Transcript Periodicity PPT

PERIODICITY
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.
Discovery of elements with correct properties.
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
TRENDS IN THE PERIODIC TABLE USE OF WHAT WE HAVE JUST
LEARNED
A)Atomic radius - a number of physical
properties of elements are related to the
size of an atom, but with our probability
picture where does an atom end?
The atomic radius is ½ the distance
between the 2 nuclei of the adjacent atoms.
Atomic Radius
 Atomic
radii
actually decrease
across a row in the
periodic table. Due
to an increase in
the effective
nuclear charge.
Fig. 8.17 Atomic Radii for
Main Group Elements (s,p)
(The more protons you have,
the harder they pull the e- to
them.)
 Within
each group
(vertical column),
the atomic radius
tends to increase
with the period
number.
Chapter 8-6
Atomic Radius 2
 If
positively charged the radius decreases.
(lost e-)
 If the charge is negative the radius increases
(gained e-).
 When substances have the same number of
electrons (isoelectronic), the radius will depend
upon which has the largest number of protons.
C) Atomic radius, in general, decreases as
we move from left to right in a row of the
periodic table.
D) Atomic radius increases from top to
bottom in a family or group.
E) These 2 trends are the result of 3
influences on size.
1) As the number of the principal energy
level “n” increases, the size increases,
they extend further from the nucleus
and the covalent radius increases.
(bigger outer level, higher floor in the
motel—the larger the radius)
2) As nuclear charge (number of protons)
increases across a row, the positive charge
on the nucleus increases to electrons are
pulled closer to the nucleus…radius get
smaller.
3) The shielding effect is The attraction for
electrons in the outermost shell by the
nucleus is shielded by electrons in lower
energy levels. *As you gain electrons it
becomes harder to pull in the farther ones.
a) The smaller size of atoms going across a row
can be attributed to minimum shielding.
b) Electrons in the same shell are attracted
more strongly as the nuclear charge (# of
protons) increases, because the shielding
effect remains the same.
c) If the shielding effect remains the same, the
Effective Nuclear Charge increases.
d) The ENC is the positive charge that an
electron experiences from the nucleus and is
equal to the nuclear charge minus the
number of shielding electrons.
For example: Li has 3protons in the nucleus, 2e in the 1s orbital (shielding)
and 1e in the 2s orbital. ENC = 3 - 2 = 1. The outermost electron "feels" a
net attraction by the inside of +1.
Ionization Energy/Electron Affinity
a) Ionization Energy = energy necessary to
remove an electron---is always
endothermic and positive +.
b) M(g) + h  M+ + e.
b) Electron Affinity = energy change upon
the addition of an electron can be either
endothermic or exothermic depending on the
element (for a gaseous atom)
A(g) + e A-(g)
*An exothermic Energy = - value
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.

Sign of the ionization energy is
always positive, for example, it
requires energy for ionization to
occur.

Ionization
Energy:
Periodic
Fig. 8.18 Ionization Energy vs atomic #
table
IONIZATION ENERGY
A(g) + energy  A+ + 1e
A(g) + energy A+ + 1e H = + kJ/mol
2) THIS IS A VERY IMPORTANT
CONCEPT because the chemical properties
of any atom are determined by the
configuration of an atom's valence electrons,
those electrons in the outermost shell.
6) The trend across from left to right is
accounted for by a) the increasing nuclear
charge.
W
H
Y?
IONIZATION ENERGY
The electrons in the outermost shell are
more strongly bound to the nucleus due to
the increasing effective nuclear charge.
a) as we go across a row of the periodic
table energy is larger nuclear charge
becomes larger as the number of protons
in the nucleus of the atom becomes
larger.
IONIZATION ENERGY
b) With an electron already in the orbital
there is repulsion between the two in the
same orbital and it comes out with less
energy input.
c) The trend from top to bottom of a
column shows a decrease in the FIE which
corresponds to an increase in the atomic
radius.
9) The 2nd, 3rd, and 4th ionization energies
are those required to remove the 2nd, 3rd,
and 4th electrons.
Electron Affinity
1) Electron affinity is the energy change
which occurs when an electron is accepted
by an atom in the gaseous state.
A(g) + e A-(g)
2) In contrast to ionization energy, what do
we observe on the following graph of EA's?
Electron Affinity
The greater the negative value of the
electron affinity, the greater the tendency of
an atom to accept an electron.
d) A +value indicates that energy must be
absorbed for an atom to gain an electron.
e) left to right on the periodic chart,
general increasing tendency to form
negative ions. However, there are more
exceptions than with Ionization Energy.
Electron Affinity
Electron affinities generally become smaller as
we go down a column of the periodic table for
two reasons.
• First, the electron being added to the atom is
placed in larger orbitals, where it spends
less time near the nucleus of the atom.
• Second, the number of electrons on an atom
increases as we go down a column, so the
force of repulsion between the electron
being added and the electrons already
present on a neutral atom becomes larger.
Electron Affinity
Electron Affinity
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.
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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.
c) What should you be able to do as a
result of this???
I should be able to give you a list of
elements and you should be able to put
them in order of size from smallest to
largest by just looking at their positions
on the chart.
You should be able to tell me the reasons
why they are smaller or larger.