Transcript ATOMIC STABILITY - Barnegat Township School District

Chapter 8 Honors Chemistry
(partial)
Covalent Bonding
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Electronegativity
• In a covalent bond, we have seen that electron
pairs are shared between two nonmetals
• Rarely are these electrons shared equally as
one of the atoms has a stronger “desire” to have
those electrons
• How can we measure which atoms wants the
electrons more?
• Electronegativity (EN) !!!!!
• It is a measure of an atoms ability to attract a
pair of electrons in a molecule
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Electronegativity
F is the most electronegative element
and is given a value of 4.0 and all
elements E.N. values are in
comparison to this
Left to right across a period =  in EN
Down a group  in EN or stays about
the same
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 The higher the EN value, the more the atom will
attract shared electrons to it
 Depending upon how great the difference in
electronegativity is between the atoms the bond
can have highly positive and negative regions
 This is called a polar bond
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Electronegativity
• The only bond that is purely 100% covalent
where the electrons are equally shared is one in
which the EN = 0
• This only occurs when the electrons are shared
by identical atoms, like H2, or any of the
diatomic molecules
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Polarity
• A bond is considered to be non-polar covalent
if the EN is 0 – 0.4
• A bond is considered to be moderately polarcovalent if the EN is 0.5 – 1.0
• A bond is considered to be very polarcovalent if the EN is 1.0 – 1.7
• Any bond with EN that is ≥ 1.7 is considered
to be ionic
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Electronegativities of the Elements
Electronegativity
 Using the chart of Electronegativities (Pg. 177), determine the type
of bond formed between the following pairs of atoms:
 C and O
 Fe and O
 N and Br
 C and H
 Na and F
 Cl and Cl
 C and O
EN = 1.0 polar
 Fe and O
EN = 1.7 ionic
 N and Br
EN = 0.2 nonpolar
 C and H
EN = 0.4
nonpolar
 Na and F
EN = 3.1
ionic
 Cl and Cl
ΔEN = 0
nonpolar
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Electronegativity
• In a molecule of H2O, a pair of electrons are
shared between each O and H
• The EN of O = 3.5 and H = 2.1
• EN = 1.4 – therefore is a polar covalent bond
• This means that O attracts the electrons
towards it and so will become slightly negative
while the electrons move away from each H
atom and they become a bit positive
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• This means the electrons are not shared evenly
and that one area is slightly positive, the other
negative.
• This is called a polar molecule
• Indicated using small delta (δ).
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Dipole Moments
• A molecule with a center of negative charge and a
center of positive charge is a DIPOLE
• (two poles),
• or has a dipole moment.
• Center of charge doesn’t have to be on an atom.
• Will line up in the presence of an electric field.
How It is drawn
d+ d-
H-F
d+ d-
H-F
d+ d-
H-F
-
+
d+ d-
H-F
-
d+ d-
d+ d-
H - F d+ d- H - F
H-F
d+ d-
d+ d-
H-F
H-F
d+ d-
H-F
+
d+ d-
H-F
d+ d-
H-F
Which Molecules Have Dipole Moments?
•
Any two atom molecule with a polar bond.
H2O or FBr
•
With three or more atoms there are two considerations.
1.
There must be a polar bond.
2.
Geometry can’t cancel it out (more about geometry later)
CH4
CO2
SO2
Ionic vs. Molecular Compounds
 There are two types of forces involved in chemistry
 Intermolecular forces are those between molecules and are
responsible for holding these molecules together (inter =
between)
 Intramolecular forces are those between atoms inside the
actual molecule and are responsible for holding the molecule
together (intra = within)
 These two forces explain many of the properties of ionic and
covalent compounds
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Ionic vs. Molecular Compounds
• Ionic compounds are formed of positive and
negative ions and these forces are very
strong
• Each ion is held in place by at least 6 other
ions and so both the inter and the intra
molecular forces are strong
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• Covalent compounds have strong
intramolecular forces holding the atoms
together to form a molecule, but rather weak
intermolecular forces holding the adjacent
molecules together
• Because the intermolecular forces are weak,
covalent compounds have low boiling and
melting points (little energy is needed to move
molecules apart from a solid to liquid to gas)
• Many are gases at room temp
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Ionic vs. Molecular Compounds
• Solid ionic compounds do not conduct electricity
as the ions are held tightly, but when in the liquid
state (called molten) the ions are free to move and
so can conduct electricity
• Ionic compounds dissolve easily in water as water
is a polar molecule and water molecules surround
the ions and pull them apart into the solution
(process called solvation)
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Metallic Bonding
• How are metal atoms held in place?
• Most metals have 1, 2 or 3 valence electrons
• The metal atoms are relatively close to each other and their valence
energy levels overlap
• This allows the valence electrons to move freely from one metal
atom to those it overlaps with
• These electrons are not bonded to one particular metal atom and are
called delocalized electrons
• This is often referred to as the “Electron Sea Model” of metallic
bonding
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Metallic Bonding
Ionic Bond, A Sea of Electrons
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Sea of Electrons
• Metals conduct electricity.
• Electrons are free to move through the
solid.
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Hydrogen Bonding
• This is a type of bonding involving
hydrogen and either F, O or N
• When hydrogen bonds with either of these
elements there is a large ΔEN
• This results in a very polar molecule with
large dipoles
• This produces relatively high inter
molecular forces to adjacent molecules they
are held together “tightly”
• This accounts for the relatively high boiling
and melting point of H2O compared to other
covalent compounds
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Hydrogen Bond
The hydrogen bond is a special dipole-dipole
interaction between the hydrogen atom in a polar
N-H, O-H, or F-H bond and an electronegative O,
N, or F atom.
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