Electron Configurations and Periodicity

1
Electron Configurations and Periodicity

• Chapter 8

2
Spin and the Pauli Exclusion Principle

• No two electrons in an atom can have the same
  four quantum numbers.
• An orbital can hold a maximum of two electrons.
• Electrons in an orbital must be spin-paired.

3
Orbital Energy Levels in Multi-Electron Systems
3d
4s
3p
3s
Energy
2p
2s
1s
4
Electron Configurations-The Aufbau Principle

• It is possible for us to designate the
  electronic configuration of any atom using a
  simple procedure.
• Fill from lowest energy to highest
• Fill all orbitals in subshell with one electron
  before pairing (Hunds rule)
• Separate core (closed shell) and valence (outer)
  electrons to create short hand

5
Order for Filling Atomic Subshells
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p
5d 5f 6s 6p 6d 6f
6
Electron Configurations

• Use the building-up principle to obtain the
  ground state configuration of phosphorus.

7
Electron Configurations

• Use the building-up principle to obtain the
  ground state configuration of phosphorus.

Phosphorus (Z 15)
1s22s22p63s23p3
8
Electron Configurations

• Write the orbital diagram for the ground state
  of nickel. The electron configuration is
  Ar3d84s2.

9
Electron Configurations

• Write the orbital diagram for the ground state
  of nickel. The electron configuration is
  Ar3d84s2.

Nickel (Z 28) Ar
4s
3d
10
An Electronic Structure Example

• What is the electron configuration of Li? 1s22s1
  or He2s1
• What about Ca? 1s22s22p63s23p64s2 or
  Ar4s2
• What about Cr? 1s22s22p63s23p64s13d5 or
  Ar4s13d5

11
Core vs. Valence Electrons

• 1s22s22p63s23p64s13d5 or Ar4s13d5

Core Electrons
Valence Electrons
12
Periodicity

• Mendeleev (1872) recognized similarities in the
  properties of the oxides
• RO
• R2O
• R2O3
• RO4

13
Physical Trends

• Metals
• metallic luster, ductile solids (except Hg), low
  ionization energy, conduct electricity and heat
• Nonmetals
• varying physical appearance, no luster, high
  electron affinity
• Metalloids
• mixture of properties, often look metallic but
  without conductivity or ductility

14
Determining Atomic Radii

• Measure diatomic bond lengths and divide by 2.
• These measurements are typically low due to
  electron overlap.

2r
15
Relative Atomic Radii

• Across any row in the Periodic Table
• Core electrons stay the same
• Nuclear charge increases
• Shielding decreases
• Within a period, atomic radii tend to decrease
  moving left to right

16
Ionization Potential

• The energy necessary to remove electrons from
  an atom.

First Ionization Potential
Al - e-
Al
Second Ionization Potential
Al - e-
Al2
Third Ionization Potential
Al2 - e-
Al3
17
Relative Ionization Energies

• Ionization energy decreases going down a group.
• Ionization energy tends to increase going across
  a row.
• Filled degeneracies tend to be higher in energy.

18
Electron Affinity

• The energy change for the process of adding an
  electron to a neutral atom in the gaseous state
  to form a negative ion.
• Commonly high for reactive non-metals

Cl (g) e- Cl- (g)
19
Ionization Energy vs. Atomic Radii

• Linear Correlation
• Why?
• SHIELDING

20
Consider Ionization of He/He
e-
2
- e-
e-
He
He 2372 kJ/mole
e-
2
- e-
He2 5248 kJ/mole
He
21
Effective Nuclear Charge

• Electron repulsions can be thought of as
  reducing the effective nuclear charge
• Zeff Zactual - effect of electron repulsion

22
Group 1a, Alkali Metals

• Largest atomic radii
• React violently with water to form H2
• Readily ionized to 1
• Metallic character, oxidized in air
• R2O in most cases

23
2a, Alkali Earth Metals

• Readily ionized to 2
• React with water to form H2
• Closed s shell configuration
• Metallic

24
Transition Metals

• May have several oxidation states
• Metallic
• Reactive with acids

25
Group 3a

• Metals (except for Boron)
• Several oxidation states (commonly 3)

26
Group 4a

• Form the most covalent compounds
• Oxidation numbers vary between and - 4

27
Group 5a

• Form anions generally(-1, -2, -3), though
  positive oxidation states are possible
• Form metals, metalloids, and non-metals

28
Group 6a

• Form -2 anions generally, though positive
  oxidation states are possible
• React vigorously with alkali and alkali earth
  metals
• Non-metals

29
Halogens

• Form monoanions
• High electronegativity (electron affinity)
• Diatomic gases
• Most reactive non-metals (F)

30
Noble Gases

• Minimal reactivity
• Monotomic gases
• Closed Shell

31
Ion Sizes

• Cations shrink -
  
• Anions increase -

Ba
Ba2
Ba
Cl
Cl-
32
Periodicity and Electron Configurations

• Which of the following is the smallest ion? and
  WHY?
• S-, Fe2, As3-, Ca2, Be

33
Ions Want to Form a Closed Shell Too!

• S- Ne3s23p4 1 electron Ne3s23p5
• This is equivalent to Cl
• Fe2 Ar4s23d6 - 2 electrons
  Ar4s23d4 Ar4s13d5
• This is equivalent to Cr
• Ca2 Ar4s2 - 2 electrons Ar
• Be He2s2 - 1 electron He2s1