Special Topics for SOL 2
rd
3 Power Point
Periodic Trends (Chap 14)
Shorthand Electron
Configurations
 Shorthand configurations are a useful tool.
 Let’s look at an example for Y, Z=39
 The electron configuration for yttrium is
1s22s22p63s23p64s23d104p65s24d1
 To do a shorthand configuration, we use the
noble gas preceding the element and we put
that in brackets (the bold and italics part)
 That’s Kr and then we also just write
whatever is left over.
 [Kr] 5s24d1
You Try…
 Do a shorthand configuration for
 Fe
 Br
 Rb
The Answers…
 Do a shorthand configuration for
 Fe = [Ar]4s23d6
 Br = [Ar]4s23d104p5
 Rb = [Kr]6s1
Objective B
http://www.rsc.org/chemsoc/visualelements/PAGES/data/intro_groupvii_data.html
 Notice that the halogens all have an ending
configuration of ns2np5. That means they have
7 valence electrons.
F
Br
At
[He]2s22p5
Cl
[Ar]3d104s2 4p5
I
[Xe]4f14 5d106s2 6p5
[Ne]3s23p5
[Kr]4d105s2 5p5
 Similarly, alkali metal have 1 valence electron.
Noble gases have 8, etc. All transition metals
have 2.
Objective B
 All of the transition metals have 2 valence electrons,
with 2 exceptions. “d” electrons are not valence
electrons. Why not?
 Transition metals are where the d orbitals are being
filled up. Here are the electron configurations for
all of them.
Sc
V
Mn
Co
Cu
[Ar]3d14s2
[Ar]3d34s2
[Ar]3d54s2
[Ar]3d74s2
[Ar]3d104s1
Ti
Cr
Fe
Ni
Zn
[Ar]3d24s2
[Ar]3d54s1
[Ar]3d64s2
[Ar]3d84s2
[Ar]3d104s2
Objective B
 Notice that Cr and Cu are “exceptions.”
 They both have 1 valence electron. They do
this because in the case of Cr, moving an
electron from the 4s level to the 3d level gives
us a half full set of d orbitals.
 That’s more stable than if Cr would have
followed the pattern, and ended with “4s23d4”
Cr
[Ar]3d54s1
Objective B
 Similarly, Cu has 1 electron in the
4s energy level and 10 in the 3d
level, because having a full set of d
electrons is also more stable.
Cu
[Ar]3d104s1
Objective B
 The “inner transition metals” are the
lanthanide and actinide series.
 That’s where the f electrons are filled up.
 That’s about all I’m going to say about that.
Objective C
 The periodic table allows you to predict trends in certain
properties.
 Get out a periodic table and put these trends as notes on your
periodic table.
 The first trend is Atomic
radius.
 Atomic radius is the size of the atom. It’s defined as ½ the
distance between two nuclei which are bonded together.
Objective C
 Ionic radius is another property
 It is the size of an ion. Ionic radius is fairly similar to
atomic radius.
 A positive ion is also called a CATION.
 A negative ion is also called an ANION.
 A cation is always smaller than the
from.
atom it is formed
 An anion is always larger than the
atom it is formed from.
Objective C
http://www.chem1.com/acad/webtext/atoms/atpt-images/ionic_radii.jpg
 Since cations lose electrons to form positive ions and
anions gain electrons to form negative ions, it
should make sense that they are SMALLER than the
atom.
Objective C
 Ionization energy is the amount of energy required
to remove an electron from a gaseous atom.
 The energy required to remove the first electron is
called the FIRST IONIZATION ENERGY.
 The energy required to remove the second electron
is the second ionization energy. And so on…
 Metals always have LOWER ionization energies
than nonmetals.
 That is because metals tend to lose electrons and
nonmetals tend to gain them.
Objective C
 It is VERY MUCH easier to remove a valence electron (an electron
in the highest energy level) than an “inner core” electron.
 The inner core electrons are ANY electrons which are not
VALENCE electrons.
Na = 1s22s22p63s1
White = inner core electrons and Blue = Valence electrons
Objective C
http://www.knowledgerush.com/wiki_image/8/87/LinusPauling.jpeg
 Electronegativity is
measured on a scale from
0.0 to 4.0.
 By definition, F is the
most electronegative
element at 4.0.
 Nonmetals have a high
electronegativity.
 Metals have a low
electronegativity.
Electronegativity
 Think of this as the “greediness” of an atom not only holding
on to it’s own electrons, but ALSO wanting to “steal”
electrons from other atoms.
The Trends
 Atomic Radius AND Ionic Radius increase as you go down
a group.
 Atomic Radius AND Ionic Radius decrease as you go from
left to right across a period.
 Electronegativity AND Ionization Energy decrease as you
go down a group.
 Electronegativity AND Ionization Energy increase as you
go from left to right across a period.
Note the trends are opposites. Draw some arrows on your
periodic table to help you remember the trends.
The End