Download L to R

Periodic Trend
• Horizontal rows are called periods
• Vertical columns are called groups
• There are 7 periods, 18 groups
Element Arrangement
• Elements in each row have different electron
energy shell level – Principle Energy Level
• Elements in each column of table have same
number of electrons in their outer energy level –
Valence Electron
• Group 1 or 1A– one valence electron
• Group 2 or 2A– two valence electron
• Group 13 or 3A– has three valence electron and
so on.
• Groups 3/ 3B to 12 / 12B have different number
of valence electron
Periodic Table
• Periodic Table
A. Metallic Character
• Metals
• Nonmetals
• Metalloids
– Metals- 80%, ___
left of stair-step line, excellent
conductors, lustrous, malleable, ductile, mostly solid
right of stair-step line, generally
– Nonmetals- ____
poor conductors, brittle, non-lustrous, some gases
borders stair-step line; intermediate
– Metalloids- ______
metal & nonmetal, semi conductor
» Stair-step line aka. “zig-zag”
• We can also look at the periodic table in
terms of electron configuration, the
arrangement of electrons in an atom
Take a look at Pg 119
• What is the electron configuration of C?
• What is the electron configuration of Pb?
– s-block elements: chemically reactive metals,
1
• alkali metals- group ___
2
• alkaline earth metals- group __
– d-Block elements:
• transition metals- group ____,
3-12
– p-Block elements: part of main block
17 most reactive nonmetal
• halogens- group ___,
18 innert, nonreactive gases
• noble gases- group ___,
– f-Block elements: To save space the lanthanides
and actinides are set off below the main portion of
the Table.
rare earth________
elements
• inner transition metals-aka ____-____
• Lanthanide Very similar physical & chemical properties
Still placed in the 6th and 7th periods, respectively (even
• Actinide
though 4f and 5f)
Periodic Trend
• As you go from left to right, the number of
protons increase.
• As you go from left to right, the number of
electrons in the same energy level
increase.
• As you go from top to bottom, more shells
are added.
• As you go from top to bottom, electrons
are farther away.
Effective Nuclear Charge
• Nucleus is positive.
• Positive charge of the nucleus attracts
electrons
• Electrons closer are more effectively drawn to
nucleus
• Electrons farther away are less drawn to
nucleus
• Attractive force from the nucleus is canceled
out by inner electrons. Outermost shell feel
less attraction – Electron Shielding
Atomic Radius
• Hard to determine- electron clouds
• Bond radius- the length that is half the distance
between the nuclei of two bonded atoms. Figure
19.
• L to R, more protons  more pulling  closer to
the nucleus  smaller the bond radius 
smaller atomic radius
• T to B, more electron shells  electrons are
farther away  inefficient pulling  bigger the
bond radius  larger atomic radius
Atomic size increases, (shielding
constant across a period)
Ionic size increases
Ionization Energy
• The energy needed to remove an electron
from the shell
• L to R, more protons  stronger pulling 
tighter the atom  more energy needed to
remove an electron  I.E increases
• T to B, more electron shells  electron
shielding  farther electrons feel less
attraction  less energy to remove an
electron  I.E decreases
Ionization energy, Electronegativity,
Electron affinity INCREASE
Electronegativity
• A measure ability to attract electrons
• L to R, more protons  more pulling from
nucleus  easier to attract another
electron  increase in electronegativity
• T to B, more electron shells  farther
away  less effective pulling  decrease
in electronegativity.
Atomic Size
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Radius
• Atomic Radius = half the distance between
two nuclei of a diatomic molecule