Download I. Periodic Trends - Golden Valley High School

Chapter
6.1-6.3
Periodic Table Lecture
Do members of the same family,
generally behave the same?
Yes
The Periodic Table
The Alkali Metals
Lithium, Sodium,
Potassium, Rubidium,
Cesium, and francium
very reactive
1 valence electron
s1 sublevel is filled
The Transition Metals
metals with
atomic numbers
21-112 highest s
& d sublevels
have electrons
Alkali Earth Metals
Beryllium,
Magnesium,
Calcium,
Strontium, Barium,
and Radium
2 valence electrons
s2 sublevel is filled
Metalloids
Like metals & nonmetals
Boron, Silicon,
Germanium,
Arsenic,
Antimony,
Tellurium, Polonium
Nonmetals
• Consists of Carbon,
Nitrogen, Oxygen,
Phosphors, Sulfur,
Selenium
• poor conductors
Noble Gases
• consists of Helium, neon, Argon, Krypton,
Xenon, Radon
• unreactive stable inert because they
already have 8 valence electrons
of heat and
electricity compared
to metals
• dull and brittle
Other Metals
Halogens
• Consists of Fluorine,
Chlorine, Bromine, Iodine,
Astatine
• nonmetals
• have 7 valence electrons
• very reactive
Inner Transition metals
• want one more electron • consists of elements with
(octet rule)
atomic numbers 58 through
71 and 90 through 103
•F sublevels partially filled
• the Lanthanide Series has
atomic numbers 58 -71 and the
Actinide Series has atomic
numbers 90-103
Define the term inert gas?
noble gas –unreactive & stable
Representative Elements #1 – Group IA-VIIA
outer s & p orb partially filled
1
Alkali
Metals
ns
Group 1A
Group 2A Alkaline Earth
ns2
Group 3A Nonmetals/Metalloids
ns2 np1
Group 4A
Group 5A
Group 6A
Group 7A Halogens
Group 0
Noble Gases
8 or 18
ns2 np2
ns2 np3
ns2 np4
ns2 np5
ns2 np6
Representative Elements #1
Lewis dot structure
Na 1s2 2s2 2p6 1s2
Group B
Transition Metals
Group 58-71
Lanthanides
Group 90-103 Actinides
Filling the “d”
orbital
Filling the “4f”
orbital
Filling the “5f”
orbital
A. Ionic Size
metals
(group 1A-3A)
lose electrons to become stable cation
non-metal (group
1A =
5A-7A)
gain electrons to become stable anions.
Loses 1 e-
2A = Loses 2 e3A =
Loses 3 e-
5A = Gains 3 e6A =
Gains 2 e-
7A = Gains 1 e-
7
P
E
R
I
O
D
S
!
!
v
A Family is a Group living between Colum
http://www.privatehand.com/flas
h/elements.html

Periodic Table Song by Tom Lehrer above
End of Lecture 6.1
Next Lecture 6.2
Who designed the 1st periodic table in 1869?
Dmitri Mendeleev
grouped w/ similar chemical and physical properties
& ordered by atomic mass.Ex: Co Ni
Ar
K
Te
I
http://www.youtube.com/watch?v=y7dmRtlXaYQ
http://www.youtube.com/watch?v=zUDDiWtFtEM
Lecture 6.3
Periodic Trends
I. Periodic Trends - Atomic Size
Atomic Radii:
Measured as 1/2 distance
between nuclei 2 atoms
Nucleus
Distance
between
nuclei
Atomic
Radius
Atomic Size generally INCREASES as you
move down a group on the periodic table.
Why?
down a group
increases # of energy levels
Example:
Ca atom larger than
a Mg atom. Why?
An energy level
is added!
Atomic Size generally DECREASES
across a row on the periodic table.
Why?
adding more p+ pulls in extra electrons
RELATIVE ELECTRONEGATIVITY, IONIZATION
ENERGY, RADII, SHIELDING ETC…
ElectroHydrogen
negativities:has the smallest atomic radius
Hydrogen Oxygen
Carbon
Sodium
2.1 3.5
2.5 0.9
Na < ionization
energy than O
because less protons
pull.
B. Ionization Energy
energy needed to pull an
electron away from an atom.
B. Ionization Energy
Example :
Na
Na+1 + e-
Ionization energy decreases
as you move down a group.
Why?
increased distance from protons reduces
attractive force
Period Trend:
Ionization energy generally increases as you
move across a period.
Why?
nuclear charge increases (more protons)
which increases attractive forces
energy required to remove the 1st outermost
electron is 1st ionization energy.
What is the second ionization energy?
Which is harder to remove?
Why?
What happens to the shielding of the
nucleus as you move across a period?
Remains constant
Why?
•ONLY adding electrons, NOT a new energy level.
What happens to the shielding of the nucleus
as you move down a group?
Increases
Why?
another energy level that shields
those valence electrons.
Ca+ions –
smaller than the original atom
Why?
When electrons lost,
a whole energy level lost
decreases radius.
Negative anions grow larger
Why?
there are more e- than p+
(increased electron repulsion),
N
atom
N-3
anion
from group 5A to the right,
anions gradually decrease in size
Why?
groups 6A &7A only gain 1 or 2 eHave Same # of e-, but increased # of p+
N-3
anion
O-2
anion
F-1
anion
.
B. Electronegativity
Noble gases no electronegative #
Why?
inert / don’t form compounds.
Can’t force a noble gas to take an
electron – they have s2 p6
3. Period Trends
left to right electronegativity increases.
Why?
High ionization energy = high electronegativity
Resists
electron loss
Attracts
electrons
Fluorine is the most electronegative!
4. Group Trends
Electronegativity decreases down a group.
Why?
Increased energy levels and shielding
Cs has the lowest electronegativity