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Chapter 8
Periodic Properties of the Elements
Development of the periodic table
• Dobreiner: grouped elements into triads: 3
elements with similar properties (ex. Barium,
calcium, strontium)
• Newlands: (about 50 years after Dobreiner)
organized elements into octaves, similar to music
• Mendeleev: arranged elements in order of
increasing MASS from left to right and placed
elements with similar properties in the same
columns
• Mendeleev’s table was great! Undiscovered
elements could be predicted! (eka-aluminum,
eka= one beyond)
• However, not every element fit this system…
• MOSELEY! Arranged by atomic number instead
of mass. This resolved much of the issue with
Mendeleev’s arrangement
• PERIODIC LAW! Stated that the elements could
be arranged with predicting patterns of behavior,
but not the WHY
Electron Configurations
• Quantum-mechanical theory describes the
behavior of electrons in atoms.
• Electron Configuration: shows the particular
orbitals that are occupied for that atom
• Ground State: lowest energy state
• Electron configuration considers four values:
Energy level, sublevel, orbital, and spin
Electron spin and the Pauli Exclusion
Principle
• Spin is a basic property of all electrons
• All electrons have the same amount of spin
• The orientation of the spin is quantized (spin
up or spin down)(ms)
• Pauli Exclusion Principle: No two electrons in
an atom can have the same four quantum
numbers!
• Spin is represented by a +1/2 or a -1/2
Sublevel Energy:
• E(s orbital) < E(p orbital)< E(d orbital) <E(f
orbital)
• ONLY WHEN THE 1s ORBITAL IS OCCUPIED!
• If 1s orbital is empty, all other orbitals are
called degenerate: same energy
Penetration and Shielding
• Shielding: repulsive effect of one electron on
another
– Think of lithium. What charge does lithium have?
WHY????
– 2 inner electrons SHIELD the total 3+ charge of the
nucleus
– Effective nuclear charge (Zeff) = 1+
• Penetration: the 2s orbital experiences more of
the nuclear charge due to more penetration,
SOOOO lower energy
• Because of penetration, sublevels of each energy
level are NOT degenerate for multi-electron
atoms
• In fourth and fifth energy levels, penetration
becomes so important that it results in 4s being
lower energy than 3d
• There can be variances in electron configuration,
especially with TRANSITION METALS
Rules for electron configuration
• Pauli exclusion principle: no two electrons will
have the same 4 quantum numbers
• Aufbau principle: aufbau = German for ‘build
up’. Electrons will fill lowest energy first
• Hund’s rule: electrons fill orbitals singly first
with parallel spins (think of kids on a bus)
Let’s try it!
• Pg 328, practice 8-1 and 8-2
• Pg. 359, 41-50
Electron Configurations, valence
electrons, and the periodic table
• Notice the periodic trend: As you move down a
group, the number of outermost electrons
remains the same.
• THIS IS THE KEY CONNECTION BETWEEN ATOMIC
STRUCTURE AND ATOMIC BEHAVIOR!
• Valence Electrons: for main group electrons:
outermost electrons in the principle energy
level. For transition elements: we also count
electrons in the outermost d electrons.
• Core electrons: all other electrons (electrons in
complete principle energy levels)
• The chemical properties of elements are
largely determined by the number of valence
electrons they contain.
• Why gain or lose? A full outermost energy
level is relatively low energy. The element
becomes inert. It cannot lower its energy any
further by reacting with other elements.
Periodic trends in the size of atoms
and effective nuclear charge
• The volume of an atom is taken up mostly by its
electrons. But, we can’t know EXACTLY where
these electrons are, so how do we determine size
of atoms?
• Nonbonding atomic radius: determined by
freezing a sample of an element and measuring
the distance between the center of two adjacent
atoms. The atoms are NOT bonded together,
just very, very close to one another.
• Also called van der Waals radius
• Another way to define the size of an atom is
called the bonding atomic radius or covalent
radius
• Nonmetals: one half the distance between
two atoms bonded together
• Metals: one half the distance between two of
thee atoms next to each other in a crystal of
the metal
General Trend
• As you move down a group, radius increases
• As you move left to right, it generally
decreases.
• WHY?
Atomic Radii and the transition
elements
• The radius stays roughly consistent across a
row. WHY????
• Outermost electrons stay roughly consistent
Let’s try it!
•
•
•
•
Pg 329, practice 8.3
Pg 332, practice 8.4 (both)
Pg 338, practice 8.5 (both
Pg 359, 51-55, 59-62
More Trends!
• Ionization energy: the energy required to
remove an electron from an atom or ion in the
gaseous state.
– Energy required to remove first electron? FIRST
IONIZATION ENERGY, and so on.
– Trends: decrease as you move down a
column/group. Increases left to right across a
period. WHY????
Exceptions in Trends in First Ionization
Energy
• Exceptions do occur!
• Boron has a smaller first ionization energy
than Beryllium. WHY????
– Jump to the p block
– In general, the trends are reliable and predictable
Electron Affinities and Metallic
Character
• These characteristics also exhibit periodic
trends!
• Electron Affinity: the energy change
associated with gaining an electron
– Usually negative because an atom usually releases
energy when it gains an electron
– There is not much of an observable trend down a
group
– Becomes more negative as you move left t right
across a group. WHY???
Metallic Character
• Conductivity, Malleability, ductility, shiny
• General Trend: As you move left to right
across a row, metallic character decreases
• As you move down a column, metallic
character increases
Chapter 8 TEST
• This really isn’t too Bad, SOOO
• 1. Have a 3 x 5 notecard if you want
• 2. Know the history of the development of the
periodic table (Mendeleev, etc.)
• 3. Be able to complete full electron configurations,
noble gas configurations, electron configuration
diagrams
• 4. Be able to summarize (radii, ionization energy,
metallic character, electron affinity)
• We will review/practice tomorrow and test on
Wednesday, so COME WITH QUESTIONS!!!!
Homework
• Pg 345: practice 8.8 and more practice 8.8
• Pg 351: practice 8.9 and more practice 8.9
• Pg 359: 71, 72, 73, 74, 75, 76, 77, 78