```Atomic Radius
• A measure of the size of an atom – how
close it lies to its neighbor.
• Half the distance between 2 adjacent,
identical atoms (crystal or molecule).
• The radius for any atom can vary
depending upon the substance it’s in.
principal energy level which is larger and
more electrons internally to shield the
valence electrons.
• Radii decrease across a period – same
principal energy level with a larger positive
charge in the nucleus.
Which has the largest atomic radius: carbon,
fluorine, beryllium or lithium?
•Lithium
•has the same outer energy level as the others
•but its nucleus has the smallest positive charge
Ions
• Form when atoms gain or lose electrons.
• If an atom gains electrons, it becomes a
negative ion.
• If an atom loses electrons, it becomes a
positive ion.
• Metals are more likely to lose electrons
while nonmetals are more likely to gain
electrons.
• When an atom forms an ion, its radius will
always change size.
• Negative ions become larger due to a
greater electrostatic repulsion among the
valence electrons.
• Positive ions become smaller due to a
smaller electrostatic repulsion among
valence electrons and the fact the outer
orbital often becomes empty in the
process.
• As you move across a period, the positive
ions (on the left) and the negative ions (on
the right) decrease in size just like the
• But negative ions (on the right) will always
be bigger than the positive ions (on the
left).
• As you move down a group, ionic radii
increase – both positive & negative.
1. For each of the following pairs, predict
which atom is larger.
a. Mg, Sr
d. Ge, Br
b. Sr, Sn
e. Cr, W
c. Ge, Sn
2. For each of the following pairs, predict
which atom or ion is larger.
a. Mg, Mg2+
c. Ca2 +, Ba2 +
e. Na +, Al3 +
b. S, S2 d. Cl -, I -
1. For each of the following pairs, predict
which atom is larger.
a. Mg, Sr
b. Sr, Sn
c. Ge, Sn
d. Ge, Br
e. Cr, W
2. For each of the following pairs, predict
which atom or ion is larger.
a. Mg, Mg2+
b. S, S2 c. Ca2 +, Ba2 +
d. Cl -, I e. Na+, Al3 +
Ionization Energy
• The energy required to remove an electron
from a gaseous atom.
• The energy required to remove the first
electron is called the first ionization energy
and so forth.
• Group 1A has lowest ion. energy w/ group
8A having the highest as you move across
a period.
Ionization Energies
• Successive ionization energies ALWAYS require
more energy because there are fewer electrons
left to repel one another.
• At some point this will take a huge jump
because you’ve reached inner level electrons.
• First ionization energies increase as you move
across a period because the nuclear charge
increases so the nucleus holds the electrons
more tightly.
• However they decrease as you move down a
family due to greater atomic size, so electrons
are further away from the nucleus.
For each of the following pairs, predict which
atom has the higher first ionization
energy.
a. Mg, Na
b. S, O
c. Ca, Ba
d. Cl, I
e. Na, Al
f. Se, Br
a. Mg, Na
b. S, O
c. Ca, Ba
d. Cl, I
e. Na, Al
f. Se, Br
Octet Rule
• Atoms tend to gain, lose or share in order
to acquire a full set of 8 valence electrons.
• “Eight is great, but two will do.”
• For whatever reason, filled s and p orbitals
at any energy level is extremely stable.
Electronegativity
• Indicates the relative ability of an atom to attract
or gain electrons in a chemical bond.
• Calculated using several factors & always have
a value of 3.98 Paulings or less.
• Greatest electronegativity = fluorine (3.98)
• Least electronegativity = franicum (.70)
• It has the same trend of ionization energy which
is losing electrons.
• Decreases as you move down a group and
increases as you move left to right across the
periodic table.
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