chapter 8s-bates/chem171/ch8presstudent.pdf · draw orbital diagrams ... periodic trends the goal...
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Chapter 8:Electron Configurations
and Periodicity
Electron Spin & the Pauli Exclusion Principle
3 quantum numbers (n, l, ml) define the energy, size, shape, and spatial orientation of each atomic orbital.
To explain how electrons populate atomic orbitals in an atom, we need a 4th quantum number.
4. spin quantum number, ms
◆ electrons behave like they are spinning around an axis
◆ spinning of a charged particle creates a magnetic field
magnetic field created has a direction
if e– spins clockwise, magnetic field is one direction (ms = +!)
if e– spins counter-clockwise, magnetic field is opposite direction (ms = –!)
Experimental Evidence of Electron Spin:Stern & Gerlach (1921)
How Do Electrons Populate Atomic Orbitals?Pauli Exclusion Principle (1925)
No 2 electrons in an atom can have the same set of 4 quantum numbers.
◆ n, l, ml define the atomic orbital
◆ ms will define the electrons in the orbital
If there are only 2 possible ms values, then each atomic orbital can hold no more than 2 electrons;
specifically, one e– must have ms = +!and the other e– must have ms = –!
◆ These 2 e–’s in an atomic orbital are said to be spin-paired.
Effective Nuclear Charge (Zeff)What is the charge felt by an electron in an atom?
◆ depend on n and l of the orbital where the e– lives
◆ the farther the electron is from the nucleus:the less the force of attraction to the nucleusthe greater the e– - e– repulsion
◆ an outer shell electron is “shielded” by inner shell e–’s
Effective Nuclear Charge (Zeff)
Zeff = Zactual – shielding factor
Zeff and Atomic Orbitals
◆ relationship between Zeff and n
probability of an e– being close to the nucleus:
n = 1 > n = 2 > n = 3
Zeff felt by an e– in an orbital:
n = 1 > n = 2 > n = 3
energy of orbital:
n = 1 < n = 2 < n = 3
Zeff and Atomic Orbitals
◆ relationship between Zeff and l (within a same shell)
probability of an e– being close to the nucleus:l = 0 > l = 1 > l = 2
or: s > p > d
Zeff felt by an e– in an orbital:l = 0 > l = 1 > l = 2
or: s > p > d
energy of orbital:s < p < d
Electron Configurations
◆ give complete electronic description for every element
◊ predict orbitals occupied by electrons
◊ write electron configurations
◊ draw orbital diagrams
◆ follow set of 3 rules: Aufbau (“building up”) Principle
◊ each successive electron will go into the lowest energy orbital available
◊ this results in the lowest energy, ground state configuration
The Aufbau Principle1. Lower energy orbitals fill before higher energy orbitals.
◆ use the Relative Energies of Atomic Orbitals diagram
2. An orbital can accommodate a maximum of 2 e–’s which must be spin-paired.
◆ Pauli Exclusion Principle
3. Hund’s Rule:If 2 or more degenerate orbitals are available, one e– will go into each orbital until all are half-full.
◆ the e–’s in the singularly populated orbitals must have the same ms
After all degenerate orbitals are half-full, then a 2nd e– may be added to fill the orbitals.
In What Order Do Atomic Orbitals Populate?Lower Energy to Higher Energy
examples:
Write the ground state electron configurations, and complete an orbital diagram for neutral atoms of the following elements.
N (7 e–) : 1s2 2s2 2p3
Al (13 e–) : 1s2 2s2 2p6 3s2 3p1
1s 2s 2p
1s 2s 2p 3s 3p
↑↓ ↑↓ ↑ ↑ ↑
↑↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
Closed Shells and Subshells & Using Noble Gas Core Symbolism in Electron Configurations
Sc (21 electrons): 1s2 2s2 2p6 3s2 3p6 4s2 3d1
1s 2s 2p 3s 3p 4s 3d
Ga (31 electrons): [Ar] 4s2 3d10 4p1
4s 3d 4p
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
↑↓↑↓ ↑↓ ↑↓ ↑↓ ↑↓
↑
↑
Some Terminology:
inner shell electrons outer shell electrons
core electrons valence electrons
paramagnetic diamagnetic
Some Anomalies in Electron Configurations:
◆ result from unusual stability of half-filled or completely filled shells or subshells
◆ heavy elements (above atomic # 40) ∆E between orbitals is smaller, so anomalies are common
4s 3d 4s 3d
4s 3d 4s 3d
!
!
ex: Cr predict: [Ar] 4s2 3d4 actual: [Ar] 4s1 3d5
ex: Cu predict: [Ar] 4s2 3d9 actual: [Ar] 4s1 3d10
↑↓
↑↓ ↑↓ ↑↓↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑
↑ ↑
Electron Configurations and the Periodic Table
Periodic Trends
The goal is to use our understanding of electron configurations and Zeff to understand trends in:
◆ atomic radius
◆ ionization energy
◆ electron affinity
How does a given property change from left to right across periodic table?
How does a given property change from
top to bottom of periodic table?
Periodic Trend in Zeff
Effective Nuclear Charge,Zeff
increase
decrease
Na Mg Al Si P S Cl Are–
configuration[Ne]3s1 [Ne]3s2 [Ne]3s2 3p1 [Ne]3s2 3p2 [Ne]3s2 3p3 [Ne]3s2 3p4 [Ne]3s2 3p5 [Ne]3s2 3p6
actual nuclear charge 11 12 13 14 15 16 17 18
# core e–’s 10 10 10 10 10 10 10 10
# valence e–’s 1 2 3 4 5 6 7 8
Zeff +1 +2 +3 +4 +5 +6 +7 +8
Atomic Radius
radius (typically in pm or Å) of neutral atoms of elements
trend:atomic radius decreases left to right across the periodic table
atomic radius increases top to bottom of the periodic table
Atomic Radius
decrease
increase
Periodic Trend in Atomic Radius
Periodic Trend in Atomic Radius
Ionization Energy
increasedecrease
Ionization Energy
ionization energy – the energy required to remove an electron from a gas phase atom or ion in its ground state
X (g) " X+ (g) + e– ; endothermic
Periodic Trend in Ionization Energy
◆ Note the unexpected changes between groups IIA & IIIA, and groups VA & VIA. Why? Think about e– configurations.
Periodic Trend in Ionization Energyconsider successive ionization energies:
M (g) ! M+ (g) + e– 1st ionization energyM+ (g) ! M2+ (g) + e– 2nd ionization energyM2+ (g) ! M3+ (g) + e– 3rd ionization energy
◆ It becomes successively harder to remove an e– from a positively charged species because of forces of electrostatic attraction.
Periodic Trend in Ionization Energy
◆ removing a core e– costs MUCH more energy than removing a valence e–
◆ Valence electrons are most easily lost during ionization, and are gained, lost, or shared during chemical reactions.
Electron Affinity
Electron Affinity
* increase*decrease
electron affinity – change in energy that occurs when an electron is added to an isolated gas phase atom.
X (g) + e– ! X– (g)
* increase means becomes larger, negative value ∴ more favorable for anion formation;
decrease means becomes smaller, negative value ∴ less energy released and less favorable for anion formation
Periodic Trend in Electron Affinity
note: where electron affinity values are > 0, anion formation is very unfavorable; alkaline earth metals & the noble gases