John A. Schreifels

1
Chapter 8

• Electron Configuration and Periodicity

2
Overview

• Electron Structure of Atoms
• Electron spin and the Pauli Exclusion Principle.
• Aufbau Principle and the Periodic Table
• Electron Configuration
• Orbital Diagram of atoms Hunds Rule
• Periodicity of the Elements
• Mendelevs periodic table predicted undiscovered
  elements.
• Periodic Properties
• Periodicity and the main group elements.

3
Orbitals in Multielectron Atoms

• Electrons are attracted to the nucleus but also
  repelled by each other.
• Repulsion from other electrons reduces the
  attraction to the nucleus by a small amount
  giving rise to an effective nuclear charge
• Effective nuclear charge the net nuclear charge
  felt by an electron after shielding from other
  electrons in the atom is taken into account.
  Zeff Zact ? Zshield.

4
Diagonal Rule for Build-up Rule

• The periodic table can also be used to determine
  the electron configuration of an element.

5
Electron configurations of multielectron atoms
(Aufbau principle)

• Electron configuration determined since electrons
  tend to be in lowest energy orbitals.
• The Aufbau principle guides us in the filling of
  orbitals
• Fill lowest energies first.
• Maximum of two electrons with opposite spins in
  each orbital.
• Degenerate orbitals (orbitals with same energy)
  follow Hunds rule
• Hunds rule If two or more orbitals have the
  same energy, fill each orbital with one electron
  before pairing electrons.
• E.g. Determine the electron configurations of H
  and He
• H ? 1s1 ?
• He ? 1s2 ??
• E.g. 2 Determine the electron configuration of
  the second row elements.
• E.g.3 Determine the electron configuration of
  the 4th row elements.
• Shorthand electron configuration of arsenic is
  Ar4s23d104p3.

6
Magnetic Properties

• Although an electron behaves like a tiny magnet,
  two electrons that are opposite in spin cancel
  each other. Only atoms with unpaired electrons
  exhibit magnetic susceptibility (see Fig. 8.2).
• A paramagnetic substance is one that is weakly
  attracted by a magnetic field, usually the result
  of unpaired electrons.
• A diamagnetic substance is not attracted by a
  magnetic field generally because it has only
  paired electrons.

7
Periodic Table and Electron Configurations

• Build-up order given by position on periodic
  table row by row.
• Elements in same column will have the same outer
  shell electron configuration.

8
Anomalous Electron Configurations

• A few exceptions to the Aufbau principles exist.
  Stable configuration
• half-filled d shell
• Cr has Ar4s13d5
• Mo has Kr 5s14d5
• filled d subshell
• Cu has Ar4s13d10
• Ag has Kr5s14d10.
• Au has Xe6s14f145d10
• Exceptions occur with larger elements where
  orbital energies are similar.

9
Electron Configuration of Excited States Ions

• Metals form cations by losing e? nonmetals
  become anions by gaining e?.
• Both adopt inert gas electron configuration.
• E.g. The alkali metals will lose a single
  electron to become M. The electron
  configuration is He, Ne, Ar, Kr, and
  Xe for Li, Na, K, Rb respectively.

10
ISOELECTRONIC SUBSTANCES and EXCITED STATES

• Substances with the same number of electrons are
  isoelectronic ions.
• Isoelectronic ions (or molecules) ions (or
  molecules) with the same number of valence
  electrons.
• Isoelectronic substances P3?, S2?, Cl?, Ar, K,
  Ca2.
• The electron configation of an element in an
  excited state will have an electron in a
  high-energy state
• E.g. Ar4s13d94p1 is an excited-state electron
  configuration for Cu.

11
Development of the Periodic Table

• Mendeleev developed periodic table to group
  elements in terms of chemical properties.
• Alkali metals develop 1 charge, alkaline earth
  metals 2
• Nonmetals usually develop negative charge (?1 for
  halides, ?2 for group 6A, etc.)
• Blank spots where elements should be were
  observed.
• Discovery of elements with correct properties.

12
Periodic Properties

• Periodic law elements arranged by atomic number
  gives physical and chemical properties varying
  periodically.
• We will study the following periodic trends
• Atomic radii
• Ionization energy
• Electron affinity

13
Atomic Radius
Fig. 8.17 Atomic Radii for Main Group Elements

• Atomic radii actually decrease across a row in
  the periodic table. Due to an increase in the
  effective nuclear charge.
• Within each group (vertical column), the atomic
  radius tends to increase with the period number.

14
Atomic Radius 2

• If positively charged the radius decreases while
  if the charge is negatively the radius increases
  (relative to the atom).
• When substances have the same number of electrons
  (isoelectronic), the radius will depend upon
  which has the largest number of protons.
• E.g. Predict which of the following substances
  has the largest radius P3?, S2?, Cl?, Ar, K,
  Ca2.

15
IONIZATION ENERGY

• Ionization energy, Ei minimum energy required to
  remove an electron from the ground state of atom
  (molecule) in the gas phase. M(g) h? ? M e?.
• Ei related to electron configuration. Higher
  energies stable ground states.
• Sign of the ionization energy is always positive,
  i.e. it requires energy for ionization to occur.
  
• The ionization energy is inversely proportional
  to the radius and directly related to Zeff.
• Exceptions to trend
• B, Al, Ga, etc. their ionization energies are
  slightly less than the ionization energy of the
  element preceding them in their period.
• Before ionization ns2np1.
• After ionization is ns2. Higher energy ? smaller
  radius.
• Group 6A elements.
• Before ionization ns2np4.
• After ionization ns2np3 where each p electron in
  different orbital (Hunds rule).
• Electron-electron repulsion by two electrons in
  same orbital increases the energy (lowers EI).

16
Ionization Energy Periodic table
Fig. 8.18 Ionization Energy vs atomic
17
HIGHER IONIZATION ENERGIES

• The energies for the subsequent loss of more
  electrons are increasingly higher. For the
  second ionization reaction written as
• M(g) h? ? M2 e? Ei2.
• Large increases in the ionization energies vary
  in a zig-zag way across the periodic table.
• States with higher ionization energies have
  1s22s22p6 (stable).

18
ELECTRON AFFINITY

• Electron Affinity, Eea, is the energy change that
  occurs when an isolated atom in the gas phase
  gains an electron.
• E.g. Cl e? ? Cl? Eea ?348.6 kJ/mol
• Energy is often released during the process.
• Magnitude of released energy indicates the
  tendency of the atom to gain an electron.
• From the data in the table the halogens clearly
  have a strong tendency to become negatively
  charged
• Inert gases and group I II elements have a very
  small Eea.

19
Fig. 8.2 Stern-Gerlach Experiment

• Hydrogen atoms split into two beams when passed
  through magnetic field. Beams correspond to spin
  on atom.

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