Electron Configurations ? Chemical Periodicity (Ch 8)

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Electron Configurations ? Chemical Periodicity
(Ch 8)

• Electron spin Pauli exclusion principle
• configurations
• spectroscopic, orbital box notation
• Hunds rule - electron filling rules
• configurations of ATOMS
• the basis for chemical valence
• configurations and properties of IONS
• periodic trends in
• size
• ionization energies
• electron affinities

Na Cl ? NaCl
Mg ?O2 ? MgO
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Arrangement of Electrons in Atoms

• Electrons in atoms are arranged as
• SHELLS (n)
• SUBSHELLS (?)
• ORBITALS (m?)

Each orbital can be assigned up to 2 electrons!
. . . Because there is a
4th quantum number, the electron spin quantum
number, ms.
WHY ?
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Electron Spin Quantum Number, ms

• It can be proved experimentally that the
  electron has a spin. This is QUANTIZED.
• The two allowed spin directions are defined by
  the magnetic spin quantum number, ms
• ms 1/2 and -1/2 ONLY.

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Electron Spin Quantum Number
MAGNETISM is a macroscopic result of quantized
electron spin
5_magnet.mov
Diamagnetic NOT attracted to a magnetic
field All electrons are paired N2
Paramagnetic attracted to a magnetic field.
Substance has unpaired electrons O2
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Pauli Exclusion Principle

• electrons with the same spin keep as far apart
  as possible
• electrons of opposite spin may occupy the same
  region of space ( orbital)
• Consequences

• No orbital can have more than 2 electrons
• No two electrons in the same atom can have the
  same set of 4 quantum numbers (n, l, ml, ms)
• OR
• Each electron has a unique address.

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QUANTUMNUMBERS
n (shell) 1, 2, 3, 4, ... ? (subshell) 0, 1,
2, ... n - 1 m? (orbital) - ? ... 0 ...
? ms (electron spin) 1/2, -1/2
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Shells, Subshells, Orbitals
n ? orbitals e-
Total PERIOD 1 0 s 1 2 2 1 (H, He) 2 0
s 1 2 1 p 3 6 8 2 (LiNe) 3 0 s 1 2 1
p 3 6 3 (Na .. Ar) 2 d 5 10 18 4 0
s 1 2 1 p 3 6 2 d 5 10 3
f 7 14 32 n 0..(n-1) (2 ?1) 2(2 ?1) 2n2
? 0 s ? 1 p ? 2 d ? 3 f
etc, for n 5, 6
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Element Mnemonic Competition
Hey! Here Lies Ben Brown. Could Not Order Fire.
Near Nancy Margaret Alice Sits Peggy Sucking
Clorets. Are Kids Capable ?
WHATs YOURs ??
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Assigning Electrons to Atoms

• Electrons are assigned to orbitals successively
  in order of the energy.
• For H atoms, E - R(1/n2). E depends only on n.
• For many-electron atoms, orbital energy depends
  on both n and ?.
• E(ns) lt E(np) lt E(nd) ...

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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
  ?) increases.

b) for subshells of same (n ?), subshell
with lower n is lower in energy.
5_manyelE.mov
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Effective Nuclear Charge

• The difference in SUBSHELL energy
• e.g. 2s and 2p subshells
• is due to effective nuclear charge, Z.

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Effective Nuclear Charge, Z

• Z is the nuclear charge experienced by an
  electron.
• Z increases across a period owing to incomplete
  shielding by inner electrons.
• For VALENCE electrons we estimate Z as

Z Z - (no. of 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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Photoelectron Spectroscopy - Measuring IE
Photoelectric effect h? A ? A
e- forms basis for DIRECT determination of IE
Kinetic energy of electron h? - IE therefore
IE h? - KE(e-)
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Electron Filling Order (Figure 8.7)
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Writing Atomic Electron Configurations

• Two ways of writing configurations.
• One is called the spectroscopic notation

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Writing Atomic Electron Configurations (2)

• A second way is called the orbital box notation.

One electron has n 1, ? 0, ml 0, ms
1/2 Other electron has n 1, ? 0, ml 0, ms
- 1/2
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Electron Configuration tool - see toolbox.
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LithiumGroup 1AZ 3 1s22s1
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BoronZ 5 1s2 2s2 2p1
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CarbonZ 61s2 2s2 2p2
The configuration of C is an example of HUNDS
RULE the lowest energy arrangement of electrons
in a subshell is that with the MAXIMUM no. of
unpaired electrons
Electrons in a set of orbitals having the same
energy, are placed singly as long as possible.
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NitrogenZ 71s2 2s2 2p3
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FluorineZ 9 1s2 2s2 2p5
Note that we have reached the end of the 2nd
period, . . . and the 2nd shell is full!
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GROUPS and PERIODS

• or neon core 3s1
• Ne 3s1 (uses rare gas notation)
• Na begins a new period.
• All Group 1A elements

Li Na K Rb Cs
have core ns1 configurations. (n period )
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Periodic Chemical Properties
Li Na K Rb Cs Alkalis
5_Li.mov
5_Na.mov
5_K.mov
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Alkaline Earths
Metals (ns2) - easily oxidized to M2 -
less reactive than alkalis of same
period reactivity Be lt Mg lt Ca lt Sr lt Ba WHY? -

• Size INCREASES as ? group
• VALENCE e- are farther from nucleus
• same Z - Valence e- less tightly held
• Therefore valence e- are easier to remove

Typical reactions / compounds Oxides M 1/2O2
(g) ? MO (s) CaO (lime) - 5 Ind.
Chem Halides M X2 (g) ? MX Carbonates
CaCO3 (limestone) ? CaO CO2
Sulfates CaSO4.2H2O (gypsum) ? CaSO4.
0.5H2O (plaster-of-paris) 3/2H2O
RECALL Solubility rules and PRECIPITATION
REACTIONS
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Relationship of Electron Configuration and
Regions of the Periodic Table
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Transition Metals Table 8.4

• Transition metals (e.g. Sc .. Zn in the 4th
  period) have the configuration argon nsx (n -
  1)dy
• also called d-block elements.

3d orbitals used for Sc - Zn
Copper
Iron
Chromium
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Ion Configurations
To form cations from elements remove 1 e- (or
more) from subshell of highest n or highest (n
?).

• P Ne 3s2 3p3 - 3e- ? P3 Ne 3s2 3p0

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Ion Configurations (2)

• Transition metals ions
• remove ns electrons and then (n - 1)d electrons.

Fe Ar 4s2 3d6 loses 2 electrons ? Fe2 Ar
4s0 3d6
E4s E3d - exact energy of orbitals depend on
whole configuration
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Ion Configurations (3)
How do we know the configurations of ions?

• From the magnetic properties of ions.
• Ions (or atoms) with UNPAIRED ELECTRONS are
• PARAMAGNETIC.
• Ions (or atoms) without unpaired electrons are
• DIAMAGNETIC.

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General Periodic Trends

• Atomic and ionic radii SIZE
• Ionization energy E(A) - E(A)
• Electron affinity E(A-) - E(A)

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Atomic Size INCREASESdown a Group

• Size goes UP on going down a GROUP
• Because electrons are added further from the
  nucleus, there is less attraction.

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Atomic Size DECREASES across a period

• Size goes DOWN on going across a PERIOD.
• Size decreases due to increase in Z.
• Each added electron feels a greater and greater
  ve charge.

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Atomic Radii
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Trends in Atomic Size (Figure 8.10)
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Sizes of Transition Elements(Figure 8.11)

• 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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Ion Sizes - CATIONS
Does the size go up or down when an atom loses an
electron to form a cation?

Forming a cation

• CATIONS are SMALLER than the parent atoms.
• The electron/proton attraction goes UP so size
  DECREASES.

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Ion Sizes - ANIONS

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

Forming an anion

• ANIONS are LARGER than the parent atoms.
• electron/proton attraction goes DOWN so size
  INCREASES.

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Trends in Ion Sizes
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Oxidation-Reduction Reactions

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

- related to IE and EA
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Ionization Energy (IE)
Mg (g) atom Ne2s

• Mg (g) 735 kJ ? Mg (g) e- Ne2s1

Mg (g) 1451 kJ ? Mg2 (g) e- Ne2s0
Mg2 (g) 7733 kJ ? Mg3 (g) e- He2s22p5

• Energy cost is very high to remove an INNER
  SHELL e- (shell of n lt nVALENCE).
• This is why oxidation. no. Group no.

Mg
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Trends in First Ionization Energy
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Trends in Ionization Energy (2)

• 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.

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

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2nd IE / 1st IE
2nd IE A ? A e-
Li
Na
K
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Electron Affinity (EA)

• A few elements GAIN electrons to form anions.
• Electron affinity is the energy released when an
  atom gains an electron.
• A(g) e- ? A-(g) E.A. DE E(A-) -
  E(A)
• If E(A-) lt E(A) then the anion is more stable
  and there is an exothermic reaction

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Trends in Electron Affinity (Table 8.5, Figure
8.14)
Atom EA (kJ) B -27 C -122 N
0 O -141 F -328

• Affinity for electron increases across a period
• (EA becomes more negative).

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SUMMARY

• Electron spin diamagnetism vs. paramagnetism
• Pauli exclusion principle - allowable quantum
  numbers
• configurations
• spectroscopic notation
• orbital box notation
• Hunds rule - electron filling rules
• configurations of ATOMS the basis for chemical
  valence
• period 2 groups
• transition metals
• configurations and properties of IONS
• periodic trends in
• size
• ionization energies
• electron affinities