Lets Explore The Formation Of A Covalent Bond

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Lets Explore The Formation Of A Covalent Bond
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How do we know when two elements can combine to
form a compound?

• One way to understand bonding deals with the
  forces of repulsion and the attraction
• Repulsion between the negative electron clouds of
  each atom
• Attraction between the positive nuclei and the
  negative electron clouds

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• When two atoms approach each other closely enough
  for their electron clouds to overlap
• The electrons of one atom repel the electrons of
  the other
• The nuclei of one repels the nuclei of the other
• The electron cloud of one atom attracts the
  nuclei of the other

As the optimum distance is achieved that balances
these forces, there is a release of energy.
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With a small energy difference the bond formed is
very weak, but if the energy difference is large
then we speak of a strong chemical bond.
There are two types of strong chemical bonds
ionic bonds and covalent bonds.
The atoms involved and the energy difference in
their bonding decides which type of bond forms.
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The formation of a chemical bond is a favored
process
By rearranging the electrons so that each atom
resembles a noble gas, these atoms become more
stable paired up.
Sometimes to develop the appearance of a noble
gas there is a transfer of electrons between 2
atoms and sometimes a sharing of electrons
between two atoms
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The type of bond that forms can also be
determined by examining the amount of difference
the atoms have in terms of electronegativity.
Electronegativity is a measure of the
ability an atom has to attract electrons in a
chemical bond.
The greater the difference in electronegativities
the more ionic the bond is.
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lets consider the compound cesium fluoride, CsF,
as an example.
the electronegativity value (EV) for Cs is .70
the EV for F is 4.00.
The difference between these values is 3.30,
which falls on a scale of ionic character.
When the difference in electro-negativities is
greater than 2.1 the bond is classified as mostly
ionic.
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The take home lesson with electronegativity is
this
The closer the atoms are on the periodic table,
the more evenly they share their electrons, and
therefore more likely to form covalent bonds
The farther apart they are on the periodic table,
the less evenly they share their electrons, and
there-fore more likely to form ionic bonds.
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Covalent Bonds
In a covalent bond The electronegativity
differences in the 2 atoms involved is not
extreme, so the electrons that are interacting
are shared It may not be an equal sharing, but
at least the electrons are being shared.
Consider the interaction between H and O. The O
has a much higher electronegativity than the H
atom.
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O has a higher electronegativity, so the pair of
electrons will not be shared equally between H
and O.
When a covalent bond forms with unequally shared
electrons, the bond is said to be a polar
covalent bond.
The uneven sharing causes the more
electronegative atom to have a partial negative
charge. The other atom will have a partial
positive charge.
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If the whole molecule has a tendency to develop a
positive end and a negative end to it, because of
differences in electronegativity it is called a
polar molecule
Sets up opposite charges on opposite sides
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These oppositely charged poles set up by the
electronegativity differences in atoms causes
interaction between molecules.
The partial negative charges of one molecule
attracts the positive end of another molecule and
so on setting up a network of loose connections
These connections lead to different chemical
properties such as higher boiling points because
it takes more energy to pull the molecules apart.
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Special kind of polar interaction called
hydrogen bonding
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Properties of covalent bonding
We looked briefly at the energy difference as two
hydrogen atoms approached each other
The balance set up by repulsion and attraction
between two atoms in a covalent bond, is what
holds the atoms together
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The distance the two atoms are from each other to
achieve this happy medium is called the bond
length, and it is inversely related to its bond
energy.
The shorter the length of the bond, then the
higher the energy necessary to pull them apart
and vice a versa.
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The bond lengths are not really fixed values
because the atoms vibrate much like theyre
attached by a tiny spring
We can measure the average length of the bonds
and place them in charts.
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As you can see from the table the larger the
atoms, then usually the longer bond length, and
therefore lower bond energy.
It is usually easier to break the bond between
two larger atoms than between two small atoms.
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Using Lewis structures we can begin to predict
the results of sharing valence electrons to
achieve the same configuration as the nearest
noble gas.
For instance how many bonds do you predict the
following nonmetals will make?
2
0
4
3
1
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Octet Rule
To be stable the two atoms that are involved in
the bond share their electrons in order to have 8
valence electrons
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COVALENT BONDS
lets look at the molecule cl2

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each atom must have 8 valence e's
Cl
Cl
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Another example
4 bonds possible
1 bond each
C is usually the central atom
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Another example
C is usually the central atom
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Another example
each atom must have 8 valence e's
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Another example
Hydrogen is happy with 2 electrons
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There are also examples of molecules where there
is more than a single pair of shared electrons
between two atoms.
Nitrogen has three spots for bonding, and shares
three pairs of electrons.
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Each nitrogen has a set of electrons that are not
available for bonding. Nitrogen can only make 3
bonds.
N
N
These triple bonds are very strong, and have a
very high bond energy. therefore it is a very
stable bond.
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The Lewis structures that you have seen so far
use a pair of dots to represent covalent bonds.
We can also use dashes to represent a pair of
bonding electrons.
There are three types of bonds, single, double,
and triple depending on how many electron pairs
are shared.
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Dots are used to show valence electrons that are
nonbonding electrons.
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We can write these structural formulas with
confidence by using these six steps for drawing
Lewis structures.
With ex CH2O

• Count the total valence electrons for the
  molecule (valence electrons for each atom and add
  them up)

C ? 4
H ?2 1
O ? 6
12 val es
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• Figure out how many octet electrons the molecule
  should have, using octet rule (everyone wants 8
  except hydrogen, which is happy with 2)

C ? 8
H ?2 2
O ? 8
20 octet es
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With ex CH2O

• Subtract the valence electrons from the octet
  electrons (these are the number of electrons
  involved in bonding)

20 octet es 12 val es 8
bonding es
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With ex CH2O

• Divide the number from 3 by two (these are the
  number of dashes or bonds in the molecule)

8 bonding es/2 4 bonds
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With ex CH2O

• Draw an arrangement of the atoms for the molecule
  that contains the of bonds you found in 4.
• Hydrogen and the halogens only form 1 bond
• Oxygen family forms 2 bonds
• Nitrogen family forms 3 bonds
• Carbon family forms 4 bonds and is usually the
  central atom

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Hint bond all the atoms together by single
bonds, and then add the multiple bonds until the
rules above are followed
C
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With ex CH2O

• Find the number of lone pair (nonbonding)
  electrons by subtracting the bonding electrons
  (3 above) from the valence electrons (1 above).
  Arrange these around the atoms until they
  satisfy the octet rule.

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Hint Hydrogen is happy with 2 electrons
12 val es - 8 bonding es
4 nonbonding es
?
?
?
C
?
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Hint Hydrogen is happy with 2 electrons
12 val es - 8 bonding es
4 nonbonding es
?
?
?
C
?
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So what we get is the structure of the molecule
from the molecular formula, we know how the
atoms are connected not just how many atoms are
involved.
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With ex C2H2

• Valence electrons

C ? 2 4
10 val es
H ?2 1

• Octet electrons

C ? 2 8
20 oct es
H ?2 2

• Of bonding electrons

20 oct es 10 val es
10 bonding es
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With ex C2H2

• Of bonds

10 bonding es/2
5 bonds

• Arrangement

• Nonbonding es

10 bond es 1o val es
0
?
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Sometimes, there is more than one possible Lewis
structure for the same molecule.
These are called resonance structures and we
usually draw the average structure
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Lewis structures can be drawn for the polyatomic
ions too
The only difference is you must alter the of
electrons to reflect the total charge of the
polyatomic ion.
H
H
H
N
O
H
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Lewis structures can be drawn for the polyatomic
ions too
The only difference is you must alter the of
electrons to reflect the total charge of the
polyatomic ion.

H
H
H
H
N
O
H
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Lewis structures can be drawn for the polyatomic
ions too
The only difference is you must alter the of
electrons to reflect the total charge of the
polyatomic ion.

H
-
H
H
H
N
O
H
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Naming molecular or covalent compounds is fairly
easy.
Its similar to naming binary ionic compounds,
except covalent compounds use prefixes.

• Prefixes to use
• Mono- (1)
• Di- (2)
• Tri- (3)
• Tetra- (4)
• Penta- (5)

• Hexa- (6)
• Hepta- (7)
• Octa- (8)
• Nona- (9)
• Deca- (10)

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A prefix is added to an element to tell how many
of that element are in the compound
We will work with binary compounds almost
exclusively when naming covalent compounds.
We then change the ending of the 2nd element in
the binary compnd to ide.
You need to follow a couple of simple rules
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Rule 1 Name the first nonmetal with the proper
prefix (if there is only one of the 1st element,
you do not use the prefix mono).
Rule 2 Name the second nonmetal with the
proper prefix and change the ending to ide.
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LETS PRACTICE
P2Br3 ? CH3 ? S2O6 ?
Diphosphorus tribromide
Carbon trihydride
Disulfur hexaoxide
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LETS PRACTICE
P2Br3 ? CH3 ? S2O6 ?
Diphosphorus tribromide
Carbon trihydride
Disulfur hexoxide
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How can you tell the shape of a molecule?
The Shape of a molecule cannot be predicted using
exclusively the molecular formulas.
Molecules with relatively simple structural
formulas tend to have relatively simple shapes.
The greatest affect on shape of the molecule is
any unshared electrons present in an atoms
valence set.
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Unshared electrons as well as the connected atoms
all tend to take up space.
They work to repel each other which pushes the
molecule into shapes and arrangements where the
atoms get as far from each other as possible.
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The theory that states that repulsion between
valence level electron pairs causes these pairs
to be oriented as far apart as possible is the
valence shell electron pair repulsion theory
(VSEPR).
You can begin to predict the basic geometry of a
molecule by looking at its structural formula
then accounting for the repulsion of the electron
clouds or domains of the attached atoms.
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