Chapter 8
Electron Configuration and Periodicity
8–1
John A. Schreifels
Chemistry 211
Chapter 8-1
Overview
• Electron Structure of Atoms
–
–
–
–
Electron spin and the Pauli Exclusion Principle.
Aufbau Principle and the Periodic Table
Electron Configuration
Orbital Diagram of atoms; Hund’s Rule
• Periodicity of the Elements
– Mendelev’s periodic table predicted undiscovered
elements.
– Periodic Properties
– Periodicity and the main group elements.
John A. Schreifels
Chemistry 211
Chapter 8-2
8–2
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.
8–3
John A. Schreifels
Chemistry 211
Chapter 8-3
Diagonal Rule for Build-up Rule
1s
2s
3s
4s
5s
6s
2p
3p
4p
5p
6p
3d
4d 4f
5d 5f
6d 6f
• The periodic table can also be used to determine
the electron configuration of an element.
John A. Schreifels
Chemistry 211
Chapter 8-4
8–4
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 Hund’s rule
• Hund’s 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 [Ar]4s23d104p3.
John A. Schreifels
Chemistry 211
Chapter 8-5
8–5
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.
John A. Schreifels
Chemistry 211
Chapter 8-6
8–6
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.
8–7
John A. Schreifels
Chemistry 211
Chapter 8-7
Anomalous Electron Configurations
• A few exceptions to the Aufbau principles
exist. Stable configuration:
– half-filled d shell:
• Cr has [Ar]4s13d5;
• Mo has [Kr] 5s14d5
– filled d subshell:
• Cu has [Ar]4s13d10
• Ag has [Kr]5s14d10.
• Au has [Xe]6s14f145d10
• Exceptions occur with larger elements where
orbital energies are similar.
John A. Schreifels
Chemistry 211
Chapter 8-8
8–8
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.
8–9
John A. Schreifels
Chemistry 211
Chapter 8-9
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. [Ar]4s13d94p1 is an excited-state electron
configuration for Cu.
8–10
John A. Schreifels
Chemistry 211
Chapter 8-10
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.
8–11
• Discovery of elements with correct properties.
John A. Schreifels
Chemistry 211
Chapter 8-11
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
8–12
John A. Schreifels
Chemistry 211
Chapter 8-12
Atomic Radius
• 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.
John A. Schreifels
Chemistry 211
Fig. 8.17 Atomic Radii for Main
Group Elements
8–13
Chapter 8-13
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+,
8–14
Ca2+.
John A. Schreifels
Chemistry 211
Chapter 8-14
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.
•
•
–
Group 6A elements.
•
•
•
Before ionization ns2np1.
After ionization is ns2. Higher energy  smaller radius.
Before ionization ns2np4.
After ionization ns2np3 where each p electron in different orbital (Hund’s rule).
Electron-electron repulsion by two electrons in same orbital
increases the energy (lowers EI).
John A. Schreifels
Chemistry 211
8–15
Chapter 8-15
Ionization Energy: Periodic table
Fig. 8.18 Ionization Energy vs atomic #
8–16
John A. Schreifels
Chemistry 211
Chapter 8-16
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).
8–17
John A. Schreifels
Chemistry 211
Chapter 8-17
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.
John A. Schreifels
Chemistry 211
8–18
Chapter 8-18
Fig. 8.2 Stern-Gerlach Experiment
• Hydrogen atoms split into two beams when passed
through magnetic field. Beams correspond to spin on
atom.
8–19
Return to slide 6
John A. Schreifels
Chemistry 211
Chapter 8-19