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The Periodic Table and Trends
Please have a periodic table out.
Use this one.
Dmitri Mendeleev
1834 – 1907
• Russian chemist and teacher
• given the elements he knew
about, he organized a
“Periodic Table” based on
increasing atomic mass (it’s
now atomic #)
• he even left empty spaces to
be filled in later
At the time the elements gallium and germanium were not
known. These are the blank spaces in his periodic table. He
predicted their discovery and estimated their properties.
Henry Moseley
1887 – 1915
• arranged the elements in
increasing atomic
numbers (Z)
– properties now recurred
periodically
Design of the Table
• Groups are the vertical columns.
– elements have similar, but not identical,
properties
• most important property is that
they have the same # of valence
electrons
• valence electrons- electrons in the highest
occupied energy level
• all elements have 1,2,3,4,5,6,7, or 8 valence
electrons
Lewis Dot-Diagrams/Structures
• valence electrons are represented as dots
around the chemical symbol for the element
Na
Cl
2
1
3
2
5
8
What two blocks will always be the highest occupied level?
Look, they
are following
my rule!
• B is 1s2 2s2 2p1;
– 2 is the outermost energy level
– it contains 3 valence electrons, 2 in
the 2s and 1 in the 2p
• Br is [Ar] 4s2 3d10 4p5
How many valence electrons are
present?
• Periods are the horizontal rows
– do NOT have similar properties
– however, there is a pattern to their properties
as you move across the table that is visible
when they react with other elements
Trends in the table
IB loves the alkali metals and
the halogens
• many trends are easier to understand if
you comprehend the following
• the ability of an atom to “hang on to” or
attract its valence electrons is the result
of two opposing forces
– the attraction between the electron and the
nucleus
– the repulsions between the electron in
question and all the other electrons in the
atom (often referred to the shielding effect)
– the net resulting force of these two is
referred to effective nuclear charge
This is a simple, yet very good picture. Do you understand it?
• ATOMIC RADII
– the distance from the nucleus to the outermost
electron
– cannot measure the same way as a simple circle
due to electrons are not in a fixed location
– therefore measure distance between two nuclei
and divide by two
– groups
• increases downwards as more levels are added
• more shielding
– periods across the periodic table
• radii decreases
– the number of protons in the nucleus
increases
McGraw » increases the strength of the positive
Hill
nucleus and pulls electrons in the given
video
level closer to it
» added electrons are not contributing to the
shielding effect because they are still in the
same level
H
Li
Na
K
Rb
– atoms tend to gain or loose electrons in order to
have the electron configuration of a noble gas
– trends
• across a period
– decreases at first when losing electrons (+ ion)
– then suddenly increases when gaining electrons
(- ion)
– then goes back to decreasing after just like
neutral atoms because of more protons pulling
in the outer level
• down a group (same as neutral atoms)
– increases as new levels are added
– more levels shielding
IONIC RADII
Looking at ions compared to their
parent atoms
• meaning does an atom become smaller
or larger as it gains or loses electrons?
– cations (+ ions) are smaller than the parent
atom
• have lost an electron (actually, lost an entire level!)
• therefore have fewer electrons than protons
+
Li
0.152 nm
Li forming a
cation
Li+
.078nm
• radii still increases downwards as more levels are
added on and shielding increases
– anions (- ions) are larger than parent atom
• have gained an electron to achieve noble gas
configuration
• effective nuclear charge has decreased since
same nucleus now holding on to more electrons
• plus, the added electron repels the existing
electrons farther apart (kind of “puffs it out”)
F 0.064 nm
9e- and 9p+
F- 0.133 nm
10 e- and 9 p+
– IONIZATION ENERGY
• the minimum energy (kJ mol-1) needed to
remove an electron from a neutral gaseous
atom in its ground state, leaving behind a
gaseous ion
– X(g)  X+(g) + e-
• first ionization energy- energy to remove
first electron
• second ionization energy- energy to remove
second electron
• third ionization energy- and so on…
don’t forget-- gaseous
• decreases down a group
– outer electrons are farther from the nucleus and
therefore easier to remove
– inner core electrons “shield” the valence electrons
from the pull of the positive nucleus and therefore
easier to remove
• increases across a period
– the nucleus is becoming stronger (effective
nuclear charge) and therefore the valence
electrons are pulled closer
• atomic radii is decreasing
• this makes it harder to remove a valence electron
since it is closer to the nucleus
– or another way to look at it… a stronger
nuclear charge acting on more contracted
orbitals
• ELECTRONEGATIVITY
– measures the attraction for a shared pair
of electrons in a bond
• Linus Pauling (1901 to 1994) came up with a
scale where a value of 4.0 is arbitrarily
given to the most electronegative element,
fluorine, and the other electronegativities
are scaled relative to this value.
• trends (same as ionization energy and for the same
reasons)
• as you go down a group electronegativity decreases
– the size of the atom increases
» the bonding pair of electrons (-) is increasingly
distant from the attraction of the nucleus (+)
» the valence electrons (-) are shielded because of core
electrons (-) interfering with the nucleus’ (+) hold on
valence electrons
H
Li
Na
K
Rb
• as you go across a period
– electronegativity increases
• the atoms become smaller as the effective
nuclear charge increases
– easier to attracts electrons as they will be
in a level closer to the nucleus moving
from L to R on the table
• next concept requires understanding of
concepts covered in later topics
• only need to know the trends, not the
reason why until later
– metals do what they do
– Van der Waals forces
– bonding
– MELTING POINT
• down group 1 (alkali metals)
– decreases as “sea of negative
electrons” are farther away
from the positive metal ions
• down group 7 (halogens)
Element
Melting
Point (K)
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
– increases as the van der Waals’ forces increase
» larger molecules have more
electrons which increases
the chance that one side of
the molecule could be negative
increases
increases
• across the table (period 3)
– from left to right
• bonding goes from strong metallic to very strong
macromolecules (network covalent) to weak van
der Waals’ attraction