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  Atomic Radius:
The distance from the centre of the nucleus to the outermost
electron.
  Two factors must be taken into consideration in
explaining this periodic trend:
1.  Increasing nuclear charge.
2.  Increasing number of shells.
  Along a Period (left to right):
  The atomic number increases (more protons) while the
valence electrons remain in the same shell.
  Due to the increasing nuclear charge (pulling electrons
closer to the nucleus):
The atomic radius increases
from right to left.
  Along a Group (top to bottom):
  The atomic number continues to increase.
  However, the shell increases from shell 1 to shell 2 etc.
  The atomic orbitals for each successive shell get larger
and larger.
  The result is:
The atomic radius increases
from top to bottom.
Atomic Radius vs. Atomic Number
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  Cations:
  Get smaller than the atomic species because you lose outer
electrons and the net positive charge draws in remaining
electrons.
  Anions:
  You gain electrons and they repel each other.
  The ionic radius expands out to accommodate the repulsive
forces.
  For Example:
  For the Sodium atom:
  The energy needed to remove an electron from the
atom.
Na(s) → Na+(s) + eEI = Ionization energy = 496 kJ/mol
  It is a minimum for the alkali metals which have a
single electron outside a closed shell.
  It generally increases across a row on the periodic table
with a maximum for the noble gases which have
closed shells.
Ionization Energy vs. Atomic Number
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  The representative elements tend to gain or lose
electrons until they become isoelectronic with the
nearest noble gas.
  Imagine a campfire…
  The same thing happens in the atom:
When a new orbital is started, every
orbital of lower energy shields these
electrons from feeling the full nuclear
charge.
  Fluorine (the most electronegative
element) is assigned a value of 4.0.
  A chemical property that describes the ability of an
atom to attract electrons towards itself in a covalent
bond.
  Values range down to Cesium and Francium
which are the least electronegative at 0.7.
  First proposed by Linus Pauling in
1932.
  He developed the Pauling Scale. It
uses a dimensionless number from
0.7 to 4.0.
Electronegativity of Atoms in the Periodic Table
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  The energy released when 1 mole of gaseous atoms
each acquire an electron to form 1 mole of gaseous 1ions.
X(g) + e- → X-(g) + energy
  The above represents the first electron affinity (Eea).
  Do:
  W.S. 5-3 (use Text Pages 174-180)
Now, where
did I leave my
keys…
  From 1916-1919, Gilbert N.
Lewis made several
important proposals on
bonding which lead to the
development of Lewis
Bonding Theory.
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  LEWIS SYMBOLS:
  A common chemical symbol surrounded by up to 8
dots.
  The symbol represents the nucleus and the electrons of
the filled inner shell orbitals.
  The dots represent the valence electrons.
This only works well for the representative
elements.
  For Example:
Transition metals, actinides and lanthanides
have incompletely filled inner shells - we can't
write simple Lewis structures for them.
  Ionic (metal/non-metal)
  Covalent (non-metal/non-metal)
  Intermolecular Bonding (between molecules):
  Hydrogen Bonding
  London Dispersion Forces
  Ionic bonds are forces that hold ionic compounds
together.
  Forming the ionic bond:
  Step 1:
  A cation forms by the LOSS of 1 or more e-.
  Representative elements become cations, isoelectronic
with the nearest Noble gas.
  For others (transition metals), not necessarily.
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  ATOMS → COMPOUNDS, properties change…
  Their reactivity decreases considerably.
  They become neutral overall.
  There are no unique molecules in many ionic
solids.
  Ionic compounds become electrically
conductive when melted or dissolved.
IONIC BONDS are STRONG so that
compounds held together by ionic bonds
have HIGH MELTING
TEMPERATURES.
  The valence electrons involved in the bond are called
the BONDING ELECTRONS or the BOND PAIR.
  Those not involved in the bond are called the
NONBONDING ELECTRONS or the LONE
PAIRS.
Lone
Pair
Bond
Pair
  Covalent bonds arise from the sharing of electrons
between atoms (generally of groups IVA, VA, VIA, and
VIIA).
  Each electron in a shared pair is attracted to both
nuclei involved in the bond.
COVALENT BONDS are VERY STRONG.
  OCTET RULE :
  An atom other than hydrogen tends to form bonds until
it is surrounded by eight valence e-.
  The pairs repel each other and thus tend to stay as far
away as possible.
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  Imagine we have one atom with a somewhat higher
electronegativity than the other in a covalent bond:
  This will cause the electrons to be shared unevenly, such
that the shared electrons will spend more time (on
average) closer to the atom that has the higher
ELECTRONEGATIVITY.
  Do:
  W.S. 7-1
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  Polar molecules, such as water molecules, have a weak,
partial negative charge at one region of the molecule
(the oxygen atom in water) and a partial positive
charge elsewhere (the hydrogen atoms in water).
  When water molecules are close together, their positive
and negative regions are attracted to the oppositelycharged regions of nearby molecules.
  Each water molecule is hydrogen bonded to four
others:
  The London dispersion force is the weakest
intermolecular force.
  It is a temporary attractive force that results when the
electrons in two adjacent atoms occupy positions that
make the atoms form temporary dipoles.
HYDROGEN BONDS are WEAK.
  Because of the constant motion of the electrons, an
atom or molecule can develop a temporary
(instantaneous) dipole when its electrons are
distributed unsymmetrically about the nucleus.
  Dispersion forces are present between any two
molecules (even polar molecules) when they are almost
touching.
LONDON DISPERSION FORCES are the
WEAKEST.
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Simple Ionic Compounds
  Easy to do…
  For Example:
  KBr
  Li2S
  The overall charge on the compound must equal zero,
that is, the number of electrons lost by one atom MUST
EQUAL the number of electrons gained by the other
atom.
  K3P
  The Lewis Structure of each ion is used to construct
the Lewis Structure for the ionic compound.
1.  Find the total number of valence e-.
  Lewis structures show how the VALENCE electrons are
distributed in a molecule.
  Covalent compounds share electrons to fill their
valence shells.
  Therefore, the Lewis structures for these compounds
are drawn a little differently.
  Go by the column or group that it’s in.
2.  After connecting your central atom to the terminal atoms with single
bonds, begin adding the remaining valence e- as lone pairs.
  First around the terminal atoms.
  Then around the central atom (if you have any left over).
3.  Satisfy the octet rule for your central atom by either:
  Replacing a lone pair on your terminal atoms with a bond (to
make a double bond).
OR:
  By replacing two lone pairs with two bonds (to make a triple
bond).
  You can’t just add double bonds without
first removing a lone pair.
  Not only are you adding more electrons
than you started with, but you’re
probably breaking the octet rule for the
terminal atoms.
  NH4+
  O2
  NO2  C2H4
  CO
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Incomplete Octets:
  In addition to H, Be, B, and Al are exceptions to the
octet rule.
  Since they have very low electonegativities, they can
only accept one electron for every one they donate.
  For Example:
BF3
  For Example:
  PCl5
  Do:
  SF4
  W.S. 7-2
  Drawing Lewis Structures W.S.
  XeF4
  Periodic Trends
  Atomic and Ionic Radii
  Ionization Energy
  Electron Affinity and
Electronegativity
  Lewis Theory
  Elements of the Theory
  Ionic Bonds
  Covalent Bonds
  Multiple Bonds
  Polarity
  Intermolecular Bonding
  Hydrogen Bonding
  London Dispersion
Forces
  Drawing Lewis structures
  Simple Ionic
Compounds
  Structures that Obey the
Octet Rule
  Structures that Violate
the Octet Rule
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