Chapter 8 Electron Configurations, Atomic Properties, and the Periodic Table General Chemistry: An...
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Transcript of Chapter 8 Electron Configurations, Atomic Properties, and the Periodic Table General Chemistry: An...
Chapter 8Electron Configurations, Atomic Properties, and the Periodic Table
General Chemistry: An Integrated Approach
Hill, Petrucci, 4th Edition
Mark P. HeitzState University of New York at Brockport
© 2005, Prentice Hall, Inc.
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
2
Multielectron AtomsElectrons are attracted to the nucleus while simultaneously repelling one another
EOS
In the hydrogen atom, all subshells of a principal shell are at the same energy level
recall En = –B/n2
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
3
Multielectron Atoms
The increasing energy order of subshells is generally:
s < p < d < fEOS
In a multielectron atom the various subshells of a principal shell are at different energy levels, but all orbitals within a subshell are at the same energy level
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
4
Multielectron Atoms
In higher numbered principal shells of a multielectron atom, some subshells of different principal shells have nearly identical energies
EOS
Orbital energies are lower in multielectron atoms than in the hydrogen atom
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
5
Electron Configurations
Electron configuration describes the distribution of electrons among the various orbitals in the atom
The spdf notation uses numbers to designate a principal shell and the letters to identify a subshell; a superscript number indicates the number of electrons in a designated subshell
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
6
Formula Mass
EOS
Each box has arrows representing electron spins; opposing spins are paired together
An orbital diagram uses boxes to represent orbitals within subshells and arrows to represent electrons:
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
7
Rules for Electron Configurations
Electrons occupy the lowest available energy orbitals
Pauli exclusion principle – no two electrons in the same atom may have the same four quantum numbers
EOS
Orbitals hold a maximum of two electronsspins must be opposed
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Rules for Electron Configurations
For orbitals of identical energy, electrons enter empty orbitals whenever possible – Hund’s rule
Electrons in half-filled orbitals have parallel spins
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Rules for Electron Configurations
EOS
Capacities of shells (n) and subshells (l)
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Rules for Electron Configurations
Subshell filling order ...
Each subshell must be filled before moving to the next level
EOS
1s22s22p63s23p6 ...
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
11
The Aufbau Principle
A hypothetical building up of an atom from the one that precedes it in atomic number
(Z = 1) H 1s1
(Z = 2) He 1s2
(Z = 3) Li 1s22s1
EOS
(Z = 3) Li 1s22s1 [He]2s1
Abbreviated electron configuration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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The Aufbau Principle ...[He]2p2
[He]2p3
[He]2p4
[He]2p5
EOS
[He]2p6
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Main Group and Transition Elements
Elements in which the orbitals being filled in the aufbau process are either s or p orbitals of the outermost shell are called main group elements
“A” group designation on the periodic tableThe first 20 elements are all main group elements
In transition elements, the subshell being filled in the aufbau process is in an inner principal shell
EOS
Fourth period transition elements have n = 4 as their outermost shell as the 3d subshell fills
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Main Group and Transition Elements
EOS
Completely filled and half-filled sublevels are more energetically favorable configurations
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Periodic RelationshipsThe valence shell is the outermost occupied shell
The period number = principal quantum number, n, of the electrons in the valence shell
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Periodic Relationships
EOS
For main group elements the number of valence shell electrons is the same as the periodic table “A” group number
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Periodic RelationshipsWe can deduce the general form of electron configurations directly from the periodic table
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Valence Electrons and Core Electrons
Valence electrons are those with the highest principal quantum number
EOS
Sulfur has six valence electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Valence Electrons and Core Electrons
Electrons in inner shells are called core electrons
EOS
Sulfur has 10 core electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Electron Configurations of IonsAnions: gain e– to complete the valence shell
Example:
EOS
-
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Electron Configurations of IonsCations: lose e– to attain a complete valence shell
Example:
(Z = 11) Na
EOS
(Z = 11) Na+
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Electron Configurations of IonsCations formed from transition metals lose e– from the highest principal energy level (n)
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Magnetic PropertiesDiamagnetism is the weak repulsion associated with paired electrons
Paramagnetism is the attraction associated with unpaired electrons
EOS
Ferromagnetism is the exceptionally strong attractions of a magnetic field for iron and a few other substances
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
24
Periodic Trends
• Use must justify the trend across the period, you cannot simply state the trend.
• A trend is an observation, not an explanation!
• You should state the trend in your answer, but you must also go further by explaining what causes the observed trend!
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
25
Periodic Atomic Properties of the Elements
Periodic law states that certain sets of physical and chemical properties recur at regular intervals when the elements are arranged according to increasing atomic number
EOS
• Consider atomic radii: distance between the nuclei of two atoms • The distance between the nucleus and the outer
edge of the electron cloud
26
Atomic Radii
• Atomic Radii decrease as atomic numbers increase in an given period (going across).– A proton and electron are added so the effective
nuclear charge increases because each proton has more of an effect than each additional electron
• As that attraction between the nucleus and electrons increases, and the atomic radius decreases
• Atomic Radii increase Sgoing down– In going from top to bottom of a group, the
valence electrons are assigned to orbitals with increasingly higher values of n (prin. Quantum number)
• The underlying electrons requires some space, so the electrons of the outer shell must be further (Your are adding energy levels)
27
Atomic Radii
•Zeff effective nuclear charge: the nuclear charge experienced by a particular electron in a multielectron atom
– Increases the attraction of the nucleus and pulls the electron cloud closer to the nucleus resulting in a smaller atomic radius
•Atomic radii of transtion metals trend a little differently•Exceptions in atomic radii also exist in the lanthanide and actinide series because of how the f subshells are uniquely filled by electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Zeff & Shielding
– The order in which electrons are assigned to subshells in an atom, as well as other properties are because of Zeff
– Shielding: electrons closer to the nucleus screen or shield the effect of nuclear charge on valence electrons• the number of shielding electrons increases
when you reach the end of the periodic table and go on to the next period.
• Shielding increases in steps as you start a new period or go down a group
• Video
29
Transition Metal Atomic Trends
• From left to right across a period, the radii initially decrease, then size remains almost the same, then slightly increases toward the end.
•The small increase in atomic radii is because of the d subshell is filled with electrons and thus the ele-eletron repulsions cause the size to increase
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Atomic Radii Properties
• The increased number of energy levels (n) increases the distance over which the nucleus must pull and therefore reduces the attraction for electrons
• Full energy levels provide shielding between the nucleus and valence electrons, so you see an increase in shielding as the level gets full EOS
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Ionic Radii
EOS
• The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion• If the size of an atom is
determined by the outermost electrons, what happens if you remove or add an electron?
32
Ionic RadiiCations are smaller than the atoms from which they are formed
– the nucleus attracts the remaining electrons more strongly
EOS
Anions are larger than the atoms from which they are formed
– the greater number of electrons repel more stronglyThink of the proton/electron ratio, -as electrons are lost, the ratio of p+/e- increases and so the electrons are held closer vv.
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Isoelectronic Configurations
Isoelectronic species are elements that all have the same number of electrons
For isoelectronic species, the greater the nuclear charge, the smaller the species
EOS
Effective nuclear charge
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
34
Atomic and Ionic Radii
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Ionization Energy
Ionization energy is the energy required to remove an electron from a ground state atom in the gaseous state• to remove an electron, energy must be supplied to overcome
the attraction of the nuclear charge (endothermic, always +)
Continual removal of electrons increases ionization energy greatly B B+ + e– I = 801 kJ mol–1
B+ B+2 + e– I = 2427 kJ mol–1
B+2 B+3 + e– I = 3660 kJ mol–1
B+3 B+4 + e– I = 25,025 kJ mol–1
EOS
B+4 B+5 + e– I = 32,822 kJ mol–1Illustration
36
Ionization energy• First ionization energy- energy is increased with
each successive removal because the electron is being removed from an increasingly positive ion– The remaining electrons are held more tightly– Notice the large jump at the 3rd level for Mg.– There is a large increase as you remove electrons from
lower (inner) energy subshells
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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First Ionization Energies
EOS
Illustration
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Ionization energy
• Ionization energy increases as atomic number increases in any given period– Zeff increases the attraction of the nucleus and holds the
electrons more tightly• Exceptions: group II to III, IE drops because the p electrons do
not penetrate the nuclear region as well as s electrons so aren’t as tightly held
• Drop in IE also occurs between V & VI because of increased repulsion created by the first pairing of electrons, that is stronger than the increase in Zeff, lowering the energy required to remove the electron
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
39
Ionization energy
• Ionization energy decreases as atomic number increases down a column or group– The increased number of energy levels (n)
increases the distance over which the nucleus must pull, reducing the attraction for electrons
– A full energy level provides some shielding between the nucleus and valence electrons
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
40
Electron Affinity
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom- How much an atom ‘likes’Electrons (+ or -)-the more negative it is the higher the EA
(energy is flowing out of the system)
EOS
Electron affinities are expressed as negative because the process is exothermic Illustration
41
Electronegativity
• A measure of the attraction of an atom for the pair of outer shell electrons in a covalent bond with another atom– Pattern is same as electron affinity for same
reasons– Both are attraction nucleus has for electrons,
one in forming an ion (EA) and one in forming a molecule (EN)
– Fluorine is the most electronegative. The closer it is to fluorine, the more electronegative it is.
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
42
Metals, Nonmetals, and MetalloidsMetals have a small number of electrons in their valence shells and tend to form positive ions
Except for hydrogen and helium, all s-block elements are metals
EOS
All d- and f-block elements are metals
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Metals, Nonmetals, and MetalloidsAtoms of a nonmetal generally have larger numbers of electrons in their valence shell than do metals, and many tend to form negative ions
Nonmetals are all p-block elements and include hydrogen and helium
EOS
Metalloids have properties of both metals and nonmetals
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Metals
• Metals react by losing electrons– A loosely held electron will result in a more
reactive metal– This is tied directly to ionization energy– With an increased # of energy levels (n), comes
increased distance from the nuclear attraction and thus a more loosely held electron available for reactions
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
45
Non-metals
• Non-metals tend to gain electrons, a strong nuclear attraction will result in a more reactive non-metal
• This means that an atom with the highest Zeff and the least number of energy levels should be the most reactive nonmetal (F) because its nucleus exerts the strong pull
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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A Summary of Periodic Trends
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
47
The Inert (Noble) Gases
The six noble gases, He, Ne, Ar, Kr, Xe, and Rn, rarely enter into chemical reactions
EOS
All have complete octets ... = stability!
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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“Periodic” Behavior of Elements
Flame tests: elements with low first ionization energies are excited in a flame
EOS
Atoms emit energy when electrons fall from higher to lower energy states
FlameTests
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
49
“Periodic” Behavior of Elements
halogens (Group 7A) are good oxidizing agents
EOS
When Cl2 is bubbled in a solution containing iodide ions, chlorine oxidizes I– to I2
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
50
The s-Block Metals as Reducing Agents
EOS
Recall activity series ...H+ is reduced by these metals
2 K + 2 H2O 2 K+ + 2 OH– + H2
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
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Acidic, Basic, and Amphoteric Oxides
Acidic oxides are oxides that produce acids by reacting the oxide with water
e.g., SO3 + H2O H2SO4
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
52
Acidic, Basic, and Amphoteric Oxides
Basic oxides are oxides that produce bases by reacting with water
e.g., MgO + H2O Mg(OH)2
EOS
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
53
Acidic, Basic, and Amphoteric Oxides
Oxides that can react with either acids or bases are amphoteric oxides
e.g., Al2O3
EOS
Behavior of Oxides
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
54
Summary of Concepts
• The wave-mechanical treatment of the hydrogen atom can be extended to multielectron atoms, but with two differences
• Electron configuration is the distribution of electrons in orbitals among the subshells and principal subshells
EOS
• There are two types of electron configuration notation:
spdf and orbital
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
55
Summary of Concepts
• The aufbau principle describes a process of hypothetically building up an atom from the atom of the preceding atomic number
• Elements in similar electron configurations fall in the same group of the periodic table
• An atom with all the electrons paired is diamagnetic; an atom with one or more unpaired electrons is paramagnetic
EOS
• Certain atomic properties, such as atomic radius, ionic radius, ionization energy, and electron affinity, vary periodically with increasing atomic number
Chapter 8: Electron Configurations, Atomic Properties and the Periodic Table
56
Summary of Concepts
• The regions of the periodic table ascribed to metals, nonmetals, metalloids, and the noble gases are related to the value of atomic properties
EOS