Review Book Topic 6
Introduction:
 BONDING is the glue that hold all compounds
together.
 There are many different bond types and these
differences in bond type give compounds their specific
properties and determine how they react with other
compounds.
 Lewis Dot Structures: picture representation to help
you see how bonds occur in specific compounds.
Energy and Chemical Bonds
 Review:
 Endothermic: energy is added and absorbed due to actions
like breaking bonds. (a +ΔH)
 Exothermic: energy is released in reactions like the making of
compounds. (a –ΔH)
*When chemical bonds are formed, the resulting compound
has less potential energy than the reactants because energy is
always released in forming compounds.
**The more energy released, the more stable that resulting
compound is.
Ex: B+C-> BC + 100 joules is less stable than I+S-> IS+ 400
joules
LOOK ON TABLE I to find out the heat released or absorbed
by a reaction.
Lewis Dot Diagrams
 Also called an electron dot diagram.
 This diagram consists of a chemical symbol surrounded by
the number of electrons in that elements valence orbital
represented by dots . REMEMBER: Valence electrons are
the only ones that participate in bonding.
 The most surrounding “dots” any element could have is 8
(octet)….your noble gases.
 Be sure you refer to the LAST number on the electron
configuration of your element.
 KERNEL: everything in the atom that is positively charged
Write the element and start to encircle it with its
valence electrons on four sides- placing one at a time in
a clockwise fashion starting at the top.
The Octet Rule
 Atoms: (other than Hydrogen and share electrons
until they are surrounded by EIGHT valence electrons.
This is called the “OCTET RULE”. The magic number
is eight, due to all elements wanting to have a full
valence shell like the noble gases….the valence shell
can hold a total of eight electrons. Be sure that all of
the elements in your Lewis Dot Diagrams have the
complete octet!
Lewis Dot Diagram: Compounds
 For covalent compounds (two non-metals that SHARE
electrons), you must first draw out the two elements
separate, take the pairs that overlap and make them
into lines.
 Every PAIR of “dots” makes ONE line and unpaired
electrons remain dots!
 EX: Two hydrogen making H2
Rules for Covalent Compounds
Determine the total number of valence electrons by
adding the valence number for each element in the
compound.
2. Arrange the atoms so that the central atoms has the
smallest electronegative value (usually the odd on in the
compound, only has one of them)
3. Draw single bonds (lines) between center element and
surrounding ones (remember: each line = 2 electrons)
4. Place the remaining electrons in pairs around elements
that need the octet
5. If there are still unpaired central electrons..you must have
it share if there outer elements have any free
6. OR you can do it Ms. Shield’s way….
(remember Hydrogen only gets 2 and B only gets 6!)
1.
Example:
 Let’s try CH3Cl:
Counting valence electrons:
1 carbon with 4 valence= 4
3 hydrogen with 1 each= 3
1 chlorine with 7= 7
Total= 14 electrons
2. If Carbon is in your compound…it is usually the
central element.
1.
H
H
C Cl
H
Example Cont…
3. Place lines in between the central and outer elements:
Cl
4. Last we must place the remaining electrons. Above we
have a total of 8 already there…so we have 6 electrons
remaining. Seeing as Hydrogen already holds two
electrons (a line), then Cl must get the rest since C has an
octet already. So,
DONE!!!!!
Multiple bond between elements
 Sometimes there are more free electrons “left over” than
can be placed around the elements…this means there are
multiple electrons being shared.
 Follow the same rules as before (#1-3), but at the last step
you may have to add another “line”.
 Lets try C2H2
 So we know our Cs must be our middles and the Hs are on the
outside and that we have a total of 10 electrons
 We get: H-C-C-H……but we have four electrons left and H is
already full!!! If you look at the C’s, they don’t have their octet
so we must make more bonds between the C’s
Example Cont.
 In the end we have to add two more “line” bonds
because each line is two electrons.
 The answer would be :
H-C
C-H
Covalent Bonding
http://www.youtube.com/watch?v=ulyopnxj
AZ8&feature=related
Your Turn!
Try these:
N2
HCl
Cl2
Answers:
 N2
N
 HCl
H-Cl
 Cl2
Cl-Cl
N
Covalent Bonds:
 We know covalent bonds are between two non-metals
and there is a sharing of electrons. There are three
types of bonds that can result from covalent bonding:
1. Non-polar covalent bonding: when the pulling
attraction is equal on all sides of the compound
(found when it’s a diatomic element or when one
element is surrounded by the same element).
2. Multiple covalent bonding: when there are multiple
bonds (lines) in between two elements (like in I2 or
O 2)
3. Polar covalent bonds: when atoms in a bond are
unequal electronegativity, the pull each exerts will
not be even.
Polar Molecules
•
Polar: when atoms in a bond are unequal, the pull
each exerts will not be even. (Think two poles…as
in the North and South pole.)
•
How can you tell which bonds are MOST or
LEAST Polar?
• Polarity is determined by the difference in
Electronegativity values. You must go to Table
S and subtract the electronegativity of the
elements in the bond. The higher the
difference, the more polar the bond is!
All bonds will form a “Geometry”,
or a shape. This will affect how
they interact and the properties it
has. The bonds we will look at new
you MUST know for your regents!
Bond Shapes
http://www.youtube.com/watch?v=keHSCASZfc&feature=related
Ionic bonding
 Lewis dot diagrams for ionic bonds are just like
covalent bonds, but they will form ions (charges) due
to which elements will lose or gain electrons.
 REMEMBER: Non-Metals gain electrons, while Metals
lose electrons!
Ionic Bonding
 Since ionic bonds have metals and non-metals
transferring electrons, ions will form. When these
ions form, they will carry a charge with them.
http://www.youtube.c
om/watch?v=k-nQD_mu7o
Metallic Bonding
 Metallic Bonds occur between two metals …it will be
between the same element though.
 So many Cr together would form metallic bonds.
Lewis Dot Diagram: Ions
 Be sure to notice the different way IONS
are expressed
 REMEMBER: Ions are elements that lose
or gain electrons, making them charged!
Lewis Dots of Polyatomic Ions:
Same rules apply, at the end they get brackets and a
charge.
When you draw an ion, don't forget [ ] and a
charge.
Distinguishing Bond Types
 All different bonding types have different properties;
 Metals have a high melting/boiling point, as do ionic compounds
(they have a metal), but are not good conductors like regular metals,
yet are usually hard in quality. The only state of a compound with a
metal present (ionic or metallic) that cannot conduct is ionic in solid
form.
 Molecular (covalent) compounds are soft and poor conductors, they
will have a low melting/boiling point
 MAKE SURE TO MEMORIZE THE CHART IN YOUR REVIEW
BOOK!
 Remember: if there is a metal present- ionic or metallic- they will
have a high boiling/melting point and will have good conductivity as
long as it is not ionic in solid form!
Intermolecular Forces:
 Hydrogen Bonding:
Is very important to living things~Keeps two strands of DNA together in
double helix
~ Causes water molecules to stick together and for
water characteristics.
~Know it is the strongest when bonding with
N, F, and O
 Dipole-Dipole: two equal and opposite electrically
charged poles that are separated by a short distance
You MUST know:
 How to form Lewis Structures from Polyatomic ions, Covalent
compounds, and Ionic compounds.
 How to name the Geometry of a compound
 Tell if the compound is “Polar” or “Non-Polar” and “least” or “most”
polar (table S)
 The more energy released from a reaction= elements stuck together
more= greater stability. To know this, LOOK AT TABLE I!
(Endothermic= +ΔH) and (Exothermic= -ΔH)
 Review Videos:
 http://www.youtube.com/watch?v=ulyopnxjAZ8&feature=related
(covalent)
 http://www.youtube.com/watch?v=k--nQD_mu7o (ionic)
 http://www.youtube.com/watch?v=keHS-CASZfc&feature=related
(vesper theory= geometry)
You’re Finished!