CHAPTER 7: PERIODIC
PROPERTIES OF THE ELEMENTS
Exploration of important properties of elements
across a row or down a column.
JANUARY 2ND, 2013
 Do

Now:
Take out your homework, compare
your answers to a neighbor.
FROM MENDELEEV TO MOSELEY

1869: Mendeleev
arranged elements by atomic mass
 Noted similarities, forced to leave holes in PT


1913: Moseley
Developed concept of atomic numbers
 Bomabrded elements with high-energy electrons



Unique frequency
Correctly identified atomic number as number of
protons
EFFECTIVE NUCLEAR CHARGE



Where do many of the properties of atoms
originate?
How does Coulomb’s law relate to effective
nuclear charge?
As we get further from the nucleus, how is our
effective nuclear charge effected?
EFFECTIVE NUCLEAR CHARGE CONT.


Why is analyzing the effective nuclear charge
precisely problematic?
How do we estimate these interactions?
EFFECTIVE NUCLEAR CHARGE
Zeff = Z − S
where:
Z = atomic number
S = a screening constant
(usually close to the
number of inner
electrons)
TRENDING



How does effective nuclear charge change across
a period?
Why?
How does effective nuclear charge change down a
period?
Covalent atomic radius(Bonding radius): is half the
distance between two atoms covalently bonded.
Van der Waals Radius (Non-bonding radius): is the
radius of an atom when it is bonded to another atom
by van der Waals forces.
ATOM SIZE

Had students read 7.4 in the textbook and then
had class discussion
JANUARY 3RD, 2012

Do Now:

In terms of effective nuclear charge, explain the
atomic radius trend both across a period and down a
group.

Arrange the following isoelectronic series in
decreasing atomic size:

K+, Cl-, Ca2+, S2-
ARE YOU A BRAINIAC?
IONIZATION ENERGY



Define Ionization Energy.
How do we differentiate between 1st and 2nd
ionization energies?
How is the difficulty of removing an electron
correlated with ionization energy?
TRENDING!
 I1


< I2 < I3 , Why?
How is the ionization energy affected w/ removal
of electrons from an inner shell?
Page 259 (table)
TRENDING CONT.



How does ionization energy trend from left to
right across a period?
How about down a group?
Generally, s + p block show larger range in
values for ionization energy. The d block
increases minimally. The f, very little.
PRACTICE



Which has a higher third ionization energy, Ca or
S? Why?
Why does Boron have a lower ionization energy
than Beryllium?
Why is it easier to remove a 2p electron from
oxygen than from nitrogen?
JANUARY 4TH, 2012
DO NOW :
 Arrange the following atoms in order of
increasing ionization energy:


Neon, Sodium, Phosphorus, Argon, and Potassium.
Arrange the following atoms in order of
decreasing ionization energy:

Boron, Aluminum, Carbon, and Silicon
SO HOW DOES IONIZATION WORK?

When electrons are removed from an atom, they
are removed from occupied orbital with highest
principle quantum number.


Try removing 2 electrons from Iron.


IE: Lithium.
How about a third?
Try removing 2 electrons from Tin.

How about 4?
PRACTICE

Write the electron configurations for the
following:
Ga3+
 Cr3+
 Br
AWKWARD?



Why is using ionization energy “awkward” when
talking about forming an anion?
Instead, we use the term “electron affinity”!
Energy is released when adding an electron. How
would we represent this numerically?
HOW DOES THIS ALL TIE TOGETHER?


How can we differentiate between ionization
energy and electron affinity?
What does the electron affinity of a noble gas look
like?
PREDICTIONS!


Predict which group would have the most
negative electron affinities. Why?
What is the relationship between the value of the
first ionization energy of Cl-(g) and the electron
affinity of Cl(g)?
BROADLY SPEAKING…


Elements organized as metals, nonmetals, metalloids.
Metals:






Luster; various colors although mostly silver
Solids malleable and ductile
Good conductors
Most metal oxides are solids that are basic
Tend to form cations in aqueous solution
Non Metals:





No luster, various colors
Solids usually brittle; some hard, some soft.
Poor conductors
Most nonmetal oxides are molecular substances that form
acid solutions
Tend to form anions in aqueous solution
METALLIC CHARACTER
Where is the “most” metallic metal located on the
periodic table?
 What is the general trend for metallic character
based on this information?

METALS



Which metals is a liquid at room temperature?
How can we use first ionization energy to
characterize metals vs. nonmetals?
Why are most metal oxides basic?
METAL OXIDES

Most metal oxides are basic.
Metal Oxides + Water  Metal Hydroxide
IE: Na2O (s) + H2O (l)  2 NaOH (aq)

Due to oxide ion, which reacts with water.
IE: O2- (aq) + H2O (l)  2 OH- (aq)

Insoluble metal oxides react with acids to form
salt plus water.
IE: NiO (s) + 2 HNO3 (aq)  Ni(NO3)2 (aq) + H2O (l)
PRACTICE

Would you expect scandium oxide to be a solid,
liquid, or gas at room temperature?


Write the balanced chemical equation for the reaction
of scandium oxide with nitric acid.
Write the balanced chemical equation for the
reaction between copper (II) oxide and sulfuric
acid.
NON METALS

How can non-metals exist at room temperature?

Why are most non-metal oxides acidic?
Nonmetal oxide + water  acid

Most nonmetal oxides dissolve in basic solutions
to form salt and water.
PRACTICE

Write the balanced chemical equation for the
reaction of solid selenium dioxide with the
following:
Water
 Aqueous sodium hydroxide


Write the balanced chemical equation for the
reaction of solid tetraphosphorus hexoxide with
water.