Transcript Chapter 12 The Periodic Table

Chemical Periodicity
Trends in the periodic table
Atomic Size
How do you measure the size of an
atom?
 The electron cloud doesn’t have a
definite edge.
 Can get around this by measuring
covalent atomic radius.

Atomic Size
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Radius
Atomic Radius = half the distance between
two nuclei of a diatomic molecule.
Atomic size is influenced by two factors:
 Energy
Level
–more occupied levels = bigger atom
 Charge
on nucleus
– More charge pulls electrons in closer
Group trends

As we go down a
group electrons
are added to
higher energy
levels so the
atoms get bigger.
H
Li
Na
K
Rb
Periodic Trends
As you go across a period the radius
gets smaller.
 Same energy level, but protons pull
electrons closer to nucleus.

Na
Mg
Al
Si
P
S Cl Ar
Trends in Atomic Radius
Questions:
Of the following elements, which has the largest
atomic radius? Why?
a) Si, Mg, S
Mg – same energy level, smallest nuclear charge
b) Al, Na, Cl
Na – same energy level, smallest nuclear charge
c) Li, Cs
Cs – higher occupied energy levels
Ionic Size
Ionic Size
Cations




Positive ions - form by losing electrons.
Metals form cations
Cations of representative elements have noble gas
configuration.
Smaller than the atom they come from because of
increased attraction by nucleus for fewer
remaining electrons
+
Ionic size
Anions




Negative ions - form by gaining electrons.
Nonmetals form anions.
Anions of representative elements have noble gas
configuration.
Larger than the atom they come from, because
nuclear attraction is less for an increased number
of electrons.
-
Group trends

Ions get bigger as you
go down (adding
energy levels)
Li+1
Na+1
K+1
Rb+1
Cs+1
Periodic Trends

Across the period nuclear charge
increases so both cations and anions get
smaller from left to right.
Li+1
B+3
Be+2
C+4
N-3
O-2
F-1
Questions:
1. Of the following ions, which ones should have
the larger radius? Why?
a) Na+ or Cs+
Cs+
It has more occupied energy levels
b) Br- or K+
Br- Anions are larger than cations
2. The Mg2+ and Na+ ions have ten electrons
surrounding the nucleus. Which ion would you
expect to have the smaller radius? Why?
Mg2+ Greater nuclear charge
Ionization Energy
Ionization Energy
The amount of energy required to
completely remove an electron from a
gaseous atom (how hard it is to pull an e- off
an atom)
 1st IE = removing 1 e-, 2nd IE=removing 2 e
Na(g)
Na+ + e-
Shielding

The electron on
the outside energy
level is shielded
from the nucleus
by the inner
electrons
Group trends
As
you go down a group first IE
decreases because the electron is
further away (more shielding)
Periodic trends


All the atoms in the same period have the same
energy level (same shielding).
As you go from left to right, nuclear charge
increases so IE generally increases.
Questions:
1. Which element in the following sets has the
lowest ionization energy and why?
a) B, C, F
B – same energy level, smallest nuclear charge
b) K, Na, Li
K – electron farther away, more shielding
Electron Affinity
Electron Affinity


The energy given off when an electron is
added to an atom
how much an atom ‘wants’ an electron
F(g) + e -
F -(g)
Electron Affinity
Group trends

Generally decreases as we go down a group
because shielding increases
Periodic trends

Increases from left to right as atoms become
smaller with greater nuclear charge
Questions:
1. Of the following elements, which ones should
have the higher electron affinity? Why?
a) Se or Te
Se – smaller atom
b) Calcium or Chromium
Chromium – greater nuclear charge
Electronegativity
Electronegativity


The tendency for an atom to attract
electrons to itself when it is chemically
combined (BONDED) with another
element.
Big electronegativity means it pulls the
electron towards itself.
Group Trends



The further down a group the farther the
electron is away from the nucleus and the
more electrons an atom has.
More willing to share = low
electronegativity
So as you go down a group
electronegativity decreases
Periodic Trends



As we go from left to right across the table,
electronegativity increases, because nuclear
charge is increasing and electrons are held
in more strongly
Metals have low electronegativity
Non-metals have high electronegativities
(they win the electron tug-of-war)
Questions:
1. Which element would you expect to have the
highest electronegativity? Why?
F
smallest nonmetal
2. Put the following elements in order of increasing
electronegativity: Na, P, Cl
Na, P, Cl