Download Periodic Table Properties

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History
Periodic Table Properties
• Russian scientist Dmitri Mendeleev taught
chemistry in terms of properties.
• Mid 1800 - molar masses of elements were
known.
• Wrote down the elements in order of
increasing mass.
• Found a pattern of repeating properties.
Dmitri Mendeleev is
considered the father
of the modern
periodic table
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Mendeleev’s Table
The Modern Table
• Grouped elements in columns by similar
properties in order of increasing atomic mass.
• Found some inconsistencies - felt that the
properties were more important than the
mass, so switched order.
• Found some gaps.
• Must be undiscovered elements.
• Predicted their properties before they were
found.
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Elements are still grouped by properties.
Similar properties are in the same column.
Order is in increasing atomic number.
Added a column of elements Mendeleev didn’t
know about.
• The noble gases weren’t found because they
didn’t react with anything.
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• Group 1A are the Alkali Metals
• Group 2A are the Alkaline Earth Metals
Groups (a.k.a. Families)
• Groups are the vertical columns found within
the periodic table
• The elements in the A Groups are called the
representative elements. Roman Numerals
found at the top of each groups correspond to
the number of valence electrons
• The current system uses traditional
numbers…8, 9, 10 etc. (Please note that the
number of valence electrons is the same as
the second digit)
1A
8A
2A
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3A 4A 5A 6A 7A
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• Group 7A is called the Halogens
• Group 8A are the Noble Gases
1A
Group B are called the transition
elements
8A
2A
3A 4A 5A 6A 7A
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Periods
These are called the inner transition
elements and they belong here
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• Periods are the horizontal rows found within
the periodic table
• Each period corresponds to an energy level
(electron orbit)
• There are 7 periods
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Periods
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3
4
5
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What is the significance of a
full outer energy level?
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• Noble gases (Group 18 or VIIIA) are the only
elements that have the maximum capacity in
their valence shell
• These are the only elements that exist as
individual atoms in nature
PERIODIC LAW
• Rule developed from observations
• “When elements are arranged in order of
increasing atomic number, their properties
show a periodic recurrence and gradual
change”
HAPPY ATOM!
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Atomic Size
Atomic Size
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• First problem where do you start measuring.
• The electron cloud doesn’t have a definite
edge.
• They get around this by measuring more than
1 atom at a time.
Radius
•Atomic Radius = half the distance between two
nuclei of a diatomic molecule.
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Going across a period the A.R. decreases
(left  right)
Group trends
• Electrons are added at the same energy level
with some repulsion between electrons
• Electrostatic attraction between positive
nucleus and negative electrons is stronger
• As positive charge of nucleus increases,
attraction increases
• Electrons are pulled closer to the nucleus,
therefore atomic size decreases
H
Li
• As we go down a group
Na
• Each atom has another
energy level,
K
• So the atoms get bigger.
Rb
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Going down a family the A.R. increases
(top bottom)
Periodic Trends
• As you go across a period the radius gets
smaller.
• Same energy level.
• More nuclear charge.
• Outermost electrons are closer.
• There are more energy levels so distance between
positive nucleus and negative electrons increases
decreasing forces of attraction
• Inner electrons shield the outer electrons from the
full positive charge of the nucleus
• Large increase from Group 18 to 1 (e.g. Ne to Na)
because the single electron added to a new energy
level is shielded from the positive charge by the
electrons in the lower level
• Overall, atomic size increases
Na
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Mg
Al
Si
P
S Cl Ar
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Rb
K
Atomic Radius (nm)
Overall
Figure 1: Atomic radius (pm)
Na
Li
Kr
Ar
Ne
H
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Atomic Number
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What about Cations (positive ions) and
Anions (negative ions)?
Cation Size as Compared to Neutral (Parent)
Atom
• The size decreases when cations form
• The effect is particularly pronounced when all
the valence electrons are lost and only the
noble gas core of electrons remain
• For example, the Mg2+ ion (65 pm radius) is
considerable smaller than the Mg atom (160
pm radius)
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Anion Size as Compared to Neutral
(Parent) Atom
• Size increases when anions form
• The added electrons are going into the same
shell
• They repel each other and so the size
increases
• For example, the Cl- ion (181 pm radius) is
considerably larger that the Cl atom (99 pm
radius)
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Figure 2: Ions and atomic radii
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First Ionization Energy (I.E.)
The energy required to remove the outermost
electron of an atom of an element in its
gaseous state
X (g) + energy

X+
+ 1 e-
(How tightly the electrons are bound)
Figure 3: Atomic Radii
Figure 4: Ionic radii
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Going across the period 1st I.E. increases
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Going down a family 1st I.E. decreases
• Energy level stays the same but number of
protons and electrons increases
• Electrostatic force of attraction increases
• Harder to remove an electron and so I.E.
increases
• More energy levels going down family
• Electrons are further from the nucleus
• By inner electrons shield the positive charge
of the nucleus
• So lower attractive forces
• Less energy to remove an electron
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Second Ionization Energy (2nd I.E.)
The energy required to remove the second
outermost electron from an atom of an
element in its gaseous state
X+ (g) + energy
Figure 5: Ionization energies vary periodically,
which is explained by changes in nuclear
attractive forces.

X2+ + 1 e-
Always greater than first ionization energy!
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Electron Affinity
Going across a period (L to R) (exclude Noble
Gases)
The amount of energy released when an atom
of an element in the gaseous state gains an
electron
X (g) + 1
e-

X-
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• The E.A. increases
• Greater attraction for the electron by the
higher positive charge in nucleus
+ energy
(How high a price the atom will pay for an electron)
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Going down a family
• E.A. decreases
• The electron to be gained would be placed
further away from the nucleus
• Less attractive forces
Figure 6: Electron affinities (kJ/mol)
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• Fluorine is considered to have the greatest
ability to attract electrons, therefore was
assigned the highest value (4.0)
• All other element electronegativities assigned
relative to fluorine
• In general, electronegativities increase
diagonally from the lower left (Cs) to the
upper right (F) of the periodic table
Electronegativity
• When two atoms form a bond, each atom
attracts the other atoms electrons in addition
to its own
• Electronegativity is the measure of an atom’s
ability to attract electrons in a chemical bond
• EN is the symbol for electronegativity
WHY?
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• Smaller atomic size, increase nuclear
attraction, electrons are pulled more tightly to
nucleus  therefore, the atoms attracts a
bonding pair of electrons more strongly and
causes bonding pair to move closer to nucleus
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Electronegativity
• 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.
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Group Trend
Periodic Trend
• The further down a group the farther the
electron is away and the more electrons an
atom has.
• More willing to share.
• Low electronegativity.
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Metals are at the left end.
They let their electrons go easily
Low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away.
High electronegativity.
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Atomic size increases,
shielding constant
Ionization energy, electronegativity
Electron affinity INCREASE
Ionic size increases
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