Solution Manual for Chemistry The Central Science 13th Edition by Theodore E. Brown H. Eugene LeMay Bruce E. Bursten Test bank

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Solution Manual for Chemistry The Central Science 13th Edition by Theodore E. Brown H. Eugene LeMay Bruce E. Bursten Test bank

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Chapter 6. Electronic Structure of Atoms Media Resources
Important Figures and Tables:
Figure 6.3 Electromagnetic Waves
Table 6.1 Common Wavelength Units for
Electromagnetic Radiation
Figure 6.4 The Electromagnetic Spectrum
Figure 6.7 The Photoelectric Effect
Figure 6.9 Creating a Spectrum
Figure 6.11 Line Spectra of Hydrogen and Neon Figure 6.12 Energy Levels in the Hydrogen Atom
from the Bohr Model
Figure 6.17 Electron-Density Distribution
Table 6.2 Relationship Among Values of n, l, and
ml throughn=4
Figure 6.18 Energy Levels in the Hydrogen Atom Figure 6.19 Radial Probability Functions for the
1s, 2s, and 3s Orbitals of Hydrogen
Figure 6.22 Probability Density [(r)]2 in the 1s,
2s and 3s Orbitals of Hydrogen
Figure 6.23 The p Orbitals
Figure 6.24 Contour Representations of the Five
d Orbitals
Figure 6.25 General Energy Ordering of Orbitals
for a Many-Electron Atom
Table 6.3 Electron Configurations of Several
Lighter Elements
Figure 6.30 Regions of the Periodic Table Figure 6.31 Outer-Shell Electron Configurations
of the Elements
Activities:
Electromagnetic Spectrum Bohr Model
Quantum Numbers Electron Configuration Periodic Table
Animations:
Photoelectric Effect
Radial Electron Distribution Electron Configuration
Movies:
Flame Tests for Metals
Section:
6.1 The Wave Nature of Light 6.1 The Wave Nature of Light
6.1 The Wave Nature of Light
6.2 Quantized Energy and Photons 6.3 Line Spectra and the Bohr Model 6.3 Line Spectra and the Bohr Model 6.3 Line Spectra and the Bohr Model
6.5 Quantum Mechanics and Atomic Orbitals 6.5 Quantum Mechanics and Atomic Orbitals
6.5 Quantum Mechanics and Atomic Orbitals 6.6 Representations of Orbitals
6.6 Representations of Orbitals
6.6 Representations of Orbitals 6.6 Representations of Orbitals
6.7 Many-Electron Atoms
6.8 Electron Configurations
6.9 Electron Configurations and the Periodic Table 6.9 Electron Configurations and the Periodic Table
Section:
6.1 The Wave Nature of Light
6.3 Line Spectra and the Bohr Model
6.5 Quantum Mechanics and Atomic Orbitals
6.8 Electron Configurations
6.9 Electron Configurations and the Periodic Table
Section:
6.2 Quantized Energy and Photons 6.6 Representations of Orbitals 6.8 Electron Configurations
Section:
6.3 Line Spectra and the Bohr Model
Copyright 2015 Pearson Education, Inc.
VCL Simulations:
Blackbody Radiation Photoelectric Effect
The Rydberg Equation
Atomic Emission Spectra Heisenberg Uncertainty Principle
Other Resources
Further Readings:
Scientific American, September 2004
Put Body to Them!
Presenting the Bohr Atom
Getting the Numbers RightThe Lonely Struggle
of Rydberg
Suitable Light Sources and Spectroscopes for
Student Observation of Emission Spectra in
Lecture Halls Niels Bohr
100 Years of Quantum Mysteries
On a Relation Between the Heisenberg and
de Broglie Principles
Introducing the Uncertainty Principle Using
Diffraction of Light Waves
Perspectives on the Uncertainty Principle and
Quantum Reality
A Students Travels, Close Dancing, Bathtubs,
and the Shopping Mall: More Analogies in
Teaching Introductory Chemistry
The Mole, the Periodic Table, and Quantum
Numbers: An Introductory Trio
The Origin of the s, p, d, f Orbital Labels Electron Densities: Pictorial Analogies for
Apparent Ambiguities in Probability
Calculations
Magnetic Whispers: Chemistry and Medicine
Finally Tune Into Controversial Molecular
Chatter Seeing Inside
The Nobel Prize in Medicine for Magnetic Resonance Imaging
Chemistry in Britain, June 1996
The Magnetic Eye
Mind Over Matter
New Schemes for Applying the Aufbau Principle A Low-Cost Classroom Demonstration of the
Aufbau Principle
Demystifying Introductory Chemistry Part 1:
Electron Configurations from Experiment Quantum Analogies on Campus
Housing Electrons: Relating Quantum Numbers,
Section:
6.2 Quantized Energy and Photons 6.2 Quantized Energy and Photons 6.3 Line Spectra and the Bohr Model 6.3 Line Spectra and the Bohr Model 6.4 The Wave Behavior of Matter
Section:
6.2 Quantized Energy and Photons 6.2 Quantized Energy and Photons 6.3 Line Spectra and the Bohr Model 6.3 Line Spectra and the Bohr Model
6.3 Line Spectra and the Bohr Model
6.3 Line Spectra and the Bohr Model 6.3 Line Spectra and the Bohr Model 6.4 The Wave Behavior of Matter
6.4 The Wave Behavior of Matter
6.4 The Wave Behavior of Matter
6.5 Quantum Mechanics and Atomic Orbitals
6.5 Quantum Mechanics and Atomic Orbitals
6.6 Representations of Orbitals 6.6 Representations of Orbitals
6.7 Many-Electron Atoms
6.7 Many-Electron Atoms 6.7 Many-Electron Atoms
6.7 Many-Electron Atoms 6.7 Many-Electron Atoms 6.7 Many-Electron Atoms 6.7 Many-Electron Atoms 6.7 Many-Electron Atoms
6.8 Electron Configurations
6.8 Electron Configurations 6.8 Electron Configurations
Energy Levels, and Electron Configurations
Copyright 2015 Pearson Education, Inc.
Electronic Structure of Atoms 75
76 Chapter 6
Pictorial Analogies VII: Quantum Numbers and Orbitals
The Quantum Shoe Store and Electron Structure Some Analogies for Teaching Atomic Structure
at the High School Level
Ionization Energies, Parallel Spins, and the Stability
of Half-Filled Shells
The Noble Gas ConfigurationNot the Driving
Force but the Rule of the Game in Chemistry The Periodic Table as a Mnemonic Device for
Writing Electronic Configurations
The Periodic Table and Electron Configurations
Live Demonstrations:
Simple and Inexpensive Classroom Demonstration of Nuclear Magnetic Resonance and Magnetic Resonance Imaging
6.8 Electron Configurations
6.8 Electron Configurations 6.8 Electron Configurations
6.8 Electron Configurations
6.8 Electron Configurations
6.9 Electron Configurations and the Periodic Table
6.9 Electron Configurations and the Periodic Table
Section:
6.7 Many-Electron Atoms
Copyright 2015 Pearson Education, Inc.
Chapter 6. Electronic Structure of Atoms Common Student Misconceptions
Some students have difficulties converting between angstroms, nanometers, etc. and meters.
Students often have difficulties switching from the language of certainties to the language of
probabilities.
Students are often frightened or put off by the mathematics, vocabulary, foreign names and an
apparent intangibility of the information.
Students are initially unaware that the quantum theory laid foundations for such areas as spectroscopy
and nanotechnology.
Students confuse Bohrs orbits with orbitals; most spellcheckers do not recognize the word orbital.
Students mistakenly think that spectral lines represent energy levels; consequently
Students have difficulties associating a given line in an emission (or absorption) spectrum with a
transition between two energy levels.
When drawing the orbital diagrams, students often draw 2, 6, 10, and 14 boxes for s, p, d, and f
orbitals, respectively.
Teaching Tips
This is often students first glimpse at the realm of quantum theory. They need to understand that the model has been built up to rationalize experimental data. They also need to know that elements of one theory are maintained in the subsequent theory.
Students may not be familiar with the common symbol for wavelength (the lowercase lambda, ) and the common symbol for frequency (the lowercase nu, ).
Using the unit s1 for frequency makes the units cancel more easily.
Many students will be familiar with the mnemonic for remembering the order of colors in the visible
spectrum: ROY G BIV, which stands for red, orange, yellow, green, blue, indigo, and violet.
A good analogy for the uncertainty principle: picture a busy intersection photographed at night. With
a short exposure, you get a clear image of the position of every car, but you cannot tell how fast they are going or whether they are going forward or backward or about to swerve or turn. With a time- lapsed exposure, you can tell from the streaks of light the speed and direction of each car, but you cannot tell where each one currently is. You can know position or path, but not both.
Lecture Outline
6.1 The Wave Nature of Light1,2,3,4
The electronic structure of an atom refers to the arrangement of electrons.
Visible light is a form of electromagnetic radiation or radiant energy.
Radiation carries energy through space.
Electromagnetic radiation is characterized by its wave nature.
All waves have a characteristic wavelength, lambda), and amplitude, A.
The frequency, nu), of a wave is the number of cycles which pass a point in one second.
The units of are hertz (1 Hz = 1 s1).
1 Figure 6.3
2 Figure 6.4
3 Electromagnetic Spectrum Activity from Instructors Resource CD/DVD
4 Table 6.1
Electronic Structure of Atoms 77
Copyright 2015 Pearson Education, Inc.
78 Chapter 6
The speed of a wave is given by its frequency multiplied by its wavelength.
For light, speed, c = ,
Electromagnetic radiation moves through a vacuum with a speed of 3.00 108 m/s.
Electromagnetic waves have characteristic wavelengths and frequencies.
The electromagnetic spectrum is a display of the various types of electromagnetic radiation arranged
in order of increasing wavelength.
Example: visible radiation has wavelengths between 400 nm (violet) and 750 nm (red).
FORWARD REFERENCES
X-ray diffraction will be discussed in Chapter 12.
Light emitting diodes will be described in Chapter 11 (section 11.7).
Different ranges of the electromagnetic spectrum will be mentioned in Chapters 18 and 23.
Gamma radiation will be further discussed on Chapter 21.
6.2 Quantized Energy and Photons
Some phenomena cant be explained using a wave model of light:
Blackbody radiation is the emission of light from hot objects.
The photoelectric effect is the emission of electrons from metal surfaces on which light shines.
Emission spectra are the emissions of light from electronically excited gas atoms.
Hot Objects and the Quantization of Energy5
Heated solids emit radiation (blackbody radiation)
The wavelength distribution depends on the temperature (i.e., red hot objects are cooler than
white hot objects).
Planck investigated blackbody radiation.
He proposed that energy can only be absorbed or released from atoms in certain amounts.
These amounts are called quanta.
A quantum is the smallest amount of energy that can be emitted or absorbed as electromagnetic
radiation.
The relationship between energy and frequency is:
E = h
where h is Plancks constant (6.626 1034 Js).
To understand quantization consider the notes produced by a violin (continuous) and a piano
(quantized):
A violin can produce any note when the fingers are placed at an appropriate spot on the bridge.
A piano can only produce notes corresponding to the keys on the keyboard.
The Photoelectric Effect and Photons6,7,8,9,10
The photoelectric effect provides evidence for the particle nature of light. It also provides evidence for quantization.
Einstein assumed that light traveled in energy packets called photons. The energy of one photon is E = h.
Light shining on the surface of a metal substance can cause electrons to be ejected from the metal. The electrons will only be ejected if the photons have sufficient energy (work function):
5 Blackbody Radiation VCL Simulation from Instructors Resource CD/DVD
6 September 2004 issue of Scientific American from Further Readings
7 Photoelectric Effect VCL Simulation from Instructors Resource CD/DVD
8 Photoelectric Effect Animation from Instructors Resource CD/DVD
9 Figure 6.7
10 Put Body to Them! from Further Readings
Copyright 2015 Pearson Education, Inc.
Below the threshold frequency no electrons are ejected.
Above the threshold frequency, the excess energy appears as the kinetic energy of the ejected
electrons.
Light has wave-like AND particle-like properties. FORWARD REFERENCES
Photoconductivity in solar energy conversions and emission of photons by semiconductor nanoparticles will be described in Chapter 12 (section 12.9).
Photodissociation, i.e., bond breaking as a result of an absorption of a photon by a molecule, as well as photodecomposition will be discussed in Chapter 18 (section 18.2).
The role of photons from sunlight in photosynthesis will be further discussed in Chapter 23 (section 23.3).
6.3 Line Spectra and the Bohr Model Line Spectra11,12,13,14,15
Radiation composed of only one wavelength is called monochromatic.
Radiation that spans a whole array of different wavelengths is called continuous.
When radiation from a light source, such as a light bulb, is separated into its different wavelength
components, a spectrum is produced.
White light can be separated into a continuous spectrum of colors.
A rainbow is a continuous spectrum of light produced by the dispersal of sunlight by raindrops or mist.
On the continuous spectrum there are no dark spots, which would correspond to different lines. Not all radiation is continuous.
A gas placed in a partially evacuated tube and subjected to a high voltage produces single colors of light.
The spectrum that we see contains radiation of only specific wavelengths; this is called a line spectrum.
Bohrs Model16,17
Rutherford assumed that electrons orbited the nucleus analogous to planets orbiting the sun.
However, a charged particle moving in a circular path should lose energy.
This means that the atom should be unstable according to Rutherfords theory.
Bohr noted the line spectra of certain elements and assumed that electrons were confined to specific energy states. These were called orbits.
Bohrs model is based on three postulates:
Only orbits of specific radii, corresponding to certain definite energies, are permitted for electrons
in an atom.
An electron in a permitted orbit has a specific energy and is an allowed energy state.
Energy is only emitted or absorbed by an electron as it moves from one allowed energy state to
another.
The energy is gained or lost as a photon.
11 Figure 6.9
12 Flame Tests for Metals Movie from Instructors Resource CD/DVD
13 The Rydberg Equation VCL Simulation from Instructors Resource CD/DVD
14 Atomic Emission Spectra VCL Simulation from Instructors Resource CD/DVD 15 Figure 6.11
16 Presenting the Bohr Atom from Further Readings
17 Bohr Model Activity from Instructors Resource CD/DVD
Electronic Structure of Atoms 79
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80 Chapter 6
The Energy States of the Hydrogen Atom18,19,20
Colors from excited gases arise because electrons move between energy states in the atom.
Since the energy states are quantized, the light emitted from excited atoms must be quantized and
appear as line spectra.
Bohr showed mathematically that
1 18 1 E(hcRH) 2(2.1810 J) 2 n n
where n is the principal quantum number (i.e., n = 1, 2, 3, ) and RH is the Rydberg constant.
The product hcRH = 2.18 10-18 J.
The first orbit in the Bohr model has n = 1 and is closest to the nucleus.
The furthest orbit in the Bohr model has n = and corresponds to E = 0.
Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in
quanta (E = h.
The ground state is the lowest energy state.
An electron in a higher energy state is said to be in an excited state.
The amount of energy absorbed or emitted by moving between states is given by
Limitations of the Bohr Model21,22
The Bohr Model has several limitations:
It cannot explain the spectra of atoms other than hydrogen.
Electrons do not move about the nucleus in circular orbits.
However, the model introduces two important ideas:
The energy of an electron is quantized: electrons exist only in certain energy levels described by
quantum numbers.
Energy gain or loss is involved in moving an electron from one energy level to another.
FORWARD REFERENCES
Absorption of sufficient amounts of energy to ionize an atom will be further discussed in Chapter 7 (section 7.4).
Emission of light with characteristic colors (flame test) by excited atoms of Li, Na and K is shown in Chapter 7 (section 7.7).
Selective absorption of light by chemicals, e.g., organic dyes, will be described in Chapter 9.
Absorption of wavelengths from the visible part of the electromagnetic spectrum by
molecules of chlorophyll and other pigments will be discussed in Chapter 23 (section 23.3).
Absorption in the visible range will be responsible for colors of many transition metal
complexes (Chapter 23, section 23.5).
18 Getting the Numbers RightThe Lonely Struggle of Rydberg from Further Readings
19 Figure 6.12
20 Suitable Light Sources and Spectroscopes for Student Observation of Emission Spectra in Lecture
Halls from Further Readings
21 Niels Bohr from Further Readings
22 100 Years of Quantum Mysteries from Further Readings
Copyright 2015 Pearson Education, Inc.
6.4 The Wave Behavior of Matter23
Knowing that light has a particle nature, it seems reasonable to ask whether matter has a wave nature.
This question was answered by Louis de Broglie.
Using Einsteins and Plancks equations, de Broglie derived:
h/m
The momentum, m is a particle property, whereas is a wave property.
Matter waves are the term used to describe wave characteristics of material particles.
Therefore, in one equation de Broglie summarized the concepts of waves and particles as they
apply to low-mass, high-speed objects.
As a consequence of de Broglies discovery, we now have techniques such as X-ray diffraction
and electron microscopy to study small objects.
The Uncertainty Principle24,25,26
Heisenbergs Uncertainty Principle: we cannot determine the exact position, direction of motion, and speed of subatomic particles simultaneously.
For electrons: we cannot determine their momentum and position simultaneously.
FORWARD REFERENCES
X-ray diffraction and X-ray crystallography will be further discussed in Chapter 12 (section 12.2).
6.5 Quantum Mechanics and Atomic Orbitals27,28
Schrodinger proposed an equation containing both wave and particle terms.
Solving the equation leads to wave functions, .
The wave function describes the electrons matter wave.
The square of the wave function, , gives the probability of finding the electron.
That is, gives the electron density for the atom.
is called the probability density.
Electron density is another way of expressing probability.
A region of high electron density is one where there is a high probability of finding an electron.
Orbitals and Quantum Numbers29,30,31,32
If we solve the Schrodinger equation we get wave functions and energies for the wave functions.
We call orbitals.
Schrodingers equation requires three quantum numbers:
Principal quantum number, n. This is the same as Bohrs n.
As n becomes larger, the atom becomes larger and the electron is further from the nucleus.
23 On a Relation Between the Heisenberg and de Broglie Principles from Further Readings
24 Heisenberg Uncertainty Principle VCL Simulation from Instructors Resource CD/DVD
25 Introducing the Uncertainty Principle Using Diffraction of Light Waves from Further Readings
26 Perspectives on the Uncertainty Principle and Quantum Reality from Further Readings
27 Figure 6.17
28 A Students Travels, Close Dancing, Bathtubs, and the Shopping Mall: More Analogies in Teaching
Introductory Chemistry from Further Readings
29 The Mole, the Periodic Table and Quantum Numbers: An Introductory Trio from Further Readings 30 Quantum Numbers Activity from Instructors Resource CD/DVD
31 Table 6.2
32 Figure 6.18
Electronic Structure of Atoms 81
Copyright 2015 Pearson Education, Inc.
82 Chapter 6
Angular momentum quantum number, l. This quantum number depends on the value of n.
The values of l begin at 0 and increase to n 1.
We usually use letters for l (s, p, d, and f for l = 0, 1, 2, and 3). Usually we refer to the s, p, d, and
f orbitals.
This quantum number defines the shape of the orbital.
Magnetic quantum number, ml.
This quantum number depends on l.
The magnetic quantum number has integer values between l and +l.
There are (2l+1) possible values of ml.
For example, for l = 1, there are (21+1) = 3 values of ml : 0, +1, and -1.
Consequently, for l = 1, there are 3 orbitals: px, py and pz.
Magnetic quantum numbers give the three-dimensional orientation of each orbital.
A collection of orbitals with the same value of n is called an electron shell. There are n2 orbitals in a shell described by a the n value.
For example, for n = 3, there are 32 = 9 orbitals.
A set of orbitals with the same n and l is called a subshell.
Each subshell is designated by a number and a letter. For example, 3p orbitals have n = 3 and l = 1.
There are n types of subshells in a shell described by a the n value. For example, for n = 3, there are 3 subshells: 3s, 3p and 3d.
Orbitals can be ranked in terms of energy to yield an Aufbau diagram. Note that this Aufbau diagram is for a single electron system.
As n increases, note that the spacing between energy levels becomes smaller. 6.6 Representations of Orbitals33,34
The s Orbitals35,36,37
All s orbitals are spherical.
As n increases, the s orbitals get larger.
As n increases, the number of nodes increases.
A node is a region in space where the probability of finding an electron is zero.
= 0 at a node.
For an s orbital the number of nodes is given by n 1.
We can plot a curve of radial probability density vs. distance (r) from the nucleus.
This curve is the radial probability function for the orbital. The p Orbitals38
There are three p orbitals: px, py and pz.
The three p orbitals lie along the x-, y-, and z-axes of a Cartesian system.
The letters correspond to allowed values of ml of 1, 0, and +1.
The orbitals are dumbbell shaped; each has two lobes.
As n increases, the p orbitals get larger.
All p orbitals have a node at the nucleus.
33 The Origin of the s, p, d, f Orbital Labels from Further Readings
34 Electron Densities: Pictorial Analogies for Apparent Ambiguities in Probability Calculations from
Further Readings
35 Radial Electron Distribution Animation from Instructors Resource CD/DVD 36 Figure 6.19
37 Figure 6.22
38 Figure 6.23
Copyright 2015 Pearson Education, Inc.
The d and f Orbitals39
There are five d and seven f orbitals.
Three of the d orbitals lie in a plane bisecting the x-, y-, and z-axes.
Two of the d orbitals lie in a plane aligned along the x-, y-, and z-axes.
Four of the d orbitals have four lobes each.
One d orbital has two lobes and a collar.
FORWARD REFERENCES
An overlap of atomic orbitals will be introduced in Chapter 9 (section 9.4).
Hybridization of atomic orbitals will be discussed in Chapter 9 (section 9.5).
Molecular orbitals will be introduced in Chapter 9 (section 9.7).
Overlap of p orbitals on C atoms will be implicated in the formation of bonds in organic
chemistry, as mentioned in Chapters 9, 22 (section 22.1; C vs. Si), and discussed in detail in
Chapter 24.
Energies of d orbitals in different crystal fields will be discussed in Chapter 23 (section 23.6).
6.7 Many-Electron Atoms40,41,42,43,44,45 Orbitals and Their Energies46,47,48
In a many-electron atom, for a given value of n,
The energy of an orbital increases with increasing value of l.
Orbitals of the same energy are said to be degenerate. Electron Spin and the Pauli Exclusion Principle49
Line spectra of many-electron atoms show each line as a closely spaced pair of lines.
Stern and Gerlach designed an experiment to determine why.
A beam of atoms was passed through a slit and into a magnetic field and the atoms were then detected.
Two spots were found: one with the electrons spinning in one direction and one with the electrons spinning in the opposite direction.
Since electron spin (electron as a tiny sphere spinning on its own axis) is quantized,
We define ms = spin magnetic quantum number = 1/2.
Paulis exclusion principle states that no two electrons can have the same set of 4 quantum numbers. Therefore, two electrons in the same orbital must have opposite spins.
FORWARD REFERENCES
The roles of screening and penetration in determining the relative energies of subshells within a shell will be explained in Chapter 7.
39 Figure 6.24
40 Magnetic Whispers: Chemistry and Medicine Finally Tune Into Controversial Molecular Chatter
from Further Readings
41 Seeing Inside from Further Readings
42 Simple and Inexpensive Classroom Demonstration of Nuclear Magnetic Resonance and Magnetic
Resonance Imaging from Live Demonstrations
43 The Nobel Prize in Medicine for Magnetic Resonance Imaging from Further Readings 44 June 1996 issue of Chemistry in Britain from Further Readings
45 The Magnetic Eye from Further Readings
46 Figure 6.25
47 New Schemes for Applying the Aufbau Principle from Further Readings
48 A Low-Cost Classroom Demonstration of the Aufbau Principle from Further Readings 49 Mind Over Matter from Further Readings
Electronic Structure of Atoms 83
Copyright 2015 Pearson Education, Inc.
84 Chapter 6
Paulis exclusion principle will also apply to hybrid orbitals in Chapter 9 (sections 9.49.5) and molecular orbitals (sections 9.79.8).
6.8 Electron Configurations
Electron configurations tell us how the electrons are distributed among the various orbitals of an atom.
The most stable configuration, or ground state, is that in which the electrons are in the lowest possible energy state.
We represent the configuration with an orbital diagram.
Each orbital is denoted with a box and each electron with a half arrow.
When writing ground-state electronic configurations:
electrons fill orbitals in order of increasing energy with no more than two electrons per orbital.
no two electrons can fill one orbital with the same spin (Pauli).
for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron.
How do we show spin?
An arrow pointing upwards has ms = + 1/2 (spin up).
An arrow pointing downwards has ms = 1/2 (spin down).
Electrons having opposite spins are paired when in the same orbital.
An unpaired electron is one not accompanied by a partner of opposite spin.
Hunds Rule50,51,52,53,54,55,56,57,58
Hundsrule:fordegenerateorbitals,thelowestenergyisattainedwhenthenumberofelectronswith the same spin is maximized.
Thus, electrons fill each orbital singly with their spins parallel before any orbital gets a second
electron.
By placing electrons in different orbitals, electron-electron repulsions are minimized.
Condensed Electron Configurations59
Electron configurations may be written using a shorthand notation (condensed electron configuration):
Write the valence electrons explicitly.
Valence electrons are electrons in the outer shell.
These electrons are gained and lost in reactions.
Write the core electrons corresponding to the filled noble gas in square brackets.
Core electrons are electrons in the inner shells.
50 Demystifying Introductory Chemistry; Part 1. Electron Configurations from Experiment from Further Readings
51 Quantum Analogies on Campus from Further Readings
52 Housing Electrons: Relating Quantum Numbers, Energy Levels, and Electron Configurations from
Further Readings
53 Pictorial Analogies VII: Quantum Numbers and Orbitals from Further Readings
54 The Quantum Shoe Store and Electron Structure from Further Readings
55 Table 6.3
56Electron Configuration Activity from Instructors Resource CD/DVD
57 Some Analogies for Teaching Atomic Structure at the High School Level from Further Readings
58 Ionization Energies, Parallel Spins, and the Stability of Half-Filled Shells from Further Readings 59 The Noble Gas ConfigurationNot the Driving Force but the Rule of the Game in Chemistry from
Further Readings
Copyright 2015 Pearson Education, Inc.
These are generally not involved in bonding. Example:
P is 1s22s22p63s23p3,
but Ne is 1s22s22p6.
Therefore, P is [Ne]3s23p3.
Transition Metals
After Ar the d orbitals begin to fill.
After the 3d orbitals are full, the 4p orbitals begin to fill.
The ten elements between Ti and Zn are called the transition metals, or transition elements.
The Lanthanides and Actinides
The 4f orbitals begin to fill with Ce.
Note: The electron configuration of La is [Xe]6s25d1.
The 4f orbitals are filled for the elements Ce Lu which are called lanthanide elements (or rare earth elements).
The 5f orbitals are filled for the elements Th Lr which are called actinide elements.
The actinide elements are radioactive and most are not found in nature.
FORWARD REFERENCES
Periodic properties associated with electron configurations, such as atomic radii, ionization energies and electron affinities, will be discussed throughout Chapter 7.
Valence electrons and the Octet Rule will be discussed in Chapter 8.
Valence electrons of atoms within molecules and ions will be added and distributed according
to the VSEPR model in Chapter 9 to determine molecular shapes.
Electron configurations and the associated chemical properties of nonmetals in groups 4A
8A will be discussed in detail in Chapter 22.
Electron configurations and the associated properties of select transition metals will be
discussed in Chapter 23.
High- and low-spin transition metal complexes will be discussed in Chapter 23 (section 23.6).
Electron configuration of the C atom will be highlighted in Chapter 24 on organic chemistry.
6.9 Electron Configurations and the Periodic Table60,61,62,63,64
The periodic table can be used as a guide for electron configurations.
The period number is the value of n.
Groups 1A and 2A have their s orbitals being filled.
Groups 3A8A have their p orbitals being filled.
The s-block and p-block of the periodic table contain the representative, or main-group, elements.
Groups 3B2B have their d orbitals being filled.
The lanthanides and actinides have their f orbitals being filled.
The actinides and lanthanide elements are collectively referred to as the f-block metals.
Note that the 3d orbitals fill after the 4s orbital. Similarly, the 4f orbitals fill after the 5d orbitals.
In general, for representative elements we do not consider the electrons in completely filled d or f
subshells to be valence electrons.
60 Figure 6.30
61 The Periodic Table as a Mnemonic Device for Writing Electronic Configurations from Further
Readings
62 The Periodic Table and Electron Configurations from Further Readings 63 Figure 6.31
64 Periodic Table Activity from Instructors Resource CD/DVD
Electronic Structure of Atoms 85
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86 Chapter 6
In general, for transition elements we do not consider the electrons in a completely filled f subshell to be valence electrons.
Anomalous Electron Configurations
There are many elements that appear to violate the electron configuration guidelines.
Examples:
Chromium is [Ar]3d54s1 instead of [Ar]3d44s2.
Copper is [Ar]3d104s1 instead of [Ar]3d94s2.
Half-full (d5) and full (d10) d subshells are particularly stable.
FORWARD REFERENCES
Electron configurations of ions of the main group elements will be covered in Chapter 8.
Electron configurations of transition metal cations will be mentioned in Chapter 8 and further
used in Chapter 23.
Copyright 2015 Pearson Education, Inc.
Further Readings:
1. The September 2004 issue of Scientific American is a special issue with numerous articles dealing with how Einsteins ideas reshaped the world.
2. Robert R. Perkins, Put Body to Them!, J. Chem. Educ., Vol. 72, 1995, 151152. This reference includes analogies for quantized states.
3. Bianca L. Haendler, Presenting the Bohr Atom, J. Chem. Educ., Vol. 59, 1982, 372376. Presenting the role of the Bohr theory within the framework of the development of quantum mechanics.
4. Mike Sutton, Getting the Numbers RightThe Lonely Struggle of Rydberg, Chemistry World, July 2004, 3841.
5. Elvin Hughes, Jr., and Arnold George, Suitable Light Sources and Spectroscopes for Student Observation of Emission Spectra in Lecture Halls, J. Chem. Educ., Vol. 61, 1984, 908909.
6. Dennis R. Sievers, Niels Bohr, J. Chem. Educ., Vol. 59, 1982, 303304. A short biography of Niels Bohr.
7. Max Tegmark and John Archibald Wheeler, 100 Years of Quantum Mysteries, Scientific American, February 2001, 6875.
8. Oliver G. Ludwig, On a Relation Between the Heisenberg and de Broglie Principles, J. Chem. Educ., Vol. 70, 1993, 28.
9. Pedro L. Muino, Introducing the Uncertainty Principle Using Diffraction of Light Waves, J. Chem. Educ., Vol. 77, 2000, 10251027.
10. Lawrence S. Bartell, Perspectives on the Uncertainty Principle and Quantum Reality, J. Chem. Educ., Vol. 62, 1985, 192196. This reading contains practical applications of the uncertainty principle.
11. Goeff Rayner-Canham, A Students Travels, Close Dancing, Bathtubs, and the Shopping Mall: More Analogies in Teaching Introductory Chemistry, J. Chem. Educ., Vol. 71, 1994, 943944. This reference includes an analogy dealing with the probability model of the atom.
12. Mali Yin and Raymond S. Ochs, The Mole, the Periodic Table, and Quantum Numbers: An Introductory Trio, J. Chem. Educ., Vol. 78, 2001, 13451347.
13. William B. Jensen, The Origin of the s, p, d, f Orbital Labels, J. Chem. Educ., Vol. 84, 2007, 757 758.
14. Maria Gabriela Lagorio, Electron Densities: Pictorial Analogies for Apparent Ambiguities in Probability Calculations, J. Chem. Educ., Vol. 77, 2000, 14441445.
15. Robert D. Freeman, New Schemes for Applying the Aufbau Principle, J. Chem. Educ., Vol. 67, 1990, 576.
16. James R. Hanley, III, and James R. Hanley, Jr., A Low-Cost Classroom Demonstration of the Aufbau Principle, J. Chem. Educ., Vol. 56, 1979, 747.
17. Robin Hendry, Mind Over Matter, Chemistry in Britain, November 2000, 3537. Copyright 2015 Pearson Education, Inc.
Electronic Structure of Atoms 87
88 Chapter 6
18. Peter Weiss, Magnetic Whispers: Chemistry and Medicine Finally Tune Into Controversial
Molecular Chatter, Science News, Vol. 159, 2001, 4244.
19. Mark Fischetti, Seeing Inside, Scientific American, August 2004, 9293. A short article on medical
imaging devices.
20. Charles G. Fry, The Nobel Prize in Medicine for Magnetic Resonance Imaging, J. Chem. Educ.,
Vol. 81, 2004, 922923.
21. Chemistry in Britain, Vol. 32, June, 1996. This issue contains several articles on the uses of NMR
its developments and its uses in medicine.
22. Lyn Gladden, The Magnetic Eye, Chemistry in Britain, November 2000, 3537. A short article about the work of Wolfgang Pauli.
23. Ronald J. Gillespie, James N. Spencer and Richard S. Moog, Demystifying Introductory Chemistry; Part 1. Electron Configurations from Experiment, J. Chem. Educ., Vol. 73, 1996, 617622. The use of experimental data to investigate electron configurations is presented in this reference.
24. Ngai Ling Ma, Quantum Analogies on Campus, J. Chem. Educ., Vol. 73, 1996, 10161017.
25. Anthony Garofalo, Housing Electrons: Relating Quantum Numbers, Energy Levels, and Electron
Configurations, J. Chem. Educ., Vol. 74, 1997, 709719.
26. John J. Fortman, Pictorial Analogies VII: Quantum Numbers and Orbitals, J. Chem. Educ., Vol. 70,
1993, 649650.
27. M. Bonneau, The Quantum Shoe Store and Electron Structure, J. Chem. Educ., Vol. 68, 1991, 837.
28. Ngoh Khang Goh, Lian Sai Chia, and Daniel Tan, Some Analogies for Teaching Atomic Structure at the High School Level, J. Chem. Educ., Vol. 71, 1994, 733734. Analogies for orbitals, Hunds Rule, and the four quantum numbers are included in this reference.
29. Peter Cann, Ionization Energies, Parallel Spins, and the Stability of Half-Filled Shells, J. Chem. Educ., Vol. 77, 2000, 10561061.
30. Roland Schmid, The Noble Gas ConfigurationNot the Driving Force but the Rule of the Game in Chemistry, J. Chem. Educ., Vol. 80, 2003, 931937.
31. Suzanne T. Mabrouk, The Periodic Table as a Mnemonic Device for Writing Electronic Configurations, J. Chem. Educ., Vol. 80, 2003, 894898.
32. Judith A. Strong, The Periodic Table and Electron Configurations, J. Chem. Educ., Vol. 63, 1986, 834.
Live Demonstrations
1. Joel A. Olson, Karen J. Nordell, Marla A. Chesnik, Clark R. Landis, Arthur B. Ellis, M.S. Rzchowski, S. Michael Condren, George C. Lisensky, and James W. Long, Simple and Inexpensive Classroom Demonstration of Nuclear Magnetic Resonance and Magnetic Resonance Imaging, J. Chem. Educ., Vol. 77, 2000, 882889.
Copyright 2015 Pearson Education, Inc.

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