ATOMIC ELECTRON CONFIGURATIONS AND ORBITAL SHAPES

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ATOMIC ELECTRON CONFIGURATIONS AND ORBITAL SHAPES
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Arrangement of Electrons in Atoms

• Electrons in atoms are arranged as
• SHELLS (n)
• SUBSHELLS (l)
• ORBITALS (ml)

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Arrangement of Electrons in Atoms

• Each orbital can be assigned no more than 2
  electrons!
• This is tied to the existence of a 4th quantum
  number, the electron spin quantum number, ms.

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Electron Spin Quantum Number, ms
Can be proved experimentally that electron has a
spin. Two spin directions are given by ms where
ms 1/2 and -1/2.
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QUANTUM NUMBERS
Now there are four!

• n ---gt shell 1, 2, 3, 4, ...
• l ---gt subshell 0, 1, 2, ... n - 1
• ml ---gt orbital -l ... 0 ... l
• ms ---gt electron spin 1/2 and -1/2

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Pauli Exclusion Principle

• No two electrons in the same atom can have the
  same set of 4 quantum numbers.
• That is, each electron has a unique address.

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Electrons in Atoms

• When n 1, then l 0
• this shell has a single orbital (1s) to which
  2e- can be assigned.
• When n 2, then l 0, 1
• 2s orbital 2e-
• three 2p orbitals 6e-
• TOTAL 8e-

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Electrons in Atoms

• When n 3, then l 0, 1, 2
• 3s orbital 2e-
• three 3p orbitals 6e-
• five 3d orbitals 10e-
• TOTAL 18e-

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Electrons in Atoms

• When n 4, then l 0, 1, 2, 3
• 4s orbital 2e-
• three 4p orbitals 6e-
• five 4d orbitals 10e-
• seven 4f orbitals 14e-
• TOTAL 32e-

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Assigning Electrons to Atoms

• Electrons generally assigned to orbitals of
  successively higher energy.
• For H atoms, E - C(1/n2). E depends only on n.
• For many-electron atoms, energy depends on both n
  and l.
• See Active Figure 8.4, Figure 8.5, and Screen 8.
  7.

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Assigning Electrons to Subshells

• In H atom all subshells of same n have same
  energy.
• In many-electron atom
• a) subshells increase in energy as value of n l
  increases.
• b) for subshells of same n l, subshell with
  lower n is lower in energy.

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Effective Nuclear Charge, Z

• Z is the nuclear charge experienced by the
  outermost electrons. See Figure 8.6 and and
  Screen 8.6.
• Explains why E(2s) lt E(2p)
• Z increases across a period owing to incomplete
  shielding by inner electrons.
• Estimate Z by --gt Z - (no. inner electrons)
• Charge felt by 2s e- in Li Z 3 - 2 1
• Be Z 4 - 2 2
• B Z 5 - 2 3 and so on!

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Effective Nuclear Charge
Figure 8.6
Z is the nuclear charge experienced by the
outermost electrons.
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What does shielding do?

• Well It causes the subshells to have unequal
  energy.
• Therefore, the energy levels fill in a different
  order.
• Shielding accounts for many periodic properties.

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Electron Filling OrderFigure 8.5
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Orbital Energies
Orbital energies drop as Z increases
CD-ROM Screens 8.9 - 8.13, Simulations
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Writing Atomic Electron Configurations

• Two ways of writing configs. One is called the
  spdf notation.

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Writing Atomic Electron Configurations

• Two ways of writing configs. Other is called the
  orbital box notation.

One electron has n 1, l 0, ml 0, ms
1/2 Other electron has n 1, l 0, ml 0, ms
- 1/2
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See Toolbox on CD for Electron Configuration
tool.
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Electron Configurations and the Periodic Table
Active Figure 8.7
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Lithium

• Group 1A
• Atomic number 3
• 1s22s1 ---gt 3 total electrons

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Beryllium

• Group 2A
• Atomic number 4
• 1s22s2 ---gt 4 total electrons

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Boron

• Group 3A
• Atomic number 5
• 1s2 2s2 2p1 ---gt
• 5 total electrons

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Carbon

• Group 4A
• Atomic number 6
• 1s2 2s2 2p2 ---gt
• 6 total electrons

Here we see for the first time HUNDS RULE. When
placing electrons in a set of orbitals having the
same energy, we place them singly as long as
possible.
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Nitrogen

• Group 5A
• Atomic number 7
• 1s2 2s2 2p3 ---gt
• 7 total electrons

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Oxygen

• Group 6A
• Atomic number 8
• 1s2 2s2 2p4 ---gt
• 8 total electrons

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Fluorine

• Group 7A
• Atomic number 9
• 1s2 2s2 2p5 ---gt
• 9 total electrons

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Neon

• Group 8A
• Atomic number 10
• 1s2 2s2 2p6 ---gt
• 10 total electrons

Note that we have reached the end of the 2nd
period, and the 2nd shell is full!
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Electron Configurations of p-Block Elements
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Sodium

• Group 1A
• Atomic number 11
• 1s2 2s2 2p6 3s1 or
• neon core 3s1
• Ne 3s1 (uses rare gas notation)
• Note that we have begun a new period.
• All Group 1A elements have corens1
  configurations.

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Aluminum

• Group 3A
• Atomic number 13
• 1s2 2s2 2p6 3s2 3p1
• Ne 3s2 3p1

All Group 3A elements have core ns2 np1
configurations where n is the period number.
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Phosphorus

• Group 5A
• Atomic number 15
• 1s2 2s2 2p6 3s2 3p3
• Ne 3s2 3p3

All Group 5A elements have core ns2 np3
configurations where n is the period number.
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Calcium

• Group 2A
• Atomic number 20
• 1s2 2s2 2p6 3s2 3p6 4s2
• Ar 4s2
• All Group 2A elements have corens2
  configurations where n is the period number.

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Electron Configurations and the Periodic Table
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Transition MetalsTable 8.4

• All 4th period elements have the configuration
  argon nsx (n - 1)dy and so are d-block
  elements.

Copper
Iron
Chromium
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Transition Element Configurations
3d orbitals used for Sc-Zn (Table 8.4)
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Lanthanides and Actinides

• All these elements have the configuration core
  nsx (n - 1)dy (n - 2)fz and so are f-block
  elements.

Cerium Xe 6s2 5d1 4f1
Uranium Rn 7s2 6d1 5f3
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Lanthanide Element Configurations
4f orbitals used for Ce - Lu and 5f for Th - Lr
(Table 8.2)
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Ion Configurations

• To form cations from elements remove 1 or more e-
  from subshell of highest n or highest (n l).
• P Ne 3s2 3p3 - 3e- ---gt P3 Ne 3s2 3p0

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Ion Configurations

• For transition metals, remove ns electrons and
  then (n - 1) electrons.
• Fe Ar 4s2 3d6
• loses 2 electrons ---gt Fe2 Ar 4s0 3d6

To form cations, always remove electrons of
highest n value first!
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Ion Configurations

• How do we know the configurations of ions?
• Determine the magnetic properties of ions.

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Ion Configurations

• How do we know the configurations of ions?
• Determine the magnetic properties of ions.
• Ions with UNPAIRED ELECTRONS are PARAMAGNETIC.
• Without unpaired electrons DIAMAGNETIC.

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PERIODIC TRENDS
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General Periodic Trends

• Atomic and ionic size
• Ionization energy
• Electron affinity

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Effective Nuclear Charge, Z

• Z is the nuclear charge experienced by the
  outermost electrons. See Figure 8.6 and and
  Screen 8.6.
• Explains why E(2s) lt E(2p)
• Z increases across a period owing to incomplete
  shielding by inner electrons.
• Estimate Z by --gt Z - (no. inner electrons)
• Charge felt by 2s e- in Li Z 3 - 2 1
• Be Z 4 - 2 2
• B Z 5 - 2 3 and so on!

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Effective Nuclear Charge
Figure 8.6
Z is the nuclear charge experienced by the
outermost electrons.
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Effective Nuclear ChargeZ

• The 2s electron PENETRATES the region occupied by
  the 1s electron.
• 2s electron experiences a higher positive charge
  than expected.

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Effective Nuclear Charge, Z

• Atom Z Experienced by Electrons in Valence
  Orbitals
• Li 1.28
• Be -------
• B 2.58
• C 3.22
• N 3.85
• O 4.49
• F 5.13

Increase in Z across a period
Values calculated using Slaters Rules
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General Periodic Trends

• Atomic and ionic size
• Ionization energy
• Electron affinity

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Atomic Radii
Active Figure 8.11
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Atomic Size

• Size goes UP on going down a group. See Figure
  8.9.
• Because electrons are added further from the
  nucleus, there is less attraction.
• Size goes DOWN on going across a period.

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Atomic Size

• Size decreases across a period owing to increase
  in Z. Each added electron feels a greater and
  greater charge.

Large
Small
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Trends in Atomic SizeSee Active Figure 8.11
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Sizes of Transition ElementsSee Figure 8.12
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Sizes of Transition ElementsSee Figure 8.12

• 3d subshell is inside the 4s subshell.
• 4s electrons feel a more or less constant Z.
• Sizes stay about the same and chemistries are
  similar!

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Density of Transition Metals
6th period
5th period
4th period
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Ion Sizes
Does the size go up or down when losing an
electron to form a cation?

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Ion Sizes
Forming a cation.
Li,152 pm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they come.
• The electron/proton attraction has gone UP and so
  size DECREASES.

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Ion Sizes

• Does the size go up or down when gaining an
  electron to form an anion?

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Ion Sizes
Forming an anion.

• ANIONS are LARGER than the atoms from which they
  come.
• The electron/proton attraction has gone DOWN and
  so size INCREASES.
• Trends in ion sizes are the same as atom sizes.

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Trends in Ion Sizes
Active Figure 8.15
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Redox Reactions

• Why do metals lose electrons in their reactions?
• Why does Mg form Mg2 ions and not Mg3?
• Why do nonmetals take on electrons?

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Ionization EnergySee CD Screen 8.12

• IE energy required to remove an electron from
  an atom in the gas phase.

Mg (g) 738 kJ ---gt Mg (g) e-
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Ionization EnergySee Screen 8.12
IE energy required to remove an electron from
an atom in the gas phase.

• Mg (g) 738 kJ ---gt Mg (g) e-

Mg (g) 1451 kJ ---gt Mg2 (g) e-
Mg has 12 protons and only 11 electrons.
Therefore, IE for Mg gt Mg.
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Ionization EnergySee Screen 8.12

• Mg (g) 735 kJ ---gt Mg (g) e-
• Mg (g) 1451 kJ ---gt Mg2 (g) e-

Mg2 (g) 7733 kJ ---gt Mg3 (g) e-
Energy cost is very high to dip into a shell of
lower n. This is why ox. no. Group no.
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2nd IE / 1st IE
Li
Na
K

• B

Al
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Trends in Ionization Energy
Active Figure 8.13
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Trends in Ionization Energy
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Orbital Energies
As Z increases, orbital energies drop and IE
increases.
CD-ROM Screens 8.9 - 8.13, Simulations
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Trends in Ionization Energy

• IE increases across a period because Z
  increases.
• Metals lose electrons more easily than nonmetals.
• Metals are good reducing agents.
• Nonmetals lose electrons with difficulty.

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Trends in Ionization Energy

• IE decreases down a group
• Because size increases.
• Reducing ability generally increases down the
  periodic table.
• See reactions of Li, Na, K

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Periodic Trend in the Reactivity of Alkali Metals
with Water
Lithium
Sodium
Potassium
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Electron Affinity

• A few elements GAIN electrons to form anions.
• Electron affinity is the energy involved when an
  atom gains an electron to form an anion.
• A(g) e- ---gt A-(g) E.A. ?E

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Electron Affinity of Oxygen

• ?E is EXOthermic because O has an affinity for an
  e-.

EA - 141 kJ
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Electron Affinity of Nitrogen

• ?E is zero for N- due to electron-electron
  repulsions.

EA 0 kJ
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Trends in Electron Affinity
Active Figure 8.14
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Trends in Electron Affinity

• See Figure 8.14 and Appendix F
• Affinity for electron increases across a period
  (EA becomes more positive).
• Affinity decreases down a group (EA becomes less
  positive).

Atom EA F 328 kJ Cl 349 kJ Br 325 kJ I 295
kJ
Note effect of atom size on F vs. Cl