Ions & periodicity

1 Periodic Trend of Ionic Charges 18

2 13 14 15 16 17

3 4 5 6 7 8 9 10 11 12
+1

+/- 4
tend to have
+2

-2
+3

-1
-3
more than one option

+3

+ 3 or + 4

Trends when atoms form Ions
Empirical Observation
“when forming ions, elements tend to lose or gain
electrons to be more like the nearest noble gas”


Metals tend to lose e-’s


Metals tend to lose e-’s

Nonmetals tend to gain e-’s

Atomic Radius
Definition: Half the distance
across a neutral atom

Atomic Radius

Radius decreases across a period

Atomic Radius

 The electrons are pulled in tighter as the # of
protons increases

Atomic Radius

protons increases
Radius increases down a group

Atomic Radius

protons increases
Radius increases down a group
 Each row on the periodic table adds a “shell” or
energy level to the atom. This makes atoms larger
with each added shell, just like an onion

Atomic Radius of Diatomic
Molecules

Ionic Radius Definition: Half the distance across an Ion

Cations

Anions


 Positively charged ions formed when
Cations an atom of a metal loses one or
more electrons

Anions


more electrons
 Smaller than the neutral atom

Anions


more electrons

 Negatively charged ions formed
when nonmetal atoms gain one
Anions
or more electrons


more electrons

 Negatively charged ions formed
when nonmetal atoms gain one
Anions
or more electrons
 Larger than the neutral atom

Trends in Atomic and Ionic Size

Metals

Metals
Group 1

Metals
Group 1

Li

152

Na

186

K

227

Metals
Group 1

e

Li

152

Na

186

K

227

Metals
Group 1

e

Li Li+

152 60

Na

186

K

227

Metals
Group 1

e

Li Li+

152 60

e

Na

186

K

227

Metals
Group 1

e

Li Li+

152 60

e
Na+
Na
95
186

K

227

Metals
Group 1

e

Li Li+

152 60

e
Na+
Na
95
186
e

K

227

Metals
Group 1

e

Li Li+

152 60

e
Na+
Na
95
186
e
K+
K

133
227

Metals
Group 1 Group 13

e

Li Li+

152 60
e
e e
e
Na+ Al3+
Na Al

95 143 50
186
e
K+
K

133
227

Metals
Group 1 Group 13

e

Li Li+

152 60
e
e e
e
Na+ Al3+
Na Al

95 143 50
186
e
K+
K

133
227
Cations are smaller than parent atoms

Metals Nonmetals
Group 1 Group 13

e

Li Li+

152 60
e
e e
e
Na+ Al3+
Na Al

95 143 50
186
e
K+
K

133
227

Metals Nonmetals
Group 1 Group 13 Group 17

e

Li Li+

152 60
e
e e
e
Na+ Al3+
Na Al

95 143 50
186
e
K+
K

133
227

Metals Nonmetals

e

Li Li+
F
152 60 64
e
e e
e
Na+ Al3+
Na Al Cl
95 143 50 99
186
e
K+
K Br

114
133
227

Metals Nonmetals

e
e
Li Li+
F F-
152 60 64
e 136
e e
e
Na+ Al3+
Na Al Cl
95 143 50 99
186
e
K+
K Br

114
133
227

Metals Nonmetals

e
e
Li Li+
F F-
152 60 64
e 136
e e e
e
Na+ Al3+
Na Al Cl Cl-
95 143 50 99
186 181
e
K+
K Br

114
133
227

Metals Nonmetals

e
e
Li Li+
F F-
152 60 64
e 136
e e e
e
Na+ Al3+
Na Al Cl Cl-
95 143 50 99
186 181
e e
K+
K Br Br-

114
133 195
227

Metals Nonmetals

e
e
Li Li+
F F-
152 60 64
e 136
e e e
e
Na+ Al3+
Na Al Cl Cl-
95 143 50 99
186 181
e e
K+
K Br Br-

114
133 195
227
Cations are smaller than parent atoms Anions are larger than parent atoms

Periodic Properties
• How does the size of an atom change
when electrons are added or
removed?

Periodic Properties
removed?

As an Atom loses
1 or more electrons
(becomes positive),
it loses an orbit or
energy level and
therefore, its radius
decreases.

Periodic Properties
removed?

As an Atom gains
1 or more electrons
(negative), it fills its
outermost energy
level, therefore, its
radius increases.

Electron Shielding
CORE electrons are
closer to the nucleus
and therefore have a
stronger attraction to
the positive nucleus.
This “shields”
valence electrons
from feeling as
strong of a pull from
the positive nucleus.

Effective Nuclear Charge
• In atoms with many
electrons, electrons are
both attracted to the
nucleus and repelled by
other electrons.

other electrons.
• Opposite charges attract.
Like charges repel.

other electrons.
• Opposite charges attract.
Like charges repel.
• The nuclear charge that
an electron experiences
depends on how close it
is to the nucleus and how
many other electrons are
around it.

Isoelectronic Series

EXAMPLES:
1. Na+ = 10 e Ne = 10 e F- = 10 e

2. P-3 = 18 e Ca+2 = 18 e Ar = 18 e

Isoelectronic Series

•Atoms and Ions with the
SAME # of electrons

EXAMPLES:
1. Na+ = 10 e Ne = 10 e F- = 10 e

2. P-3 = 18 e Ca+2 = 18 e Ar = 18 e

Electronegativity
Definition: The ability of an atom to attract or hold
onto electrons

Electronegativity
onto electrons

o Electronegativity INCREASES across a
period

Electronegativity
onto electrons

period
o WHY? As the atomic radius decreases,
electrons get closer to the atom’s nucleus

Electronegativity
onto electrons

period
o Electronegativity tends to DECREASE
down a group

Electronegativity
onto electrons

period
o Electronegativity tends to DECREASE
down a group
o WHY? As the atomic radius increases,
electrons are farther from the atom’s
nucleus

Periodic Trend:
Electronegativity

Ionization Energy
• Amount of energy required to
remove an electron from an atom.

Ionization Energy
First ionization energy is that
energy required to remove first
electron.

Ionization Energy
First ionization energy is that
energy required to remove first
electron.
Second ionization energy is that
energy required to remove second
electron, etc.

Trends in First Ionization Energies

Trends in First Ionization Energies
• Down a GROUP,
ionization energy
decreases
• For atoms in the same
group, Zeff is essentially
the same, but the valence
electrons are farther from
the nucleus.
• Across a PERIOD,
ionization energy
increases

Ions & periodicity

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Editor's Notes