Transcript Atomic Radius

Periodic Trends
Periodic Trends
 Atomic Radius
 Ionic Radius
 Ionization Energy
 Electron Affinity
 Electronegativity
Atomic Size or Radius
 The distance from the center of the nucleus
to the edge of the atom.
© 1998 LOGAL
Where do you start measuring
from?
The electron cloud doesn’t have
a definite edge.
They get around this by
measuring more than 1 atom at a
time.

Atomic Size
}
Radius
Atomic
Radius = half the distance between
two nuclei of a diatomic molecule.
Trends in Atomic Size
Influenced
by three factors:
1. Energy Level
• Higher energy level is further
away.
2. Charge on nucleus
• More charge pulls electrons in
closer.
3. Shielding effect (blocking effect?)
Shielding
 The
electron on the
outermost energy
level has to look
through all the other
energy levels to see
the nucleus.
 Second electron has
same shielding, if it
is in the same period
Group trends
As
we go down
a group...
each atom has
another energy
level,
so the atoms
get bigger.
H
Li
Na
K
Rb
Periodic Trends



As you go across a period, the radius gets
smaller.
Electrons are adding to the SAME energy level.
The effect of adding 1 proton to the nucleus
overpowers the effect of adding 1 electron to the
cloud. Nuclear charge increases, and pulls
electrons in.
Therefore, the atoms get SMALLER.
Na
Mg
Al
Si
P
S Cl Ar
Atomic Radius
Atomic Radius
Increases to the LEFT and DOWN
1
2
3
4
5
6
7
Examples
Which atom has the larger radius?
Be or Ba
Ba
Ca or Br
Ca
Trends in Ionic Size
Cations
form by losing electrons.
Cations are smaller that the atom
they come from.
Metals form cations.
Cations of representative
elements have noble gas
configuration.
Ionic size
Anions
form by gaining
electrons.
Anions are bigger that the atom
they come from.
Nonmetals form anions.
Anions of representative
elements have noble gas
configuration.
Configuration of Ions
 Ions
always have noble gas
configuration.
 Na is: 1s22s22p63s1
 Forms a 1+ ion: 1s22s22p6
 Same configuration as neon.
 Metals form ions with the
configuration of the noble gas
before them - they lose electrons.
Configuration of Ions
Non-metals
form ions by gaining
electrons to achieve noble gas
configuration.
They end up with the
configuration of the noble gas
after them.
Examples
Which particle has the larger radius?
S or
2S
2S
Al or
3+
Al
Al
Ionization Energy
 The energy required to remove one electron from
a neutral atom in the gaseous state.
 Removing one electron makes a 1+ ion.
 The energy required to remove the first
electron is called the first ionization
energy.
Ionization Energy
First Ionization Energy
trend is opposite of atomic radius
Increases UP and to the RIGHT
1
2
3
4
5
6
7
Ionization Energy
 Down a group:
 Electrons are located in  Across a period:
higher energy levels,
 Electrons are located
farther from the nucleus,
in the SAME energy
so they are easier to
level. Nuclear
remove.
charge is increasing.
 Also, the “shielding
 Therefore, electrons
effect” is increasing
are harder to remove.
(more electrons from
lower energy levels
blocking nuclear charge),
making it even easier to
remove an electron.
Ionization Energy
Successive Ionization Energies
Large jump in I.E. occurs when a CORE
e- is removed.
Mg
Core e-
1st I.E.
736 kJ
2nd I.E.
1,445 kJ
3rd I.E.
7,730 kJ
Ionization Energy
Successive Ionization Energies
Large jump in I.E. occurs when a CORE
e- is removed.
Al
Core e-
1st I.E.
577 kJ
2nd I.E.
1,815 kJ
3rd I.E.
2,740 kJ
4th I.E.
11,600 kJ
Ionization Energy
 The 2nd I.E. is always higher than the 1st,
the 3rd I.E. is always higher than the 2nd,
and so on.
 Why? Removing each electron decreases
the “shielding effect” around the nucleus.
Nuclear charge increases, making it harder
to remove another electron.
Ionization Energy
 The noble gases have extremely high
ionization energies.
 Why? Due to the stability that comes with
filled s and p orbitals.
 Therefore, elements that have achieved
noble gas configurations (by gaining or
losing electrons) also have very high I.E.’s.
Examples
Which atom has the higher 1st I.E.?
N or Bi
N
Ba or Ne
Ne
Electron Affinity
 The energy given off when an electron is
gained by a neutral atom.
Electron Affinity
Electron Affinity
trend is same as ionization energy
Increases UP and to the RIGHT
1
2
3
4
5
6
7
Electron Affinity
 Energy is usually given off when an atom gains
an electron.
These values for E.A. are negative.
 Sometimes, an atom must be “forced” to accept
an electron through the addition of energy.
These values for E.A. are positive.
 Adding a second electron to an already negative
ion is always more difficult, so 2nd E.A.’s are
always positive!
Electron Affinity
 Down a group:
 Across a period:
 Electrons are located in
 Electrons are located
higher energy levels,
in the SAME energy
farther from the nucleus.
level. Nuclear charge
Nuclear charge is
is increasing.
decreased, “shielding
 Therefore, electrons
effect” is increased.
add more easily
 Therefore, electrons add
with greater difficulty
The group most likely to…
 LOSE electrons?
Alkali metals (group 1)
Lowest ionization energy!
 GAIN electrons?
Halogens (group 17)
Highest (most negative) electron affinity!
Electronegativity
 The ability of an atom in a chemical
compound to attract electrons.
 trend is same as electron affinity
Electronegativity
 Fluorine is the most electronegative
element!
 Fluorine is assigned an electronegativity
value of 4.0 on the Pauling Scale.
 All other elements are assigned values
relative to this.
Electronegativity
 In general, elements that tend to lose electrons to
form positive ions (metals) have lower e-neg’s.
 Elements that tend to gain electrons to form
negative ions (nonmetals) have higher e-neg’s.