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Topic 3: Periodicity
3.1 The periodic table
3.1.1
Describe the arrangement of elements in the periodic table in order
of increasing atomic number
3.1.2 Distinguish between the terms group and period
3.1.3 Apply the relationship between the electron arrangement of
elements and their position in the periodic table up to z=20.
3.1.4 Apply the relationship between the highest occupied energy level
for an element and its position in the periodic table.
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Groups: vertical columns (18)
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Have similar properties because have same
number of electrons in outer shell
Periods: horizontal row (7)
Family Names:
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Group 1: alkali metals
Group 2: alkaline earth metals
Group 17: halogens
Group 18: noble gases
Group 3-12: Transition metals
Groups 1,2, 13-18: representative elements
3.2 Physical properties
3.2.1
Define the terms first ionization energy and
electronegativity
3.2.2
Describe and explain the trends in atomic radii,
ionic radii, first ionization energy, electronegativities and
melting points for alkali metals (Li  Cs) and the halogens
(F  I).
3.2.3
Describe and explain the trends in atomic radii,
ionic radii, first ionization energy, and electronegativities for
elements across period
3.2.4
Compare the relative electronegative values of two
or more elements based on their position on the periodic
table.
Atomic Size
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The electron cloud doesn’t have a definite
edge.
They get around this by measuring more than 1
atom at a time.
Summary: it is the volume that an atom takes
up
http://www.mhhe.com/physsci/chemistry/essen
tialchemistry/flash/atomic4.swf
Group trends
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As we go down a
group (each atom has
another energy level)
the atoms get bigger,
because more protons
and neutrons in the
nucleus
H
Li
Na
K
Rb
Periodic Trends
atomic radius decreases as you go from left to right across a
period.
 Why? Stronger attractive forces in atoms (as you go from
left to right) between the opposite charges in the nucleus
and electron cloud cause the atom to be 'sucked' together a
little tighter. Remember filling up same energy level, little
shielding occurring.
Na
Mg
Al
Si
P
S Cl Ar
Ionic Size
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Cations form by losing electrons.
Cations are smaller than the atom they come
from.
Metals form cations.
Cations of representative elements have noble
gas configuration.
Ionic size
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Anions form by gaining electrons.
Anions are bigger than the atom they come
from.
Nonmetals form anions.
Anions of representative elements have noble
gas configuration.
Periodic Trends
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Metals losing from outer energy level, more protons
than electrons so more pull, causing it to be a smaller
species.
Non metals gaining electrons in its outer energy level,
but there are less protons than electrons in the
nucleus, so there is less pull on the protons, so found
further out making it larger.
Li+1
N-3
B+3
Be+2
C+4
O-2
F-1
Size of Isoelectronic ions
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Positive ions have more protons so they are
smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
Electronegativity
Electronegativity
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The tendency for an atom to attract electrons to
itself when it is chemically combined with
another element.
How fair it shares.
Big electronegativity means it pulls the
electron toward it.
Atoms with large negative electron affinity
have larger electronegativity.
Group Trend
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The further down a group the farther the electron is
away and the more electrons an atom has.
So as you go from fluorine to chlorine to bromine and
so on down the periodic table, the electrons are
further away from the nucleus and better shielded
from the nuclear charge and thus not as attracted to
the nucleus. For that reason the electronegativity
decreases as you go down the periodic table.
Period Trend
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Electronegativity increases from left to right
across a period
When the nuclear charge increases, so will
the attraction that the atom has for electrons in
its outermost energy level and that means the
electronegativity will increase
Period trend
Electronegativity increases as you go from left to right
across a period.
 Why? Elements on the left of the period table have 1
-2 valence electrons and would rather give those few
valence electrons away (to achieve the octet in a
lower energy level) than grab another atom's
electrons. As a result, they have low electronegativity.
Elements on the right side of the period table only
need a few electrons to complete the octet, so they
have strong desire to grab another atom's electrons.
Group Trend
electronegativity decreases as you go down a group.
 Why? Elements near the top of the period table have few
electrons to begin with; every electron is a big deal. They have
a stronger desire to acquire more electrons. Elements near the
bottom of the chart have so many electrons that loosing or
acquiring an electron is not as big a deal.
 This is due to the shielding affect where electrons in lower
energy levels shield the positive charge of the nucleus from
outer electrons resulting in those outer electrons not being as
tightly bound to the atom.
Shielding
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Shielded slightly from the
pull of the nucleus by the
electrons that are in the
closer orbitals.
Look at this analogy to help
understand
Melting Points of Group 1
Element
Melting Point (K)
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
Metallic bonding
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Collective bond, not a single bond
Strong force of electromagnetic attraction between
delocalized electrons (move freely).
This is sometimes described as "an array of positive
ions in a sea of electrons
Why does the melting point decrease going
down the alkali metals family?
 Atoms are larger and their outer electrons are
held farther away from the positive nucleus.
 The force of attraction between the metal ions
and the sea of electrons thus gets weaker down
the group.
 Melting points decrease as less heat energy is
needed to overcome this weakening force of
attraction.
Melting Points for halogens
Element
Melting Point (K)
Fluorine
85
Chlorine
238
Bromine
332
Iodine
457
Astatine
610
Why does melting point increase
going down the halogens?
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The halogens are diatomic molecules, so F2,
Cl2, Br2, I2
As the molecules get bigger there are more
electrons that can cause more influential
intermolecular attractions between molecules.
The stronger the I.A, the more difficult it will
be to melt. (more energy needed to break the
I.A)
What are these I.A?
van der Waals forces:
 Electrons are mobile, and although in a
diatomic molecule they should be shared
equally, it is found that they temporarily
move and form slightly positive end and
negative end.
 Now that one end is + and the other -, there
can be intermolecular attractions between the
opposite charges of the molecules
van der Waals forces
IB requires knowledge specifically for
halogens. Check out this site for more detail.
http://www.chemguide.co.uk/inorganic/group7/p
roperties.html
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