Chapter Eight

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Chapter Eight

• Electron Configurations, Atomic Properties, And
  The Periodic Table

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Contents

• Multielectron Atoms
• Electron Configurations
• (1) An Introduction
• (2) Aufbau (Building Up) Principle
• (3) Electron Configurations Demonstrations
• (4) Periodic Relationships
• Magnetic Properties Paired and Unpaired
  Electrons
• Periodic Atomic Properties of the Elements
• Metals, Nonmetals, Metalloids, and Nobel Gases
• Using Atomic Properties and the Periodic Table To
  Explain the Behavior of the Elements

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• Multielectron Atoms
• (1) Orbital Energy Diagrams

Single electron species (H, He, Li2)
Multielectron species
subshells of a principle (n) have same Elevel
various subshells of a principle (n) are at
different Elevel
Elevel of same orbitals are lower in a
multielectron species
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(2) More about Elevel of Orbitals

• Either single electron or multielectrons species
• The lower n, the lower energy
• For multielectrons species with same n
• The lower ?, the lower energy, i.e.,
• s lt p lt d lt f
• Two or more orbitals that have the same energy
  level called degenerate orbitals.
• Example 1
• E2px E2py E2pz
• E3dxy E3dyz E3dxz E3dz2 E3dx2-y2
• Example 2
• For H-atom E2s E2p, E3s E3p E3d

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2. Electron Configurations (1) An Introduction

• Definition
• Electron configuration The distribution of
  electrons among the various orbitals in the atom
  (or atomic ion)
• For example

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• Electron configuration presentation ways
• N atom (Z 7) for example

• spdf notation
• 1s22s22p3
• Noble-gas-core abbreviated notation
• He2s22p3
• Expanded spdf notation
• 1s22s22px12py12pz1
• Orbital (box) diagram

- arrows representing electron spins - opposing
spins are paired together
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• Aufbau (Building-Up) Principle
• Rules for (Ground State Neutral Atom) Electron
  Configurations

1) Total number of the building electrons same as
the atomic number (Z) of the particular element.
2) Add electrons to subshells in order with
acquiring electrons
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21
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57
72
57
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Numbers are atomic number
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• Observe Hunds Rule
• Of a group of orbitals of degenerate (identical
  energy) orbitals, the electrons are added into
  each empty orbital with parallel spins before a
  second electron is placed.

4) Observe Pauli exclusion principle (A) No two
electrons in the same atom may have all four
quantum numbers alike (B) An atomic orbital can
accommodate only two electrons, and these
electrons must have opposing spins
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• Electron Configurations Demonstrations
• 1) Maim-Group Elements
• Main-group elements representative elements
• A-group elements

(A) For examples 4Be
Orbital diagram
spdf notation
1s22s2
He2s2
diamagnetic
magnetism
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(B) For examples 5B
Orbital diagram
spdf notation
1s22s22p1
He2s22p1
paramagnetic
magnetism
(C) For examples 7N
Orbital diagram
spdf notation
1s22s22p3
He2s22p3
magnetism
paramagnetic
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(D) For examples 10Ne
Orbital diagram
spdf notation
1s22s22p6
He2s22p6
Ne
magnetism
diamagnetic
(E) For examples 37Rb
Orbital diagram
Kr
5s
spdf notation
Kr 5s1
magnetism
paramagnetic
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• Transition Elements (d-blocks)
• Transition Elements B-group elements
• Some examples for first transition series
• Arrange in order with acquiring electrons

Ar
21Sc
4s
3d
Ar
22Ti
4s
3d
Ar
23V
4s
3d
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• Complete electron configurations for first
  transition series
• All subshells with same principle shell (n) are
  grouped together

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• 3) Exceptions to the Aufbau principle
• Examples for first transition elements
• Expected Observed
• Cr(Z24) Ar4s23d4 Ar4s13d5
• Cu(Z29) Ar4s23d9 Ar4s13d10

Half-filled and complete filled are favorable

• Can be explained by the observations of various
  orbital energies for different atoms

ELevel in general
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• 4) Inner-Transition Elements (f-blocks)
• Inner-Transition Elements Lanthanide elements
  Actinide elements
• Some examples for lanthanide elements
• Ce(Z58) Xe6s24f2
• Yb(Z70) Xe6s24f14

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5) Maximum ecapacities of principle shells (n)
and subshells (l)
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6) A complete periodic table style electron
configurations
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• Electron Configurations of Ions
• Anions examples
• Br(Z35) Ar4s23d104p5 Br Ar4s23d104p6 or
  Kr
• N(Z7) He2s22p3 N3 He2s22p6 or Ne

(B) Cations examples (2nd and 3rd period
metals) Na(Z11) Ne3s1 Na Ne or
1s22s22p6 Al(Z13) Ne3s23p1 Al3 Ne
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(C) Cations examples (4th period and beyond
metals)
ELevel in general

• Ga(Z31) Ar3d104s24p1
• Ga3 Ar3d10
• Sn(Z50) Kr4d105s25p2
• Sn2 Kr4d105s2
• Cr(Z24)Ar4s23d4
• Cr2Ar3d4
• Cr3Ar3d3

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(D) Classified electron configurations of some
metal ions
a style of (n1)s2(n1)p6(n1)d10 b style of
(n1)s2(n1)p6(n1)d10ns2
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(4) Periodic Relationships
1)
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• 2) Valence Electrons and Core Electrons
• Valence shell The outermost (highest n) occupied
  principal shell, valence shell filling electrons
  called valence electrons.
• For main group elements (A-group), the number of
  valence shell electrons is the same as the
  periodic table group number
• The period number is the same as the principal
  quantum number n of the electrons in the valence
  shell.
• Core electrons The electrons in inner shells.

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• Example 1 S(Z16)
• 1s22s22p63s23p4
• Core electrons 1s22s22p6
• Valence electrons 3s23p4 6 electrons
• 3rd period element
• 6A group element
• Example 2 Br(Z35)
• Ar3d104s24p5
• Core electrons Ar3d10
• Valence electrons 4s24p5 7 electrons
• 4th period element
• 7A group element

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• Magnetic Properties Paired and Unpaired
  Electrons
• Definition
• Diamagnetism All the electrons in atoms, ions,
  or molecules of a substance are paired, the
  substance is weekly repelled by a magnetic field,
  the weak repulsion associated with paired
  electrons called diamagnetism.
• Paramagnetism One or more unpaired electrons in
  atoms, ions, or molecules of a substance, the
  substance is attracted by a magnetic field, the
  attraction associated with unpaired electrons
  called paramagnetism.
• Ferromagnetism Exceptionally strong attractions
  of a magnetic field (iron and a few other
  substances).

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2) For paramagnetic or ferromagnetic substances
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• Periodic Atomic Properties of the Elements
• (1) Effective nuclear charge (Zeff)

• 1) Some definition
• Shielding (Screening) Decrease in attraction
  between electrons and the nucleus in
  multielectrons atom because of repulsive forces
  between electrons.
• Electron penetration The ability of electron
  approach the nucleus.
• Shielded Weekly Shielded Strongly
• Strong Penetration Weak Penetration

lower n shell e higher n shell e s
p d f

• Effective nuclear charge (Zeff) The nuclear
  charge acting on an electron

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2) Estimate Zeff of the valence electron
Zeff Z - s Zeff effective nuclear charge Z
actual nuclear charge (atomic number) s
shielding constant (numbers of core electrons)

• Examples for representative elements
• 11Na 1s22s22p63s1 Zeff 11 10 1
• 12Mg 1s22s22p63s2 Zeff 12 10 2
• For same period Z increase Zeff increase
• (B) Examples for transition elements
• 26Fe Ar3d64s2 Zeff 26 24 2
• 27Co Ar3d74s2 Zeff 27 25 2
• 28Ni Ar3d84s2 Zeff 28 26 2
• For 1st transition elements, Zeff are about the
  same

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• Atomic Radii
• 1) Define atomic radii

Metallic radius Half the distance between the
nuclei of adjacent atoms in a solid metal
Covalent radius Half the distance between the
nuclei of two identical atoms joined into a
molecule
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• Trend of atomic radii
• Increase from top to bottom within a group of the
  periodic table. (valence electron penetration
  ability decreased)
• Decrease from left to right in a period of the
  periodic table for representative elements
  (A-group). (Zeff increase)

• Atomic radii of the element

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(3) Ionic Radii 1) Define ionic radii
Determined by crystal structure (Chapter 11)

• Cations Smaller than the atoms from which they
  are formed. (nucleus attracts the remaining
  electrons more strongly)

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3) Anions Larger than the atoms from which they
are formed. (greater number of electrons, repel
more strongly)

• 4) Isoelectronic species The species that all
  have the same number of electrons.
• For a series of isoelectronic species, the
  greater the nuclear charge, the smaller the
  species
• 7N3, 8O2 , 9F are isoelectronic with electron
  configurations Ne, the radius N3 gt O2 gt F
• 11Na, 12Mg2 , 13Al3 are isoelectronic with
  electron configurations Ne, the radius Na gt
  Mg2 gt Al3

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5) Some representative atomic and ionic radii
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• (4) Ionization Energy (IE)
• Definition The energy required to remove an
  electron from a species in the gaseous state.
• First ionization energy (I1)
• X(g) I1 ? X(g) e
• Second ionization energy (I2)
• X(g) I2 ? X2(g) e
• I1 lt I2 lt I3.....
• Removing a core electron takes much more energy
  than removing a valence electron

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• Trend of ionization energy
• I1 decreases from top to bottom on the periodic
  table. (n increases valence electron is farther
  from nucleus).
• I1 generally increases from left to right (with
  exceptions). (Greater Zeff from left to right
  holds electrons more tightly).

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3) Explanation of the irregularity

• Why actually the I1 of group 2A gt I1 of group 3A?
• Ans Be (2A) and B (3A) for example
• The 2p electron (of B) is at the higher energy
  level than 2s electron (of Be), thus it is
  relatively easy to be removed, therefore, I1 of
  Be gt I1 of B.
• Why actually the I1 of group 5A gt I1 of group 6A?
• Ans N (5A) and O (6A) for example
• Repulsion between paired electrons in a 2p
  orbital (of oxygen), lead the electron relatively
  easy to be removed therefore, I1 of N gt I1 of O.

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• (5) Electron affinity (EA)
• Definition The energy change that occurs when an
  electron is added to a gaseous species.
• M(g) e ? M(g) ?H EA

2)

• Negative value exothermic Positive value
  endothermic
• Trend of EA much more irregular than IE

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5. Metals, Nonmetals, Metalloids, and Nobel Gases
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5. Metals, Nonmetals, Metalloids, and Nobel Gases

• Metals
• Small number of electrons in their valence shells
  and tend to form positive ions
• Except for H and He, all s block elements are
  metals
• All d block and f block elements are metals
• A few p block metals are metals
• Metallic character
• - decreases from left to right across a period
• - increases from top to down through a group

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• 2) Nonmetals
• Large number of electrons in their valence shells
  and tend to form negative ions
• Nonmetals are all p block element plus H and He
• Nonmetallic character
• - increases from left to right across a period
• - decreases from top to down through a group
• Metalloids
• B, Si, Ge, As, Sb, Te, Po, At
• Have properties of both metals and nonmetals
• Noble gases (Inert elements)
• He, Ne, Ar, Kr, Xe, Rn
• Rarely enter into chemical reactions

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5) A summary of trends
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• Using Atomic Properties and the Periodic Table To
  Explain the Behavior of the Elements
• From Atomic Properties to chemical properties
• IE The attraction capability of an atom for its
  own electrons
• EA The attraction capability of addition
  electron from other source atom for an atom
• IE and EA can help understand the reaction of the
  elements, and nature of the compound

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2) Flame Colors

• Elements with low IE1 excited in a Bunsen burner
  flame, then emit light in the visible region
  (colored)

• For higher values of IE, higher excitation
  temperatures required, and the emitted light is
  in the UV region.

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• Halogens as oxidizing agents
• The halogens (Group 7A) are good oxidizing agents
• Electron affinities somewhat correlate with the
  oxidizing power
• F2 is the most powerful oxidizing agent
• EA1 for Cl 349 kJ/mole
• EA1 for I 295 kJ/mole
• Spontaneous reaction
• Cl2(g) 2I(aq) ? 2Cl(aq) I2(g)

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• s-block metals as reducing agents
• The s-block metals are powerful reducing agents
• Ionization energy somewhat correlate with the
  reducing power
• The lower IE, the more easily itself is oxidized,
  the better reducing agent.
• All alkali metals and the heavier alkaline earth
  metals are able to displace H2.
• Example
• Mg(s) 2H(aq) ? Mg2(aq) H2(g)

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• Acidic, Basic, and Amphoteric Oxides
• Acidic oxide The oxides that produces an acid
  when reacts with water, generally the oxides of
  nonmetals.
• Example
• SO3(g) H2O(l) ? H2SO4(aq)
• (B) Basic oxides The oxides that produce bases
  when react with water, often, the metal oxides.
• Li2O(s) H2O(l) ? 2LiOH(aq)
• (C) Amphoteric oxide The oxides can react with
  either an acid or a base.
• Al2O3(s) 6HCl(aq) ? 2AlCl3(aq) 3H2O(l)
• Al2O3(s) 2NaOH 3H2O(l) ? 2NaAl(OH)4(aq)

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End of Chapter 08