Download Section 5.3

Section 5.3
Electron Configuration
And
Periodic Properties
Trends
• Trends are approximations of
overall patterns. This section
discusses trends that we see in
elements in the periodic table.
Remember that exceptions can, and
do, occur.
Explaining trends…
• Properties such as the size of an atom (atomic
radius), the energy required to remove an electron
from an atom (ionization energy), and the energy
associated with the addition of an electron to a
gaseous atom (electron affinity) can be
understood in terms of the electron
configuration of the atom and a competition
between electron-nucleus attraction and
electron-electron repulsion.
Explaining trends…
Explaining trends…
Explaining trends…
Atomic Radii (see p. 140)
• The size of any atom is
defined by the edge of its
orbital. But since this
edge is fuzzy the atomic
radius is defined as onehalf the distance between
the nuclei of identical
atoms that are bonded
together.
Period Trends of Atomic Radii
• Trend: Smaller atoms across a
period (from left to right, see
p. 141).
• Caused by the increasing
positive charge of the nucleus
(increased electron-nucleus
attraction).
• Somewhat offset by the
repulsion among the number
of electrons in the same outer
energy level.
Group Trends of Atomic Radii
• In general, the atomic radii
increase down a group.
• n increases and so does the
distance from the nucleus
• Core electrons shield outer
electrons
• attractive force b/w nucleus
and outer electrons is less
• Trend does not always hold
(see Al to Ga, it decreases).
• Difference is small, for our
purposes we will ignore it.
Ionization Energies
• Electrons from an element can
be removed if enough energy
is supplied.
• Shown by:
A + E → A+ + e –
An ion is an atom or group of
bonded atoms that has a
positive or negative charge.
Ionization
• Any process that results in the formation of an
ion.
Ionization Energy (see p. 143)
• The energy required to
remove one electron
from a neutral atom of
an element (often called
the first ionization
energy).
• To remove a second e –
is called the 2nd
ionization energy, etc..
Ionization Energy Period Trends (see p. 144)
• In general,
ionization energies
of the main-group
elements increase
across each
period.
• Caused by
increasing nuclear
charge  increased
attraction b/w
nucleus and
electrons.
Combination
Removing Electrons From Positive Ions
• See p. 145
• 2nd, 3rd, 4th, etc. ionization energies are successively
higher. Notice the pattern at the stairstep.
• Due to the remaining electrons feeling an increasingly
stronger effective nuclear charge (the nuclear charge
minus the electron shielding).
Electron Affinity (p. 147)
• The energy change that occurs when an
electron is acquired by a neutral atom.
Exothermic Reaction
• Most elements
release energy
when they gain an
electron. The
numerical value is
negative.
A + e- → A- + E
Endothermic Reaction
• Some elements must be
“forced” to gain an electron
by the addition of energy:
• A + e - + E → A• The energy value is
positive. Such ions are
unstable, and will lose the
electron spontaneously.
Period Trends
• Increase in negativity from left to right, as electrons
add to the p sublevel. Not true between Groups 14
and 15.
14: ns2 np2
15: ns2 np3
Group Trends – Electron Affinity
• Electrons add with greater difficulty down a
group (more positive values). Mainly due to
the increase in atomic radius.
Cations – positive ions
• Loss of e- always leads
to a decrease in size
because removal leads
to a smaller electron
cloud.
• The remaining e- are
drawn closer to the
nucleus by the
unbalanced positive
charge.
Anions – negative ions
• Adding electrons leads
to larger ionic size
because the e- are not
drawn as strongly to the
nucleus and the e- cloud
spreads out due to
increasing repulsion
between the e-.
Group Trends
• In general, because of the increase in atomic radii, both the
anions and cations of a group are larger as one goes down a
group. Exceptions occur in some of the p-block groups as
elements change from non-metals to metals across the
staircase. Why?
Period
Trends
• To the left of the staircase, metals form cations
(+ charge).
• All metals tend to lose the electrons from the s
orbitals first.
• To the right and above the staircase, nonmetals tend
to form anions (- charge)
Cations
• Across a period
cationic radii
decrease because
the electron cloud
shrinks due to the
increasing nuclear
charge.
Anions
• Anionic radii decrease across a period for
the same reason.
Valence Electrons
• Valence electrons are the electrons available to be
lost, gained, or shared in the formation of a chemical
compound.
Numbers of Valence Electrons
• For the main-group
elements:
Groups 1 and 2;
equal to the group
number.
Groups 13-18; equal
to the group
number minus 10.
Common Ionic Charge
• Atoms gain/lose electrons to achieve a noble
gas electron configuration.
Electronegativity
• A measure of the
ability of an
atom in a
chemical
compound to
attract electrons.
Fluorine Reference Point
• The most electronegative element was arbitrarily
assigned a value of four and all other element values
are calculated in relation to this value.
Period & Group Trends
• Electronegativities tend to increase across each period
and up each group, although there are exceptions.
Metallic Character
• Metallic character is the name given to the set of chemical properties
associated with elements that are metals. These chemicals properties
result from how readily metals lose their electrons.