and on x and y axes 3 2 2 3d x2-y2 4 pears four quadrants of xy plane 59
The nine n = 3 orbitals 3s, 3p y, 3p z, 3p x, 3d xy, 3d yz, 3d z2, 3d xz, 3d x2-y2 l = 0 l = 1 l = 2 n=4 n = 3 n = 2 n = 1 60
The nine n = 3 orbitals 3s, 3p y, 3p z, 3p x, 3d xy, 3d yz, 3d z2, 3d xz, 3d x2-y2 l = 0 l = 1 l = 2 3s (cutaway 3p y 3p z 3p x view) 3d xy 3d yz 3d z2 3d xz 3d x2-y261
The five 3d orbitals y x z y x x z y z 3dxz 3dxy 3dyz x z x y 3dx2-y2 y z 3dz2 62
The 1s, 2s, and 3s orbitals n = 1 n = 2 n = 3 l = 0 l = 0 l = 0 m l = 0 m l = 0 m l = 0 63
06 0.6 0.5 1s radial pro obability 0.4 03 0.3 0.2 2s 0.1 3s 00 0.0 0 5 10 15 20 25 30 Distance from nucleus 64
The n = 4 orbitals n Allowed values for l: integers from 0 to n-11 Allowed values for m l : integers from -l to +l n = 2 n = 3 = 4 Shell # (n) Subshell # (l) Orbital # (m l ) Orbital name Orbital shape n=1 Orbital orientation 4 0 0 4s spherical Symmetric with 3 wave nodes 4 1-1, 0, 1 4p y, 4p z, 4p x Dumbbell Centered about and on x, y and z axes 4 2-2, -1, 4d xy,, 4d yz,, 4 pears and As for 3d orbitals 0, 1, 2 4d z2, 4d xz, 2 pears to 4d x2-y2 one torus 4 3-3, -2, - 1, 0, 1, 2, 3 4f z3, 4f xz2, 4f yz, 4f xyz, 4f 2 2 z(x y ), 4f x(x 3y ), 4f ( 2 3 2 ) Pretty complicated! Pretty complicated! 4f y(3x 2 y 2 ) 65
The sixteen n = 4 orbitals 4s, 4p y, 4p z, 4p x, 4d xy, 4d yz, 4d z2, 4d xz, 4d x2-y2 4f x 7 l = 0 l = 1 l = 2 l = 3 n = 4 n = 3 n = 2 n = 1 66
The seven f orbitals: difficult to draw, rarely encountered in your daily life 67
0 to n-1 allowed values -l to +l allowed values for each value of l n l Subshell m l value # of orbitals #oforbitals orbitals value value name in subshell in shell 1 0 1s 0 1 1 2 0 2s 0 1 4 2 1 2p -1, 0, 1 3 3 0 3s 0 1 9 3 1 3p -1, 0, 1 3 3 2 3d -2, -1, 0, 1, 2 5 4 0 4s 0 1 16 4 1 4p -1, 0, 1 3 4 2 4d -2, -1, 0, 1, 2 5 4 3 4f -3, -2, -1, 0, 1, 2, 3 7 68
Principal quantum Number of subshells (l) Subshell name (type of number n (shell) orbital) m l 1 1 s 2 2 s, p 3 3 s, p, d 4 4 s, p, d, f 5 5 s, p, d, f, g Type of subshell Number of orbitals in that subshell s 1 p 3 d 5 f 7 69
Atomic orbitals summary n = 1: One orbital 1s n = 2: Four orbitals: 2s 2p x, 2p y, 2p z n = 3: nine orbitals: 3s 3p x, 3p y, 3p z 3d xy, 3d yz, 3d z2, 3d xz, 3d x2-y2 n=4:sixteenorbitals: 4s 4p x, 4p y, 4 pz 4d xy, 4d yz, 4d z2, 4d xz, 4d x2-y2 4f x 7 70
Orbitals for the 1 st four energy levels of the H atom Zero point E = 0 n = Increasing E 4s 4p 4d 4f n = 4 3s 3p 3d n = 3 2s 2p n = 2 1s n = 1 Ground state E = (-) 71
Energies of the orbitals in n=1 to 4 for the H atom n = n = 4 n = 3 4s 3s 4p 4d 4f 3p 3d creasing E n = 2 2s 2p In n = 1 1s 72
Energies of the orbitals in n=1 to 4 for the H atom creasing E n = n = 4 n = 3 n = 2 Shells (energies) In 4s 3s 2s 4p 4d 4f 3p 3d subshells 2p Specific orbitals: defined orientations in space (orbital shapes) subshells (orbital shapes) n = 1 1s 73
Multielectron atoms All of the considerations to now were developed for the H atom with a single electron How does the quantum mechanical model apply to multielectron atoms? Same quantum numbers, shapes, orbitals Maximum of 2 electrons can occupy each orbital Energies of orbitals are affected by presence of other electrons in other orbitals 74
Principle energies are lower in multielectron atom Energies of subshells are split in multielectron atoms Some crossover of energy levels in multielectron atoms results in certain subshells having lower energies than subshells with lower n values Increasing E n = n = 4 n = 3 n =2 4s 3s 2s 4p 4d 4f 3p 3d 2p ΔE 1 2 ncreasing E I 6s 5s 4s 3s 5p 4p 3p 2p 4d 3d 4f 2s ΔE 1 2 n =1 1s 1s Energy levels of the single electron atom Energy levels of the multielectron atom, where additional electrons occupy higher energy orbitals in succession 75
Electrostatic interactions are responsible for the differences in energies for the single- and multielectron atoms Zp+ n = 1 n = 4 n = 3 n = 2 76
The spin quantum number, m s N S S N m s = +1/2 m s = -1/2 77
How do electrons populate orbitals in a multielectron atom? The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n energy shell l shape subshell m l orientation orbital m s electron spin clockwise (CW) or CCW 78
Populating orbitals in a multielectron atom 6s 5s 5p 4p 4d 4f 4s 3d Increas sing E 3s 2s 3p 2p The Pauli exclusion principle: p No two electrons in an atom can have the same set of four quantum numbers Electrons want to occupy the lowest energy orbital available 1s
4p Increa asing E 4s 3s 2s 3p 2p 3d The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n energy shell l shape subshell m l orientation orbital l m s electron spin clockwise (CW) or CCW Electrons want to occupy the lowest energy orbital available Represent an electron by a single sided arrow: and 1s 80
4p Increa asing E 4s 3s 2s 3p 2p n = 1 3d The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n energy shell l shape subshell m l orientation orbital m s electron spin clockwise (CW) or CCW Electrons want to occupy the lowest energy orbital available l = 0 The H atom: one e- 1s m l = 0 m s = +1/2 81
n value l value Subshell name m l value # of orbitals in subshell # of orbitals in shell Increa asing E 4p 3d 4s 3p 3s 2p 2s n = 1 n =1 l = 0 l = 0 m l = 0 m l = 0 m s = +1/2 m s = -1/2 1 0 1s 0 1 1 2 0 2s 0 1 4 2 1 2p -1, 0, 1 3 3 0 3s 0 1 9 3 1 3p -1, 0, 1 3 3 2 3d -2, -1, 0, 1, 2 5 4 0 4s 0 1 16 4 1 4p -1, 0, 1 3 4 2 4d -2, -1, 0, 1, 2 5 4 3 4f -3, -2, -1, 0, 1, 2, 3 7 The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n energy shell l shape subshell m l orientation orbital m s electron spin clockwise (CW) or CCW Electrons want to occupy the lowest energy orbital available 1s The He atom: two e- 82
4p Increa asing E 4s 3s 2s 3p 2p n = 2 l = 0 m l = 0 3d The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n energy shell l shape subshell m l orientation orbital m s electron spin clockwise (CW) or CCW 0 Electrons want to occupy the lowest energy orbital available m s = +1/2 1s n = 1 n = 1 l = 0 l = 0 m l = 0 m l = 0 m s = +1/2 m s = -1/2 The Li atom: 3 e- 83
4p Increa asing E 4s 3s 2s 3p 2p n = 2 l = 0 m l = 0 m s = +1/2 n = 2 l = 0 m l = 0 3d The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n energy shell l shape subshell m l orientation orbital m s electron spin clockwise (CW) or CCW 0 Electrons want to occupy the lowest energy orbital available m s = -1/2 1s n = 1 n = 1 l = 0 l = 0 m l = 0 m l = 0 m s = +1/2 m s = -1/2 The Be atom: 4 e- 84
4p Increa asing E 4s 3s 2s 3p 2p n = 2 l = 0 m l = 0 m s = +1/2 3d n = 2 l = 0 m l = 0 m s = -1/2 The Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers n = 2 l = 1 m l = 1 m s = +1/2 Electrons want to occupy the lowest energy orbital available 1s n = 1 n = 1 l = 0 l = 0 m l = 0 m l = 0 m s = +1/2 m s = -1/2 The B atom: 5 e- 85
Thegistofallthis: this: An electron will occupy the lowest energy orbital that is available A maximum of two electrons can occupy any given orbital (Pauli exclusion principle) 86
4p The B atom: 5 e- Increa asing E 4s 3s 3p 2p 3d An electron will occupy the lowest energy orbital that is available A maximum of two electrons can occupy any given orbital (Pauli exclusion principle) 2s 1s 87
Hund s rule For degenerate orbitals (orbitals of the same energy), the lowest energy is attained when the number of electrons with the same spin is maximized Two e- with different spins Two e- with same spins 88
Increa asing E 4s 3s 2s 4p The C atom: 6 e- 3d An electron will occupy the lowest 3p energy orbital that is available A maximum of two electrons can occupy any given orbital (Pauli exclusion principle) 2p For degenerate orbitals (orbitals of the same energy), the lowest energy is attained when the number of electrons with the same spin is maximized 1s 89
Increa asing E 4s 3s 2s 4p The N atom: 7 e- 3d An electron will occupy the lowest 3p energy orbital that is available A maximum of two electrons can occupy any given orbital (Pauli exclusion principle) 2p For degenerate orbitals (orbitals of the same energy), the lowest energy is attained when the number of electrons with the same spin is maximized 1s 90
4p The O atom: 8 e- 4s 3d Increa asing E 3s 2s 1s 3p An electron will occupy the lowest energy orbital that is available 2p A maximum of two electrons can occupy any given orbital (Pauli exclusion principle) For degenerate orbitals (orbitals of the same energy), the lowest energy is attained when the number of electrons with the same spin is maximized 91
4p The O atom: 8 e- Increa asing E 4s 3s 2s 3p 2p 3d A faster way to draw orbital filling: 1s 1s 2s 2p 3s 3p 4s 92
Electron configuration for the O atom 1s 2s 2p 3s 3p 4s Electron configuration: 1s 2 2s 2 2p 4 4p 4s 3d Incr reasing E 3s 2s 3p 2p 1s 93
Write electron configurations for n = 2 and n = 3 elements 4p Increasing E 4s 3s 2s 3p 2p 3d 1s 1s 2s 2p 3s 3p 4s 3d 4p 94
Write electron configurations for n = 4 elements 4p Increasing E 4s 3s 2s 3p 2p 3d 1s 1s 2s 2p 3s 3p 4s 3d 4p 95
Rules for adding electrons to orbitals Electrons go into the lowest energy orbital available For a given n shell, s orbitals are lower in energy than p which are lower in energy than d which are lower than f Only two electrons can occupy a given orbital For orbitals of the same energy, one electron will occupy each orbital before they start pairing up Some energy level crossovers occur due to electrostatic effects 96
Two easy ways to keep track of the order of orbital filling in multielectron atoms Use the periodic table to guide you The principle quantum number of d block elements is one less than that of the adjacent s and p block elements The principle quantum number of f block elements is two less than that of the adjacent s and p block elements, and one less than the adjacent d block Even easier: use the Auf-Bau principle 97
The periodic table color coded by the type of subshell being filled 1s 1s 2s 2p 3s 4s 5s 6s 7s 3d 4d 5d 6d 3p 4p 5p 6p 7p 4f 5f 98
The periodic table color coded by principle quantum number of the orbital being filled 1s2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 99
The periodic table color coded by principle quantum number of the orbital being filled 1s2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 100 100
The Auf-Bau Rule : The Order in which the Orbitals Fill in Polyelectronic Atoms 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 6f 14 7s 2 7p 6 Follow this rule and you can t go wrong!! 1s2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 101
Write electron configurations for n = 5 elements Increasing E 6s 5p 5s 4d 4p 3d 4s 3p 3s 2p 2s 4f 1s 2 2s 2 2p 6 1s 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 6f 14 7s 2 7p 6 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 102
Write electron configurations for n = 6 elements Increasing E 6s 5p 5s 4d 4p 3d 4s 3p 3s 2p 2s 4f 1s 2 2s 2 2p 6 1s 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 6f 14 7s 2 7p 6 1s2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 103
Electron configurations of the elements Anomolous electron configurations are in red 104
Electron configurations of the elements, color coded by subshell type 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 Anomolous electron configurations are in red 6s 2 6p 6 6d 10 6f 14 7s 2 7p 6 105