Matter takes many forms in nature In this chapter we are going to learn to distinguish the type of compound that we have already studied the ionic compound which contains oppositelycharged particles metal ID: 700413
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Slide1
Chapter 8: Covalent Bonding
Matter takes many forms in nature
: In this chapter, we are going to learn to distinguish the type of compound that we have already studied, the “
ionic
compound
” (which contains oppositely-charged particles: metal
cations
and non-metal anions), from a different type of compound – a “
molecular
compound
”. Additionally, we are going to focus on a type of molecular compound known as a
binary
molecular compoundSlide2
II. Binary compounds
:
A “binary” compound contains
atoms
from
two
different elements.
A.
NaCl
”, “CaF
2
”, and “Al
2
O
3
” (3 ionic compounds) are binary
ionic
compounds. “NH
4
Cl” is an ionic compound, but because it contains more than 2 elements, it is
not
a binary ionic compound.Slide3
B. “N
2
O
5
”, “SF
6
”, and H
2
O” (3 molecular compounds) are binary molecular compounds.
“C
6
H
12
O
6
” is a molecular compound, but because it contains more than 2 elements it is
not
a binary molecular compound.Slide4
Comparison/Contrast between an ionic compound and a molecular substance.
A. Molecular substances are made of
molecules
.
1. There is no “molecule” in an ionic compound.Slide5
B. A
“molecule” contains a specific
number
of atoms, connected in a specific manner, to give a specific shape.
If even one atom is “missing” or “different”, the molecule would be an entirely different substance. Slide6
1. Not so with an ionic compound: In an ionic compound there is a specific
ratio
of atoms.. In salt (
NaCl
) for example, there is a ratio of 1 Na for every 1 Cl.
If a clump of salt lost 1 Na and 1
Cl
, it would still be the same original substance:
NaClSlide7
C. The formula of a “molecule” should never be simplified. C
2
H
8
is not the same substance as CH
4
.
1. The formula of an ionic compound should always be simplified. Ba
2
O
2
is the same substance as BaO.
D. A molecule will
not
crack apart.
1. Ionic compounds can crack apart if hammered….. If the cations come close too close together and the anions come too close together, the structure cracks apart.Slide8
E. Molecular substances may have
low
melting points.
1. All ionic compounds have a very high melting point.
F. Molecular substances may, at room temperature, be found as
solids, liquids,
or
gases.
1. All ionic compounds are solids (at room temperature).Slide9
G. Molecular substances contain atoms which are held together by
covalent
bonds.
1. Ionic substances are held together by ionic bonds. Slide10
IV. Covalent bonding:
The type of bonding that occurs within a molecular substance, in which atoms
share
their valence electrons in order to become more stable.
A. Occurs between
atoms
of
nonmetallic
elements.Slide11
B. Not all “molecules” or “molecular substances” are compounds!
In addition to the binary molecular compounds that we will study, there are 7 nonmetallic elements found in nature (in their elemental form) as pairs of atoms. These are the 7 “
diatomic
” elements:
N
2
, O
2
, F
2
, Cl
2
, Br
2
, I
2
, H
2
.Slide12
V. An important review:
A. Metallic elements: Found to the
left
side of the staircase boundary on the periodic table.
Non-metallic elements
:
Elements found to the
right
side of the staircase boundary on the periodic table.
2.
Hydrogen
is a nonmetallic element also.Slide13
VI. The octet rule:
When a molecule is formed: “
Nonmetal atoms
share
electrons in covalent bonds in order to obtain a full
octet
of electrons.” An octet = 8 valence electrons.
A. Exception: A hydrogen atom will end up with a total of
two
electrons by sharing with 1 other atom. Slide14
B. There are a few other notable
exceptions
to the octet rule:
1. A few molecular compounds which contain an odd number of valence electrons are known to exist.
2. A few molecular compounds have either a boron or an aluminum atom with 6 valence electrons.Slide15
2. A
few molecular compounds have a central atom with 10 or 12 valence electrons.
(1) One common example is “
sulfur hexafluoride
”.
In this compound, the central sulfur atom contains 6 x 2 =
12
valence electrons. Be sure to remember that this compound is an example in which the
central
atom does
not
follow the octet rule
.Slide16
VII. Types of covalent bonds.
A.
Single
covalent bond
–
1 shared pair of valence electrons:
2 dots, or a single dash
, represent 2 electrons that are simultaneously being attracted by, or “shared” by, the nuclei of two neighboring atoms. Slide17
1. The
formula in the center is a type of structural formula called a
“
Lewis dot structural formula
”.
The formula on the right is the
molecular
formula.
H
–
H
H
H
H
2
Slide18
B.
Double
covalent bond
–
two pairs of shared valence electrons: 4 dots or 2 parallel dashes
.
C C C C C
2
H
4
H H
H H
H H
H HSlide19
C.
Triple
covalent
bond
–
three pairs of shared valence electrons: 6 dots or 3 parallel dashes
.
N N or N N N
2
Notice the
two
“
unshared
pairs
” of electrons (one pair is to the far left and one pair to the far
righ
)t of the nitrogen structure.
You may never use a long dash to represent an unshared pair of electrons
.Slide20
Unshared pairs of electrons don’t bond the atoms together….but, the repulsive forces of unshared pairs of electrons do dramatically influence the
shape
of a molecule! Slide21
D. Notice how an ion can react with a molecule to generate a polyatomic ion. In the example below, a hydrogen ion bonds to a molecule of ammonia(NH
3
) to make the
ammonium ion (NH
4
)
+
:
H
+
+ N H
H
N H
H
H
H
H
+
]
[Slide22
VIII. Drawing a Lewis Dot Structure:
A. Certain elements are known as “central” atoms…. They will be found in the center of a structure. The first element given in a formula is usually the central atom (exception: hydrogen and the halogens).
1. Position the central atom in the center of your work space.Slide23
B. Hydrogen
and the halogens are known as “peripheral” atoms. They will be found only connected to one other atom.
Position
hydrogen and halogen atoms so that they “touch”, or “go around” only 1 other atom.Slide24
C. Add up all the valence electrons. Position the valence electrons as dots around the atom they belong to - the valence electrons may never leave the original atom.
Position the dots to form a “doorway” with 4 sides, in which the symbol of the element appears centered in the doorway.
Start with no more than 2 dots on each side of the 4 sided doorway.Slide25
D. If you can’t easily achieve a Lewis dot structure which has each atom (other than hydrogen) surrounded by 8 dots by doing what is described above, then you either need a double bond (2 pairs of shared electrons) or a triple bond (3 pairs of shared electrons).
For CO
2
, you will need two double bonds.Slide26
1. To make a double bond, move
one
“
un
-shared electron”
simultaneously from each of two neighboring atoms
, and place those 2 electrons in between the two neighboring atoms.
2. To make a triple bond, start with a double bonded pair of atoms, and simultaneously move
one more
unshared electron
from each of the two
atoms. Reposition those two electrons in between the atoms.Slide27
Important points regarding nonmetal atoms and their bonding charcteristics:Slide28
Atoms of the Following Nonmetallic Elements
:
Have This Number of Electrons (dots)
when
in a Stable Structure:
Type of Bonds Permitted by these atoms:
Things to Remember about atoms of these elements; or to remember about a specified molecule.
Hydrogen
2
Single covalent
Diatomic element;
H
2
is a “linear” diatomic
molecule
Boron, Aluminum
8 or 6
Single covalent
When only 6 electrons surround a boron or an aluminum atom, the molecule’s shape will be “
trigonal
planar” (a flat pancake).
Slide29
Atoms of the Following Nonmetallic Elements
:
Have This Number of Electrons (dots)
when
in a Stable Structure:
Type of Bonds Permitted by these atoms:
Things to Remember about atoms of these elements; or to remember about a specified molecule.
Sulfur
8 EXCEPT with
“SF
6
” when there are 12
Single, double, and/or triple covalent
In sulfur hexafluoride sulfur does NOT follow the octet rule. This is one “
exception”
to the octet rule.
The Halogen family: F, Cl, Br, I, At
8
Single covalent
All are diatomic elements; and,
F
2
, Cl
2
, Br
2
, I
2
, At
2
are all linear molecules, with only single bonds.Slide30
Atoms of the Following Nonmetallic Elements
:
Have This Number of Electrons (dots)
when
in a Stable Structure:
Type of Bonds Permitted by these atoms:
Things to Remember about atoms of these elements; or to remember about a specified molecule.
The Noble Gas family:
8….Except for helium (2). Noble gas atoms don’t form compounds.
Do NOT form compounds easily (no bonds).
Always found as single atoms in the gaseous state.
Nitrogen and oxygen
8
Single, double, and/or triple covalent
Diatomic elements; linear molecules. N
2
has one triple bond, while O
2
has 1 double bond.
All other nonmetal atoms
8
Single, double, and/or triple covalent
Slide31
IX. Lewis Dot structural formulas for polyatomic ions:
A. Covalent bonds occur
within
a polyatomic ion (
not
between
polyatomic ions).
B. When
drawing polyatomic ions, place the
first
element in the
center
of the structure, and place the second element
around
the first element (placing 1 atom of the second element along each different
side
of the first element).Slide32
C. When the charge of a polyatomic ion is +, you need to
subtract
the indicated number of electrons from the total of the valence electrons in the molecule. So, for +1 ions: take away
1
electron from the molecular ion’s number of valence electrons.
D. When the charge of a polyatomic ion is –, you need to
add
the indicated number of electrons to the molecular ion’s number of valence electrons. Slide33
1
. If the charge is 1
-
, then add
1
more electron to the molecule’s total number of valence electrons.
2. If the charge is 2
-
, then add
2
more electrons; if the charge is 3
-
, then add
3
more electrons.
E. Last, for a polyatomic ion: Draw a large
bracket
around the ion; and, place its charge at upper right.Slide34
# of N valence electrons: 5 5
# of H valence electrons: 4 x 1 = 4 4
Charge of ion = +1, therefore less 1 -1
Therefore, total = 8
Ammonium ion
(NH
4
)
+
H N H
H
H
+
]
[Slide35
Steps for Dot Structures:
Step 1: total # valence electrons.
Step 2. Position central atoms:
carbon atoms form a straight line;
assume only single bonding.Slide36
Step 3. Position other atoms; remember “special” molecules.
A. Peripheral atoms:
Hydrogen and the halogens-
connect to only 1 other atom
,
use only 1 single bond
.
B. binary polyatomic ions: first element is central, second element is peripheral.
assume all single bonds.
C. hydrocarbons” – molecular formula gives list of atoms (from left to right) connecting to each central atom (usually carbon)
CH
3
CH
2
OH means “first carbon touches 3 H atoms, second carbon touches 2 H atoms, then there is an O touching an H.
Assume all
single bonds
.
D. Memorize
: CO
2
C in the middle;
use two double bondsSlide37
Step 4: Make each atom stable.
Work from left to right:
Assume all single bonds.
Position unshared pairs to provide octets.
Exception: Hydrogen atoms = only 2 dots.
Step 5: Count the dots you’ve used.
Make sure the # you used = the # you were supposed to. Erase extras.Slide38
Step 6: Make corrections -
If your structure “needs” 2 extra dots, it really needs a double bond….
Erase 2 unshared dots, and share them
(as part of a double bond).
If your structure “needs” 4 extra dots, it really needs a triple bond.....
Erase 4 unshared dots, and share them (as part of a triple bond).
Every atom should now be stable.Slide39
X. VSEPR Theory –
V
alence
S
hell
E
lectron
P
air
R
epulsion theory
. [Remember: Like charges repel!]
A. A theory to
predict
the 3-dimensional geometry,
ie
.
the“
shape
” of a moleculeSlide40
1. The theory is based on “electrostatic repulsion”: Molecules will adjust their
shape
to keep the negatively-charged pairs of
valence
electrons as
far apart
as possible from each other.
B. When NOT to use
VSEPR
theory:
When there are only 2 atoms in a molecule. These molecule’s shapes are called
linear
– it doesn’t matter if there are single bonds, double bonds, triple bonds, or unshared electron pairs.Slide41
C. Using VSEPR theory:
1. Draw the Lewis
dot structure
for the molecule.
2. Identify its
central
atom.
3. Identify the sets of valence electrons as one of two possibilities
:
A. Those
connecting
two atoms.
B. Those that do not connect two atoms
. These are
called “
unshared
pairs”.Slide42
4. The unshared pairs found on a central atom strongly
repel
each other; and molecules that would otherwise be linear, will be forced into a
bent
(or
angular
) shape
.
5. Unshared pairs also cause a molecule that would be shaped like a flat triangle
(
trigonal
planar
), to be forced into a
not flat
(
trigonal
pyramidal
) shape.Slide43
6.
Count
the number of connections
separately
from the number of unshared pairs
.
1 single bond counts as
1
connection.
1 double bond counts as
1
connection.
1 triple bond counts as
1
connection.
Each unshared
set of 2 dots
counts as
1
un-shared pair.Slide44
D Predicting Shapes Using VSEPR Table
Read horizontally across the table.
Connections To
the Central Atom
Unshared Pairs of Electrons Around Central Atom
Molecular Shape
Around Central Atom
2
0
Linear
3
0
Trigonal
Planar
4
0
Tetrahedral, 109.5
o
2
1 or 2
Bent
3
1
PyramidalSlide45
Shapes:
,
E.
Shap
e
s:
Linear diatomic
Linear triatomic
Trigonal Planar
Bent
Pyramidal
Linear diatomic
Tetrahedral
Linear triatomic
Trigonal planar
bent
Trigonal pyramidal
tetrahedralSlide46
“Molecular Polarity”
– A term that is used to distinguish two types of molecules…. Based on the presence or absence of a
separation
of
charge
.
Some
molecules
show
characteristics indicating that they have
oppositely
-charged ends (a positive end and a negative end). This is called a separation of the charges (or “separation of charge”).Slide47
Other molecules
show characteristics indicating that their structure doesn’t have a separation of charge, or their
structure
hides
the presence of their oppositely-charged ends.Slide48
How to determine a molecule’s polarity.
The first part of determining a molecule’s polarity is to calculate each individual bond’s polarity.
Be careful with the vocabulary being used –
An
individual
bond’s polarity is called the “
bond polarity”
The polarity of the entire molecule is called the “
molecular polarity”Slide49
To calculate a bond polarity, first identify the “
electronegativity
value” of each of the 2 atoms in the bond you are working on.
The
electronegativity value
is number (from 0 to 4) which informs us
of an atom’s ability to
attract
electrons when in a compound
. Slide50
The electronegativity value is given on your periodic table, side 2, within each element’s square…..upper right corner of the square, in black print.
The closer an element’s electronegativity is to “
4
, the better that an atom of that element will attract electrons when that atom is found in a compound
.
The closer an element’s electronegativity is to “
0
”, the
less
likely it is for that atom to be able to attract electrons when in a compound.Slide51
After identifying the two electronegativity values, subtract the two electronegativity values. Take the absolute value of the answer (make the answer positive). Slide52
After subtracting and taking the absolute value of the two electronegativity values,
then
think about your answer,
and determine whether your answer indicates that there is, or that there is not, a situation in which one of the two atoms in the bond “overpowers” the other atom in terms of electron attracting ability
. If an atom is able to overpower the other atom, it will “hog” the electrons, as opposed to sharing the electrons equally with the other atom. You will see (on the next page) that I’ve placed a “δ–”sign next to an oxygen atom, and a “δ+” sign next to 2 hydrogen atoms, in a sketch of a water molecule. By doing this, I am indicating that the oxygen atom is hogging the negatively-charged valence electrons belonging to the 2 hydrogen atoms; and, the 2 hydrogen atoms are both overpowered and “partially lose” their own valence electrons to the oxygen atom.Slide53
δ
-
indicates the “partial negative” atom; and,
δ
+ indicates the “partial positive” atom
. Slide54
Finally, you are now able to conclude that a bond is either “nonpolar covalent” or “polar covalent”.
A bond is to be called “
non-polar
covalent
” when
the two atoms share the electrons more-or-less
equally
.
Labeling a bond as nonpolar covalent means that
the difference in electronegativity values fell between 0 and 0.4
Labeling a bond as nonpolar covalent means that
the electrons are distributed practically equally along the bond
.
Labeling a bond as nonpolar covalent means
that there is no separation of charge
in that particular bond.Slide55
A bond is to be called
“
polar
covalent
” when the electrons are NOT distributed equally along the bond, rather,
the electrons are found much or most of the time toward the atom that has the higher
electronegsativity
value
.
This will occur when difference in electronegativity values is greater than 0.4, but less than 2.
We place a δ – next to an atom to identify it as the “partial negative” atom; and, we place a δ + next to an atom to identify it as the “partial positive” atom in a polar covalent bond.Slide56
A bond is designated as
ionic
when one atom has stripped the other of some of its valence electrons, yielding a
cation
and an anion. Then, the
cation
and anion are attracted, and more
cations
and anions are attracted and an ionic compound forms.
This should occur when the difference in electronegativity values is greater than 2.
We studied this type of compound previously. Ionic compounds have very different characteristics than compounds in which polar covalent and nonpolar covalent bonds exist.Slide57
The second part of determining the “molecular polarity” (a
molecule’s
polarity
)
is visualizing the effect of the individual bond polarities in conjunction with the
shape
of the molecule.
Some common occurrences are listed below:
If the molecule’s shape is either
tetrahedral
,
linear triatomic
or
trigonal
planar
and if
all
the bonds in the structure are
identical
, then the overall molecule’s polarity (the molecular polarity) is
nonpolar
Slide58
This is because, if every bond is nonpolar, there isn’t a substantial separation of charge to begin with.
And, if every bond is polar (as long as each bond
has the same 2 atoms
), the polarity of each bond will be
cancelled
due to the symmetry of the polar bonds around a central atom. When this symmetry exists, the
δ
– charge(s)
cancel
the
δ
+ charge(s) in each bond.Slide59
If the molecule’s shape is either
tetrahedral ,
trigonal
planar,
or
linear
, and the structure contains
non
-identical bonds in which 1 is polar, then the overall (molecular) polarity is polar (because 1 polar bond won’t have another bond to be cancelled with).
If the molecules’ shapes are pyramidal or bent, and one or more of the bonds is polar, the overall (molecular) polarity is “
polar
”. This is because in molecules having these shapes, the bond polarity is not situated symmetrically relative to the molecule’s center, so the polarity doesn’t cancel.Slide60
Polar bonds
NonPolar molecule
Polar bonds Polar molecule.
Polar bonds combine to cancel out; Polar bonds do
NOT
cancel out;
Producing
Non-polar molecules
: Producing
Polar Molecules
:Slide61
“Like dissolves like”:
An expression stating that
ionic
compounds and molecules whose molecular polarity is
polar
will be able to dissolve only in solvents containing
polar
molecules; and, conversely molecules whose molecular polarity is
nonpolar
will be able to dissolve only in solvents containing
nonpolar
molecules
.Slide62
Some characteristics of a water molecule
:
A
H
2
O
molecule has 2 identical polar bonds, a bent structure, and its molecular polarity is “polar”.
The partial negative (or slightly negative) region of a water molecule is the area closest to the
oxygen
atom.
The partial positive (or slightly positive) region of a water molecule lies within the area closest to the two
hydrogens
atoms.Slide63
Because of having the above characteristics, water behaves in an unusual manner:
When placed between a positively charged metal plate and a negatively charged metal plate, the water molecules all line up, with their
positively
-charged region attracted toward the negative plate and their
negatively-
charged region attracted toward the positive plate.Slide64
Inter
- molecular Bonds (or forces of attraction)
are DIFFERENT from the covalent bonds you have been studying so far!
The covalent bonds you have been studying so far (single, double, triple bonds) are bonds
within
a molecule; these would be called______
intra-molecular
bonds.
Inter
-molecular bonds (forces) are the attractions
between
2 molecules
.
They hold
separate
molecules together.Slide65
They are
weaker
than polar and non-polar covalent bonds (single/double/triple).
They are
weaker
than the ionic bonds which connect
cations
to anions.
Breaking an inter-molecular bond is a
physical
change; whereas, breaking an intra-molecular bond is a chemical change.Slide66
When you boil water (to generate water vapor), or melt ice (to generate liquid water), you are breaking inter-molecular bonds. When you treat a water molecule with electricity, you destroy the water molecule and generate from it oxygen gas and hydrogen gas….this is the breaking of intra-molecular bonds….this is a chemical change.Slide67
Types of intermolecular bonds:
Hydrogen
bond: Is the
strongest
intermolecular force; and, it is perhaps the most important intermolecular bond, as it is necessary for life as we know it.
Hydrogen bonds take effect when you have a combination of a few certain atoms in a
polar molecule
.
Slide68
The
polar
molecule must contain a
hydrogen
atom, and that hydrogen atom must be connected to one of the highly electronegative atoms listed below:
Fluorine, oxygen
,
OR
nitrogen
.
Ex: H
2
O, HF, NH
3
Slide69
Hydrogen bonds determine the properties of water and biological molecules (such as proteins).
Cause water to predominate as a
liquid
(rather than as a gas) on earth.
Cause ice to
expand
upon freezing (rather than contract as the kinetic molecular theory would predict).
Holds the
DNA
double helix structure together.Slide70
Time-permitting, we will also learn about one of the weaker types of intermolecular forces of attraction by doing a lab in which we use evaporation rates to identify between hydrogen bonding and a weaker force, present between nonpolar molecules.