Creativity at its best!
 Consider the periodic table. People could
have invented various tables in the 17th
century but Mendeleev created the periodic
table of elements. Ponder on your interests
and hobbies in life. If you could create your
own periodic table, what would it be in
relations to? How would it be organized?
Brainstorm this and write about it.
Development of
the Periodic
Table
Sizes of Atoms
and Ions
Ionization
Energy
Electron
Configurations
E.Q.: Can I compare and contrast trends in the chemical and
physical properties of elements and their placement on the
Periodic Table (SC4b)?

Metals (+)  Metalloids (+)  Nonmetals (-)

Developed by Dimitri Mendeleev and Julius
Meyer.

Mendeleev is given most credit.

Among the first discovered were iron (Fe),
copper (Cu), gold (Au), and silver (Ag).

Initially arranged in order of mass number.

Henry Moseley later developed the concept
of atomic numbers.

The table was then arranged in order of
increasing atomic number.

Majority of elements exist in nature only as
compounds.
Except for H,
elements left of
the zigzag line
are metals.
To the right of
the line we find
nonmetals,
including the
noble gases.
Some
elements
adjacent to
the line are
called
metalloids.
6

Because e- are negative, they’re attracted to
nuclei, which are _____________.

Coulomb’s Law – The strength of attraction
between two electric charges depend on...

Thus, the force of attraction between
electrons and nuclei depend on...

In many – e- atoms, each e- is attracted to the
nucleus and repelled by...

The charge resulting from positive charge of
the nucleus and the shielding effects of e- is
called the effective nuclear charge.
 The net electric charge of the nucleus after the
shielding effects of its surrounding e-

Effective nuclear charge increases as we
move across any period of the table.

Effective nuclear charge increases as we
move down a group...HOWEVER, not as
much as moving across a period.

Indicate which would you expect to
experience a greater effective nuclear charge:
a Neon atom or a sodium atom.



Creation of the periodic table...
Development of the modern periodic table...
Effective nuclear charge...
Study your notes!

Fill in the diagram below and tell me what
this diagram is demonstrating.

s: _____

p: _____ _____ _____

d: _____ _____ _____ _____ _____

f: _____ _____ _____ _____ _____ _____ ____
Development of
the Periodic
Table
Sizes of Atoms
and Ions
Ionization
Energy
Electron
Configurations
E.Q.: How do I use the periodic table to predict periodic trends
including atomic radii, ionic radii, ionization energy, and
electronegativity (SC4a)?

Remember this concept?

Explain how effective nuclear charge
increases and decreases.

Which would you expect to experience a
greater effective nuclear charge, helium or
magnesium?

One of the most important properties of
atoms/ions – SIZE!!!

What determines size of atoms/ions?
 Atomic radius (plural: radii)
 Distance from nucleus to valence electrons.
 Measured in Angstroms (1.0 Å = 10-10m)

Decreases across table; increases down table.

Look at the trends...

Loss/gain of an electron...

s: _____

p: _____ _____ _____

d: _____ _____ _____ _____ _____

f: _____ _____ _____ _____ _____ _____ ____

What determines size of ions?
 Formed when atoms lose/gain e Less e- are pulled more closely to nucleus.
 More e- - larger radius, less e- - smaller radius.


Positive ions are smaller than their parent
atoms!
Negative ions are larger than their parent
atoms!




Atomic radii________ as you move from left
to right across a period.
Atomic radii ________ as you move down a
group.
Ionic radii __________ as you gain e-.
Ionic radii __________ as you lose e-.

Arrange these atoms and ions in order from
smallest to largest: B, F, and O.



Sizes of Atoms and Ions...
Periodic Trends in Atomic Radii...
Periodic Trends in Ionic Radii...
Study Your Notes

Based on what you know, opine why atoms
have different sizes.

Observe the equation below. Opine what’s
going on in the chemical equation:
Development of
the Periodic
Table
Sizes of Atoms
and Ions
Ionization
Energy
Electron
Configurations
E.Q.: How do I use the periodic table to predict periodic trends
including atomic radii, ionic radii, ionization energy, and
electronegativity (SC4a)?

Minimum energy required to remove an efrom an atom.

First ionization energy (I1) – Minimum energy
required to remove the first e- from an atom.

Second ionization energy (I2) – Minimum
energy required to remove the second efrom an atom. So on and so fourth!

The greater the ionization energy, the more
difficult it is to remove an electron!
 Ionization energy values increase
as successive electrons are
removed.
 I1 < I 2 < I3

Referring to the periodic table, rank the
following atoms in order of smallest to
largest ionization energy: Ne, Na, P, Ar, and
K.


Ionization Energy...
Periodic trends in Ionization Energy...
Study your notes!

Fill in the electron orbital diagram below

s: ____

p: ____ ____ ____

d: ____ ____ ____ ____ ____

f: ____ ____ ____ ____ ____ ____ ____
Development of
the Periodic
Table
Sizes of Atoms
and Ions
Ionization
Energy
Electron
Configurations
E.Q.: How do I use the periodic table to predict periodic trends
in electron configuration?

Electron configuration - Distribution of
electrons among orbitals in atoms.
 Represented in two ways:
1. spdf notation
2. Orbital box notation

Coefficient – Shell

s,p,d,f – designates
orbital.

Superscript – How
many e- are in that
orbital.
Represents electrons within orbitals and orbitals
within shells. The arrows’ directions represent
electron spins; opposite spins are paired.
N:




e- ordinarily occupy orbitals of the lowest
energy available.
No two e- in the same atom may have all four
quantum numbers alike.
Pauli exclusion principle: one atomic orbital
can accommodate only two e-, and these emust have opposing spins.
e- in half-filled orbitals have parallel spins
(same direction).
The electron configuration
of Si ends with 3s2 3p2
The electron
configuration of Rh
ends with 5s2 4d7











4d
5p
5s
3d
4p
4s
3p
3s
2p
2s
1s

Try study guide questions (#5).
2
2
6
2
6
2
1
 Sc: 1s 2s 2p 3s 3p 4s 3d
 Ti:

Illustrate the electron configuration for
hydrogen and lithium.

Indicate the element with the following
electron configuration: 1s22s22p4

Electron configuration...
Study Your Notes!