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Compounds
 A combination of two or more
elements
 Elements form compounds because
they want to be _____________.
 Three types of Compounds:
1.
2.
3.
 Ionic compounds are made of ______
bonds, molecular compounds are
made up of _____________ bonds,
metallic compounds are made of
____________ bonds.
 Ionic –
 Molecular/Covalent –
 Metallic -
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Try It Out
 Which type of compound is:
o NaCl
o H2SO4
o PCl5
o K2S
Valence Electrons
 An atom’s…
 They are the ones that are available to
be …
 The number of valence electrons
corresponds to an atom’s __________
____________.
 The Octet Rule –
 Atom’s gain, lose, or share electrons
(BOND) so that they have an octet.
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Electron Dot Structures
 An element’s symbol with its valence
electrons drawn around it
 To Draw:
1.
2.
3.
 After you write the symbol, the areas
around the symbol represent the
orbitals and axis:
S
Example
 Carbon
 Aluminum
 Bromine
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Try it Out!
 Nitrogen
 Krypton
Ionic Bonds
 Bond in which one atom ________ an
electron and one ________ an
electron
 Between a ___________ (positive ion)
and an ___________ (negative ion).
 Bonds are made when opposite
charges are attracted to one another
by __________________________.
 Though ionic compounds are made
from __________ (charged particles)
they are _______________________.
 The total charge of the compound is
___.
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 Chemical Formula – The ___________
in the formula tell you how many
atoms of each element are present in
the compound.
o H2 0 =
o C6H12O6 =
o NaCl =
o
 Example 1 – Lithium + Bromine
 Example 2 – Calcium + 2 Fluorine
Review: Draw the bonding that
takes place between Barium and 2
Iodines. Determine the charge for
each ion and write the formula:
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Covalent Bonds
 ______________ of electrons
o
o
 Molecular Formula –
 Diatomic Molecule –
 The formula does not tell you about a
molecule’s structure
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 Covalent bonds can be …
 Bonds between atoms are represented
by a _________. A double bond is
two lines and a triple bond is three.
o Ex: H-H
 Example 1 – Iodine + Iodine
 Example 2 – Nitrogen + Nitrogen
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Energy in Bonds
 When bonds are made, energy (heat)
is _________________.
 When bonds are broken, energy is
__________________.
 It takes more energy to break
_____________ bonds. Triple bonds
are the strongest. Single bonds are
the ________________.
 Bond dissociation energy –
 The larger the bond dissociation
energy, the ____________ the bond
Lewis Structures
 Using electron dot structures, we can
draw Lewis structures – formulas that
show how elements bond and where
their electrons go.
 Only work for COVALENT compounds
because they show how atoms
____________ electrons
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 For ionic compounds, we draw dot
models demonstrating how electrons
move
o Realize electrons are ________
being __________
o The more electronegative atom
__________ electrons
 Lone Pair –
 Bonds –
Preview of Lewis Structures
 Counting Valence Electrons – NH3
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 Organize the atoms so that there is a
____________ atom (usually the least
________________) surrounded by
the outer atoms.
 __________________ is never the
central atom.
Lewis Structure Steps
1. Count the total valence electrons for the molecule:
To do this, find the number of valence electrons for
each atom in the molecule, and add them up.
2. Write a plausible skeletal structure and connect
the atoms by single dashes (covalent bonds):
Remember, a bond is made of 2 electrons.
3. Place pairs of electrons as lone pairs around the
terminal atoms to give each terminal atom (except
H) an octet: The octet rule tells us that all atoms want
eight valence electrons (except for H which only wants
two) so they can be like the nearest noble gas.
4. Assign any remaining electrons as lone pairs
around the central atom:
There are a few weird elements:
 Boron wants six
 Phosphorus, sulfur, chlorine, and a few others
can have 10 or 12 valence electrons
5. If necessary (if there are not enough electrons),
move one or more lone pairs of electrons from a
terminal atom to form a multiple bond to the
central atom.
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Some handy rules to remember are these:
 Hydrogen and the halogens bond once.
 The family oxygen is in, can bond twice.
 The family nitrogen is in, can bond three times.
So can Boron.
 The family carbon is in, can bond four times.
Example
 CH4
 CO2
 PCl5
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Try it Out!
 SiS2
 NH3
For Charged Compounds
 Negative Charge =
 Positive Charge =
 Write charge outside of the structure.
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Example
 OH-1
Resonance Structures
 Definition –
 Example – NO2-1
Wrap Up: Ionic or Covalent?
1.
2.
3.
4.
5.
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VSEPR
 Stands for..
 The repulsion between electron pairs
causes molecular shapes to adjust so
that the electrons are as far away from
each other as possible.
 Shows _____________ of molecules.
VSEPR Rules
1.
2.
 Lone pairs still occupy space, but do
not determine the shape of the
molecule.
3.
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ss-chemistry-textbook/advanced-concepts-of-chemicalbonding-10/molecular-geometry-79/table-of-geometries354-10522/
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Linear
o
Bent
o
Triangular Planar
o
Triangular Pyramidal
o
Tetrahedral
o
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Other Shapes:
https://drushapchem.wikispaces.com/%E2%80%A2+VSEPR
Polarity
 The different electronegativities of the
atoms in a compound determine
whether the molecule is polar or
nonpolar
 Polarity only exists in
______________________
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Nonpolar Bonds
 Exists when 2 electrons are shared
______________ by 2 nonmetallic
atoms.
 ________________ distribution of
charge.
 Nonpolar Example – Diatomic
Molecules
Polar Bonds
 Bonded atoms have an
________________ attraction for the
shared electrons.
 Charge _________ __________
distributed equally.
Example –
 Whether a molecule is polar or not can
be determined by its _____________.
 If it is _______________ (uniform all
around) it is ________________.
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 If it is NOT _______________(ununiform around) it is
______________.
 If a compound has charge, no matter
it’s shape, it is __________.
Bond Dipoles
 A polar covalent bond has a ______
____________; a separation of
positive and negative charge in an
individual bond.
 Bond dipoles have both a magnitude
and a direction (they are
____________ quantities).
 Ordinarily, a polar molecule must have
polar bonds, BUT … polar bonds are
not sufficient.
 A molecule may have polar bonds and
be a nonpolar molecule – IF the bond
dipoles cancel.
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 To predict polarity:
1. Use electronegativity values to predict
bond dipoles.
2. Use the VSEPR method to predict the
molecular shape.
3. From the molecular shape, determine
whether bond dipoles cancel to give a
nonpolar molecule, or combine to
produce a resultant dipole moment for
the molecule.
PRACTICE- Use VSEPR to predict the
shape & polarity of these molecules:
 BCl3
 SiCl3Br
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 PI3
 CCl4
Wrap Up: Draw as many resonance
structures as you can for SiS3-2.
Indicate the shape of the molecule.
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Metallic Bonds
 Best characterized by the phrase,
“_________________________”

The valence electrons are
____________ and can drift freely
from one part of the metal to another
 The fact that electrons ________
freely helps to explain some of the
_________________ of metals:
o
o
o
 Alloys –
 Their properties are often superior to
those of their component elements
 Examples:
o
o
o
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Types of Forces
 Thus far, we’ve talked about
___________________ forces (the
forces ________________ molecules.
o Ionic bond
o Covalent bond
o Metallic bond
 There is another, equally important
force that takes place ____________
molecules.
 These forces are known as
_______________________ forces.
Intermolecular Attractions
 Intermolecular forces occur in
covalent/molecular compounds.
 The attractions between molecules are
much ______________ than the
covalent bonds within molecules.
 These attractions vary depending on
the nature of the molecules.
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Dipole-Dipole Forces
 Many molecules have permanent
_______________.
 These ____________________
molecules form attractions between
the positive end of one molecule and
the negative end of the adjacent
molecule.
 Dipole-dipole attractions are effective
over short distances between
molecules and decrease as the
distance between molecules
__________________.
Hydrogen Bonding
 Strong dipole-dipole attractions occur
in molecules containing…
 The hydrogen in these bonds has a
partial positive charge, since there are
no inner core electrons and the shared
electrons are strongly attracted to the
small, very electronegative atoms.
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 The attractions in hydrogen bonds are
_________ to ________ times
stronger than other dipole-dipole
attractions.
 Hydrogen bonding is responsible for
the unusual properties of water and is
very important in biological systems
such as _______________ and
_______________.
London Dispersion Forces
 ___________________ do not have
permanent dipoles.
 At any given instant, the electrons
may be unevenly distributed within an
atom or molecule giving it a partial
negative charge.
 This results in atoms being attracted
to one another for an instant (can
even happen in _________________)
 The greater the molar mass (more
_____________) the stronger the
dispersion forces.
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Let’s look at some molecules:
Properties of Compounds
 Compounds have different properties
because of the types of bonds and
intermolecular forces that they have.
 The bond types help to explain the
behavior of many household items.
Ionic
Covalent
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Melting Point
 Ionic compounds have a higher
melting point because there is such a
___________________ in the bonds.
 It takes a higher temperature
(_________________) to break them.
Conducting Electricity
 Ionic compounds are able to conduct
electricity because they contain _____
(charged particles).
 Electricity is the movement of
electrons and the ions spread out in
solution and allow electricity to be
transferred.
Wrap Up: Draw the Lewis structure
for CCl3H and indicate which
inermolecular force(s) are present.
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