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    John E. McMurry Robert C. Fay

    Lecture NotesAlan D. EarhartSoutheast Community College Lincoln, NE

    General Chemistry: Atoms First

    Chapter 3Periodicity and the Electronic Structure ofAtoms

    Copyright 2010 Pearson Prentice Hall, Inc.

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    Chapter 3/2

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/3

    Light and the Electromagnetic

    SpectrumElectromagnetic energy (light) is characterized by

    wavelength, frequency, and amplitude.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/6

    Light and the Electromagnetic

    Spectrum

    Wavelength x Frequency = Speed

    =

    m

    s

    m

    s1

    cx

    cis defined to be the rate of travel of all

    electromagnetic energy in a vacuum

    and is a constant valuespeed of light.

    c= 3.00 x 108

    sm

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/7

    Light and the Electromagnetic

    Spectrum

    The light blue glow given off by mercury streetlamps

    has a wavelength of 436 nm. What is the frequency in

    hertz?

    436 nm

    3.00 x 108 sm

    1 x 109 nm

    1 m

    c=

    =

    = 6.88 x 1014 s-1 = 6.88 x 1014 Hz

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/8

    Electromagnetic Energy and

    Atomic Line Spectra

    Line Spectrum: A series of discrete lines on an

    otherwise dark background as a result of light emitted

    by an excited atom.

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    Chapter 3/9

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    Chapter 3/10

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/11

    1= R

    n2

    1

    m2

    1-

    Electromagnetic Energy and

    Atomic Line Spectra

    Johannes Rydberg later modified the equation to fit

    every line in the entire spectrum of hydrogen.

    Johann Balmer in 1885 discovered a mathematical

    relationship for the four visible lines in the atomic line

    spectra for hydrogen.

    R (Rydberg Constant) = 1.097 x 10-2 nm-1

    1

    = R n21

    22

    1-

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/12

    Particlelike Properties of

    Electromagnetic Energy

    Photoelectric Effect: Irradiation of clean metal

    surface with light causes electrons to be ejected from

    the metal. Furthermore, the frequency of the light used

    for the irradiation must be above some threshold

    value, which is different for every metal.

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    Chapter 3/13

    Particlelike Properties of

    Electromagnetic Energy

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/14

    Particlelike Properties of

    Electromagnetic Energy

    Einstein explained the effect by assuming that a beam

    of light behaves as if it were a stream of particlescalledphotons.

    Photoelectric Effect: Irradiation of clean metal

    surface with light causes electrons to be ejected from

    the metal. Furthermore, the frequency of the light used

    for the irradiation must be above some threshold

    value, which is different for every metal.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/15

    Particlelike Properties of

    Electromagnetic Energy

    E

    Quantum: The amount of energy corresponding

    to one photon of light.

    h (Plancks constant) = 6.626 x 10-34 J s

    Electromagnetic energy (light) is quantized.

    E= h

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/16

    Particlelike Properties of

    Electromagnetic Energy

    Niels Bohr proposed in 1914 a model of the hydrogen

    atom as a nucleus with an electron circling around it.

    In this model, the energy levels of the orbits are

    quantized so that only certain specific orbitscorresponding to certain specific energies for the

    electron are available.

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    Chapter 3/17

    Particlelike Properties of

    Electromagnetic Energy

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/18

    Wavelike Properties of Matter

    The de Broglie equation allows the calculation of a

    wavelength of an electron or of any particle or object

    of mass m and velocity v.

    mv

    h=

    Louis de Broglie in 1924 suggested that, iflightcan

    behave in some respects like matter, then perhaps

    mattercan behave in some respects like light.

    In other words, perhaps matter is wavelike as well asparticlelike.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/19

    Quantum Mechanics and the

    Heisenberg Uncertainty Principle

    In 1926 Erwin Schrdinger proposed the quantum

    mechanical model of the atom which focuses on the

    wavelike properties of the electron.

    In 1927 Werner Heisenberg stated that it is impossibleto know precisely where an electron is and what path

    it followsa statement called the Heisenberg

    uncertainty principle.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/20

    Wave Functions and Quantum

    NumbersProbability of finding

    electron in a region

    of space ( 2)

    Wave

    equation

    Wave function

    or orbital ( )

    solve

    A wave function is characterized by three parameterscalled quantum numbers, n, l, ml.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/21

    Principal Quantum Number (n)

    Describes the size and energy level of the orbital

    Commonly called shell

    Positive integer (n = 1, 2, 3, 4, )

    As the value ofn increases: The energy of the electron increases

    The average distance of the electron from the

    nucleus increases

    Wave Functions and Quantum

    Numbers

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/22

    Wave Functions and Quantum

    NumbersAngular-Momentum Quantum Number (l)

    Defines the three-dimensional shape of the orbital

    Commonly called subshell

    There are n different shapes for orbitals

    Ifn = 1 then l= 0 Ifn = 2 then l= 0 or 1

    Ifn = 3 then l= 0, 1, or2

    etc.

    Commonly referred to by letter (subshell notation)

    l= 0 s (sharp) l= 1 p (principal)

    l= 2 d(diffuse)

    l= 3 f(fundamental)

    etc.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/23

    Wave Functions and Quantum

    NumbersMagnetic Quantum Number (ml )

    Defines the spatial orientation of the orbital

    There are 2l+ 1 values ofmland they can have

    any integral value from -lto +l

    Ifl= 0 then ml= 0 Ifl= 1 then ml= -1, 0, or 1

    Ifl= 2 then ml= -2, -1, 0, 1, or2

    etc.

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    Wave Functions and Quantum

    Numbers

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    Chapter 3/25

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    Chapter 3/26

    The Shapes of Orbitals

    Node: A surface of zero

    probability for finding

    the electron.

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    Chapter 3/27

    The Shapes of Orbitals

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    The Shapes of Orbitals

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/31

    Quantum Mechanics and

    Atomic Line Spectra

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    Electron Spin and the Pauli

    Exclusion PrincipleElectrons have spin which gives rise to a tiny

    magnetic field and to a spin quantum number (ms).

    Pauli Exclusion Principle: No two electrons in an

    atom can have the same four quantum numbers.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/34

    Electron Configurations of

    Multielectron AtomsElectron Configuration: A description of which

    orbitals are occupied by electrons.

    Degenerate Orbitals: Orbitals that have the same

    energy level. For example, the threep orbitals in agiven subshell.

    Ground-State Electron Configuration: The lowest-

    energy configuration.

    Aufbau Principle (building up): A guide for

    determining the filling order of orbitals.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/35

    Electron Configurations of

    Multielectron AtomsRules of the aufbau principle:

    1. Lower-energy orbitals fill before higher-energy

    orbitals.

    2. An orbital can only hold two electrons, which musthave opposite spins (Pauli exclusion principle).

    3. If two or more degenerate orbitals are available,

    follow Hunds rule.

    Hunds Rule: If two or more orbitals with the same

    energy are available, one electron goes into each until

    all are half-full. The electrons in the half-filled orbitals

    all have the same value of their spin quantum number.

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/36

    Electron Configurations of

    Multielectron Atoms

    n = 1s orbital (l= 0)

    1 electronH: 1s1

    Electron

    Configuration

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/37

    1s2

    n = 1

    s orbital (l= 0)

    2 electrons

    Electron Configurations of

    Multielectron Atoms

    H:

    He:

    Electron

    Configuration

    1s1

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/38

    n = 2

    s orbital (l= 0)

    1 electrons1s2 2s1

    Electron Configurations of

    Multielectron Atoms

    H:

    Li:

    Lowest energy to highest energy

    He:

    Electron

    Configuration

    1s2

    1s1

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    Electron Configurations of

    Multielectron Atoms

    N:

    n = 2

    p orbital (l= 1)

    3 electrons

    H:

    1s2 2s2 2p3

    Li:

    He:

    Electron

    Configuration

    1s2 2s1

    1s2

    1s1

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/40

    Electron Configurations of

    Multielectron Atoms

    N:

    H:

    Li:

    1s

    He:

    Electron

    Configuration

    Orbital-Filling

    Diagram

    1s2 2s2 2p3

    1s2 2s1

    1s2

    1s1

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    Electron Configurations of

    Multielectron Atoms

    N:

    H:

    Li:

    1s

    1sHe:

    Electron

    Configuration

    Orbital-Filling

    Diagram

    1s2 2s2 2p3

    1s2 2s1

    1s2

    1s1

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/42

    Electron Configurations of

    Multielectron Atoms

    N:

    H:

    Li:

    1s

    1s

    2s1s

    He:

    Electron

    Configuration

    Orbital-Filling

    Diagram

    1s2 2s2 2p3

    1s2 2s1

    1s2

    1s1

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    Electron Configurations of

    Multielectron Atoms

    N:

    H:

    He:

    Li:

    Orbital-Filling

    Diagram

    1s

    1s

    2s1s

    1s 2p2s

    Electron

    Configuration

    1s2 2s2 2p3

    1s2 2s1

    1s2

    1s1

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/44

    Electron Configurations of

    Multielectron Atoms

    Na: [Ne] 3s11s2 2s2 2p6 3s1

    Ne configuration

    Electron

    Configuration

    Shorthand

    Configuration

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    Electron Configurations of

    Multielectron Atoms

    Na:

    P:

    [Ne] 3s11s2 2s2 2p6 3s1

    1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3

    Electron

    Configuration

    Shorthand

    Configuration

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    Electron Configurations of

    Multielectron Atoms

    Na:

    P:

    K:

    [Ne] 3s11s2 2s2 2p6 3s1

    1s2 2s2 2p6 3s2 3p3

    1s2 2s2 2p6 3s2 3p6 4s1

    [Ne] 3s2 3p3

    [Ar] 4s1

    Ar configuration

    Electron

    Configuration

    Shorthand

    Configuration

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/47

    Electron Configurations of

    Multielectron Atoms

    Sc:

    Na:

    P:

    1s2 2s2 2p6 3s2 3p6 4s2 3d1

    K:

    Shorthand

    Configuration

    Electron

    Configuration

    [Ar] 4s1 3d1

    [Ne] 3s11s2 2s2 2p6 3s1

    1s2 2s2 2p6 3s2 3p3

    1s2 2s2 2p6 3s2 3p6 4s1

    [Ne] 3s2 3p3

    [Ar] 4s1

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    Chapter 3/48

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    Copyright 2010 Pearson Prentice Hall, Inc. Chapter 3/49

    Some Anomalous Electron

    Configurations

    [Ar] 4s1 3d5Cr:

    Cu: [Ar] 4s1 3d10

    Actual

    Configuration

    Expected

    Configuration

    [Ar] 4s2 3d4

    [Ar] 4s2 3d9

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    Electron Configurations and

    the Periodic TableValence Shell: Outermost shell.

    Br: 4s2 4p5Cl: 3s

    2

    3p5

    Na: 3s1Li:

    2s

    1

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    Chapter 3/52

    Electron Configurations and

    Periodic Properties: Atomic Radii

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    Chapter 3/53

    Electron Configurations and

    Periodic Properties: Atomic Radii

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    Electron Configurations and

    Periodic Properties: Atomic Radii

    radiusrow

    radiuscolumn

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    Chapter 3/55


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