The document discusses periodic trends in atomic properties including atomic radius, ionization energy, and electronegativity. It explains that as you move down a group in the periodic table, atomic radius increases while ionization energy and electronegativity decrease due to increased electron shielding. However, as you move across a period, atomic radius decreases while ionization energy and electronegativity increase due to the stronger effective nuclear charge pulling electrons in more tightly.

1. MORE PERIODIC TRENDS

2. Key Concepts
• Effective Nuclear Charge: The attractive force exerted by
the atoms positively charged nucleus (+) on an electron (-
).
• Holds the electron in orbit.
• Opposites attract.
• As you add protons (+) the Effective Nuclear Charge grows
larger.
• Electron Shielding: The reduction in the attractive force
(Effective Nuclear Charge) between a positively charged
nucleus and its outermost electrons due to the cancellation of
some of the positive charge by the negative charges of the
inner electrons.
• Each primary energy level added moves the valence electrons farther
from the nucleus.
• Each successive energy level also puts more inner shell electrons “in
the way” between the nucleus and valence (outer) shell.
• Both of which decrease the amount of attractive force the nucleus can
exert.

3. TREND I: ATOMIC RADIUS

4. Atomic Radius
• The exact size of an atom is
hard to determine.
• The volume the electrons
occupy is thought of as an
electron cloud, with no
clear-cut edge.
• One method for calculating
the size of an atom involves
calculating the bond
radius, which is half the
distance from center to
center of two like atoms
that are bonded together.

6. Atomic Radius Increases as You Move Down a Group
•As you proceed from one element down to
the next in a group, a new principal energy
level is added.
•The addition of another level of electrons
increases the size, or atomic radius, of an
atom.
•Effect is further magnified by electron
shielding.

7. Atomic Radius Decreases as You Move Across a Period
• As you move from left to right across a period, each
atom has one more proton and one more electron than
the atom before it has.
• All additional electrons go into the same principal
energy level—no electrons are being added to the inner
levels – changes in electron shielding are minimal.
• As the nuclear charge increases across a period, the
effective nuclear charge acting on the outer electrons
also increases.
 Pulls the electrons in tighter, creating a smaller
atom.

9. Atomic Radii

10. TREND II: IONIZATION ENERGY

11. Ionization
• We’ve previously looked only at neutral atoms
with equal numbers of protons and electrons.
• Ionization: The creation of an atom with a net
charge, by the removal or addition of electrons.
– Ions have more or less electrons than neutral atoms.
• Anion: An atom that has gained an electron and
taken on a net negative charge.
• Cation: An atom that has lost an electron and
taken on a net positive charge.

12. Ionization Energy
•The ionization energy is the energy required
to remove an electron from an atom or ion.
•The Higher the Ionization Energy the
harder it is to remove an electron.
A + ionization energy ® A + + e -
neutral atom ion electron
First ionization energy: +1
Second ionization energy : +2

13. Multiple Ionization Energies
• Multiple ionization energies: If you want to pull off
more than one electron from an atom, more
energy is required for each additional one you
want to grab.
• Once you’ve reached a noble gas configuration (for
example, once magnesium has lost two electrons
to become like neon), any further electrons you
pull off will require a huge amount of energy.
First ionization energy: +1
Second ionization energy : +2

14. Ionization Energy Decreases as You
Move Down a Group
• Each element has more occupied energy levels than the
one above it has.
• The outermost electrons are farthest from the
nucleus in elements near the bottom of a group.
• As you move down a group, each successive element
contains more electrons in the energy levels between
the nucleus and the outermost electrons- more
electron shielding.
• Therefore it gets easier to remove electrons as you go
down.

15. Ionization Energy Increases as
You Move Across a Period
• Ionization energy tends to increase as you move from
left to right across a period.
• From one element to the next in a period, the number
of protons and the number of electrons increase by
one each- increases the effective nuclear charge.
 A higher nuclear charge more strongly attracts the
outer electrons in the same energy level, but the
electron-shielding effect from inner-level electrons
remains the same.

18. Electron Affinity
• Basically the opposite of ionization energy.
• The amount of energy released or absorbed
when an atom accepts an electron giving it a
negative charge. For most elements, energy is
released when an atom adds an electron. This is
also the measure of an element to attract an
electron to form a negative ion.
• How much it wants to gain an electron.
• Electron affinity increases from left to right
and decreases from top to bottom in a group or
family.

19. TREND I: REVISITED

20. Ion Radii
• Negative Ions: Always larger than the neutral
atom. Gaining electrons.
• Positive Ions: Always smaller that the neutral
atom. Loss of outer shell electrons.

23. TREND III: ELECTRONEGATIVITY

24. Electronegativity
• Electronegativity: is a measure of an atom’s
attraction for another atom’s electrons when in a
chemical compound.
– It is an arbitrary scale that ranges from 0 to 4.
• The units of electronegativity are Paulings.
– The atom with the higher electronegativity will pull on
the electrons more strongly than the other atom will.
– Generally, metals are electron givers and have low
electronegativities.
– Nonmetals are electron takers and have high
electronegativities.
– What about the noble gases?
• They have no electronegativity at all because they are noble
gases, and don’t need any more electrons

26. Electronegativities of Some Elements
Element Pauling scale
F 4.0
Cl 3.0
O 3.5
N 3.0
S 2.5
C 2.5
H 2.1
Na 0.9
Cs 0.7

27. Electronegativity Decreases as You Move Down a Group
• Electronegativity values generally decrease as you
move down a group.
• As you move from higher to lower in a group, the
electronegativity decreases due to the increase in
separation and effects of electron shielding.
• Generally, if an atom doesn’t hold the electrons it
already has very strongly (low ionization energy), it
won’t want to grab electrons from other atoms.

28. Electronegativity Increases as You Move Across a Period
• Electronegativity usually increases as you move left to
right across a period.
• As you proceed across a period, each atom has one
more proton and one more electron—in the same
principal energy level—than the atom before it has.
 Therefore the effective nuclear charge increases
across a period, resulting in an increase in
electronegativity.
• The increase in electronegativity across a period is much
more dramatic than the decrease in electronegativity
down a group.

31. Nuclear charge increases
Shielding increases
Atomic radius increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
Shielding is constant
Summary
Atomic Radius decreases
Nuclear charge increases
Electronegativity increases
Ionization energy increases