Organization of the Periodic Table


PERIODS
Columns of the periodic table
Atoms of elements in the same
group have the same # of valence
electrons and therefore behave
similarly
•Rows of the periodic table
•All elements in a period have
their valence electrons in the
same energy level.
1.
2.
3.
4.
H, He, C, Li
K, Ca, As, Br
He, Ne, Kr, Ar
B, Al, Ge, Sn
0%
0%
0%
0%
1.
2.
3.
4.
H, He, C, Li
K, Ca, As, Br
He, Ne, Kr, Ar
B, Al, Ge, Sn
0%
0%
0%
0%


Valence Electrons are electrons in the
outermost energy level.
- s or p electrons only (even when d and f
electrons are present they are not in the
outermost energy level)



Electron Dot Diagrams show the valence
electrons of an element.
Draw the electron dot diagrams for the
following:
Mg
N
F
1.
2.
3.
4.
1
2
7
8
0%
0%
0%
0%
1.
2.
3.
4.
1
2
7
8
0%
0%
0%
0%
1.
2.
3.
4.
5.
1
2
7
8
Unable to tell
0%
0%
0%
0%
0%


Think of a ball of an onion. What happens
with each layer?
As you go down the periodic table, the energy
levels increase and the size of the radius of
the atom increases. (each energy level is like
another layer of the onion)
1.
2.
3.
Na
K
Unable to be
determined
0%
0%
0%
1.
2.
3.
Br
Cl
Unable to be
determined
0%
0%
0%
•
•
•
As you move from left to right in the periodic
table, what happens to the number of
protons in the nucleus?
What effect do these protons have on the
electrons?
What effect do the electrons have on each
other?
•
Electron Shielding (or Screening) – These
inner electrons shield the valence electrons
from receiving the entire attractive nuclear
charge because they repel the valence
electrons.



Within a period, as you go from left to right,
the positive nuclear charge increases, and
attracts the electrons more strongly.
As the electrons are more attracted to the
nucleus, the atomic radius decreases.
Summary: as you go from left to right, the
atomic radius generally decreases.
1.
2.
3.
Li
Be
Unable to be
determined
0%
0%
0%
1.
2.
3.
Si
Ar
Unable to be
determined
0%
0%
0%
1.
2.
3.
Be
Mg
Unable to be
determined
0%
0%
0%
1.
2.
3.
C
Si
Unable to be
determined
0%
0%
0%


The octet rule states that all atoms attempt to
become stable by having a full valence
electron shell (generally 8 electrons, hence
octet rule).
Atoms will gain, lose, or share electrons in
order to attain this stability.
1.
2.
3.
4.
Alkali metals
Transition metals
Halogens
Noble Gases
0%
0%
0%
0%
1.
2.
3.
4.
5.
6.
7.
Alkali metals
Transition metals
Halogens
Alkaline earth
metals
Both alkali metals
and halogens
Noble gases
Both alkali metals
and alkaline earth
metals
14%
14%
14%
14%
14%
14%
14%



Electrons are held in atoms by their attraction
to the positively charged nucleus.
To remove an electron requires energy.
Ionization energy is the energy required to
remove the least tightly bound (or outermost)
electron from an atom.

Compare Li and K.

How many valence electrons?

What is the relative size of the atoms?

Which has a higher ionization energy?


As you go down a group, the ionization
energy decreases because it takes less energy
to remove an electron.
The least tightly bound electrons are further
from the positive nucleus, and can therefore
be removed more easily.
1.
2.
3.
4.
Halogens
Alkali Metals
Alkaline Earth
Metals
All the same
0%
0%
0%
0%
1.
2.
3.
4.
Li
Be
F
Ne
0%
0%
0%
0%
1.
2.
3.
4.
Na
Mg
S
Ar
0%
0%
0%
0%

Ionization energy generally decreases as you
go down a group and from right to left in a
period.
1.
2.
3.
4.
5.
1
3
4
6
8
0%
0%
0%
0%
0%
1.
2.
3.
4.
5.
6.
7.
K
Na
Li
Be
Cs
Ca
F
0%
0%
0%
0%
0%
0%
0%
1.
2.
3.
4.
1
2
3
None of the above
0%
0%
0%
0%
1.
2.
3.
4.
5.
6.
7.
K
Na
Li
Be
Cs
Ca
F
0%
0%
0%
0%
0%
0%
0%
1.
2.
3.
4.
1
4
7
14
0%
0%
0%
0%

Second ionization is the energy required to
remove a second electron.
Ex. Sodium has a lower (first) ionization
energy than Magnesium but Mg has a lower
second ionization energy than Na.

Why?

1.
2.
3.
4.
K
Ca
Ga
C
0%
0%
0%
0%


Electronegativity is the tendency of an
element to attract electrons in a bond.
Therefore, elements that want to gain
electrons will have higher electronegativity.
attractive
1.
2.
3.
4.
Alkali Metals
Transition Metals
Halogens
Noble Gases
0%
0%
0%
0%
1.
2.
3.
4.
5.
6.
7.
K
Na
Li
Be
Cs
Ca
F
0%
0%
0%
0%
0%
0%
0%
1.
2.
3.
4.
Alkali metals
Alkaline earth
metals
Halogens
Transition Metals
0%
0%
0%
0%
The pattern for
increasing electronegativity
(except for noble gases).
The pattern for increasing
ionization energy.
The pattern for increasing
atomic radius.

When an atom loses or gains electrons, it
gains a charge.
An ion is a charged atom.

A positive ion is called a cation. (ex. Na+)

A negative ion is called an anion. (ex. S2-)


Cations (positive) – lose valence electrons in
the outermost energy level.
◦ They lose an energy level so they get smaller.

Anions (negative) – gain valence electrons but
their Zeff does not change.
◦ They get bigger because of having more electrons
and the same Zeff
1.
2.
3.
Mg
Mg+
Mg2+
0%
0%
0%
1.
2.
3.
O
OO2-
0%
0%
0%
1.
2.
3.
Ca
Ca+
Ca2+
0%
0%
0%
1.
2.
3.
S
SS2-
0%
0%
0%


Groups 1-13 become cations by losing
electrons because of their low ionization
energy.
Their positive charge corresponds to their
group number.


Groups 15-17 become anions by gaining
electrons because of their high
electronegativity.
Their negative charge corresponds to how
many electrons they must gain to have the
same electron configuration as a noble gas.
1.
2.
3.
4.
Ionization energy
Atomic radius
Thermal capacity
Electronegativity
0%
0%
0%
0%
1.
2.
3.
4.
Aluminum
Oxygen
Magnesium
Lithium
0%
0%
0%
0%
1.
2.
3.
4.
Atomic radius
Ionization Energy
Atomic Mass
Ionic Charge
0%
0%
0%
0%

4 Different kinds of metals
◦ Alkali metals: soft, shiny and very reactive
 Group 1: not found in nature as elements
◦ Alkaline earth-metals: less reactive
 Group 2: have two valence electrons
◦ Transition Metals: many uses
 Groups 3-12

3 Different kinds of metals
◦ Noble Gases: mostly non-reactive, very stable
 Group 8: He, Ne, Ar, Kr, Xe, Rn
◦ Halogens: very reactive, gain one electron to form a
stable compound
 Group 7: F, Cl, Br, I









Seven elements are called diatomic and never
exist alone in nature.
Have No Fear Of Ice Cold Beans?????
H2
N2
F2
O2
I2
Cl2
Br2
1.
2.
3.
4.
5.
Alkali metals
Alkaline earth
metals
Transition metals
Halogens
Noble gases
0%
0%
0%
0%
0%
1.
2.
3.
4.
5.
Alkali metals
Alkaline earth
metals
Transition metals
Halogens
Noble gases
0%
0%
0%
0%
0%
1.
2.
3.
4.
5.
Alkali metals
Alkaline earth
metals
Transition metals
Halogens
Noble gases
0%
0%
0%
0%
0%
1.
2.
3.
4.
5.
Alkali metals
Alkaline earth
metals
Transition metals
Halogens
Noble gases
0%
0%
0%
0%
0%
1.
2.
3.
4.
5.
Alkali metals
Alkaline earth
metals
Transition metals
Halogens
Noble gases
0%
0%
0%
0%
0%

Metals
◦ Shiny
◦ Solids
◦ Stretched and
Shaped
◦ Conductors of
heat and
electricity

Nonmetals
◦ Solids, liquids or
gases
◦ Solids – dull and
brittle
◦ Poor conductors of
heat and electricity
****Semiconductors /
Metalloids – exhibit
properties of both
metals and
nonmetals

Metals
◦ On the left hand side of
the zigzag line (except
for Hydrogen –
exception)
- Metalloids or Semimetals
- Touching zigzag line
(Except for Al)
- Exhibit properties of
both metals and
nonmetals

Nonmetals
◦ On the right hand
side of the zig zag
line (plus
Hydrogen)
The difference in electronegativity between atoms A and B is given by: