Transcript Covalent Bonding-compounds & Structures

Covalent bonding
Molecules & Structures
1
What do you already know
about Covalent bonding?
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Bonding Review
 What kind of bonding occurs when a metal
and a nonmetal transfer electrons?
– Ionic bonding
 What is made when two metals just mix and
don’t react?
– An alloy
 What do two nonmetals or some metalloids
with nonmetals form when they bond
together?
– Covalent bond
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Covalent bonding makes molecules
 Molecules are
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– Specific atoms, usually nonmetals, joined by
sharing electrons
Two major kinds of molecules:
Molecular compound
– Sharing e- between different elements
– Example: CH4
Diatomic molecules
– Sharing e- between two of the same atom
– These atoms occur naturally as compounds b/c
they are more stable that way
– Examples:
Diatomic elements
 There are 7 elements that always form
molecules
 H2 , N2 , O2 , F2 , Cl2 , Br2 and I2
 1 + 6 pattern on the periodic table
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1 and 6
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Properties of Molecular Compounds
 Tend to have low melting and boiling
points
 Have a molecular formula which
shows type and number of atoms in a
molecule
– Not necessarily the lowest ratio of
elements
 Ex: C6H12O6 or H2O
 The molecular formula doesn’t tell you
how bonded atoms are arranged
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How does H2 form?
 The nuclei repel
+
8
+
How does H2 form?
 The nuclei repel
 But they are attracted to electrons
 They share the electrons
 So the bond forms when the attractive
forces balance the repulsive forces
+
9
+
How do we show bonding?
 Lewis structures
– Use electron-dot diagrams to show
how electrons are arranged in
molecules
– Ex: H2
 These are called structural formulas
– They show what atoms are bonded
together and the type(s) of bond(s).
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Covalent bonds
 Nonmetals hold onto their valence
electrons.
 Need noble gas configuration to be stable
– Usually with 8 valence electrons
 Get it by sharing valence electrons with
each other.
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Covalent bonding
l
l
l
l
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons
Both end with full orbitals
F
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F
Covalent bonding
l
l
l
l
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons
Both end with full orbitals
F F
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8 Valence
electrons
Covalent bonding
l
l
l
l
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons
Both end with full orbitals
8 Valence
electrons
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F F
Single Covalent Bonds:
 A sharing of two valence electrons.
 Ex: H2
Double Covalent Bonds:
 A sharing of four valence electrons
 Ex: O2
Triple Covalent Bonds:
 A sharing of six valence electrons
 Ex: N2
As the number of shared electrons pairs increase, bond
length decreases
The shorter the bond length, the stronger the bond
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
Drawing Structural Formulas
1.
We use Lewis Structures to show this:
Predict the location of atoms
Terminal atoms will be an end atom on the structure
because they can only form one bond
-
-
2.
3.
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Ex: H & F ALWAYS
The central atom is usually the one that is less
electronegative
Make sure you have the correct number of valence e- for all of
your atoms
–
For polyatomic ions, add one e- for each negative charge
& subtract one e- for each positive charge
Start bonding by creating single bonds between the central
atom(s) and each of the terminal atoms
4.
5.
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Make sure to share electrons between atoms
as needed to get an octet on EACH atom.
– SOME EXCEPTIONS:
H wants only 2 e-, Be is OK with only 4 e-,
and B is OK with only 6 e– Remember C, N, O, & S can form double
or triple bonds with the same element or
another element if cannot get an octet with
only single bonds
Redraw with lines for each shared pair of
electrons (covalent bonds). Remember to
enclose a polyatomic ion in brackets and
indicate the overall charge on the ion.
Some things to think about:
 Too few e- for octets? Consider a
double or triple bond.
 Too many electrons? Can an atom
have an expanded octet? Any
atom in rows 3-7 of the periodic
table can have an expanded octet.
–Why? B/c of empty d-orbitals
–The atom with the expanded
octet is usually the central atom.
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 PH3
 H 2S
 CCl4
 SiO2
 NO320
Examples:
Practice
 Draw Lewis structures for the following:
 PCl3
 CH2O
 C3H6
 SO42-
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Practice on Molecules with
More than One Central Atom
 C 2H 2
 CH3COOH
 H 2O 2
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Dealing with
Exceptions to the Octet Rule
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Resonance Structures
 When more than one correct structure can be written
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for a molecule or ion.
– Usually happens with molecules or polyatomic
ions that have both double and single bonds.
They only differ in the position of the electron pairs
NOT the atom positions.
EX: NO2Which one is the true structure?
Does it go back and forth?
Double bonds are shorter than single
In NO2- all the bonds are the same length
The actual molecule behaves as if it only has one
structure.
Let’s Draw Some Resonance Structures:
 SO2
 SO32 O3
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Coordinate Covalent Bond
 When one atom donates both electrons
in a covalent bond.
 Ex: Carbon monoxide
CO
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Coordinate Covalent Bond
l
l
When one atom donates both electrons
in a covalent bond.
Ex: Carbon monoxide
C O
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Coordinate Covalent Bond
l
l
When one atom donates both electrons
in a covalent bond.
Ex: Carbon monoxide
C O
C
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O
Another Example:
 The ammonium ion (NH4+), which is
formed from the combination of
ammonia and a H+ ion.
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Molecules with odd numbers
of electrons
 Molecules when valence e- are counted
cannot form octets around each atom
 Examples: NO2, ClO2 and NO
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Don’t forget Expanded Octets
 The central atoms contain more than 8
valence electrons
 Ex: XeF4
 Ex:
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SF6
Bond Dissociation Energy
 The energy required to break a bond
 C - H + 393 kJ  C + H
 You ALWAYS have to add energy to break bonds (it
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will always be a positive number)
Double bonds have larger bond dissociation energies
than single
Triple bonds are even larger
– Examples of bond dissociation energy between
different types of carbon-carbon bonds
– C-C 347 kJ
– C=C 657 kJ
– C≡C 908 kJ
Bond Dissociation Energy
 The larger the bond energy, the harder it is to
break the bond
 Large bond energies make chemicals less
reactive
 In chemical reactions, bonds in the reactants
are broken and new bonds are formed to
make products.
 The total energy change of a chemical
reaction is determined from the energy of the
bonds broken and formed.
– This is where Endothermic and Exothermic
reactions come from
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Let’s Do An Example:
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First balance the equation for the
combustion of methane:
CH4 +
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O2 →
CO2 +
H2O
Draw the Lewis structures of the reactants
and products
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Refer to the structures and add up the energy released from
forming all the bonds in the products. (For example: if 2 moles
of water are formed, that’s 4 O-H bonds)
Add up the bond energies for the moles of the reactants.
Subtract the energy used to break all the bonds in the reactants
from the energy used to make the bonds in the products. This
is how much energy released in the combustion reaction.
Is the reaction endothermic or exothermic?
Bond Polarity
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Electronegativity and Polarity
 Covalent bonds are the sharing of electrons
between atoms.
 The amount of sharing can change
depending on how strongly an atom holds
onto its electrons.
 We use the periodic table and values of
electronegativity to determine how strongly
an atom will pull electrons in a bond.
 The electronegativity of the atoms was
assigned by Linus Pauling when he studied
the bonding abilities of atoms in molecules.
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 When the atoms in a bond are the same, the
electrons are shared equally.
 Also when there is little difference in
electronegativity, the electrons are essentially
shared equally
 These are considered nonpolar covalent
bonds.
 When two different atoms are connected, the
electrons may not be shared equally.
 This is a polar covalent bond.
 How do we measure how strong the atoms
pull on electrons?
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Electronegativity
 A measure of how strongly the atoms attract
electrons in a bond.
 The bigger the electronegativity difference the
more polar the bond.
 Use Figure 9-15 Pg. 263 in text to get
electronegativities of atoms.
 We use general guidelines to determine if a
bond is polar, nonpolar or ionic.
– Chemical bonds between different
elements are never completely ionic or
covalent
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 Use the following differences in
electronegativities of bonded atoms as
general guidelines for bond polarity:
 0.0 - 0.4 nonpolar covalent bond
 0.5 - 1.7 polar covalent bond
 >1.7 Ionic
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How to show when a bond is polar
 Isn’t a whole charge just a partial charge
 d+ means a partially positive
 d- means a partially negative
d+
H
d-
Cl
 The Cl pulls harder on the electrons
 The electrons spend more time near the Cl
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For each pair of elements, calculate the
electronegativity difference and label the
bond type (polar covalent, nonpolar
covalent, or ionic).
 H, Cl
 H, S
 S, Cl
 Na, F
 Cl, Br
 Al, Br
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Examples with compounds
 Let’s determine if the bonds in the following
compounds are polar, nonpolar, or ionic. You
will need to show your calculations!
 HCl
 CH4
 HSF
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Molecular Shapes
Some Theories
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VSEPR Theory
 Valence Shell Electron Pair Repulsion.
 Predicts three dimensional geometry of
molecules based on the number of pairs of
valence electrons both bonded and unbonded.
 Name tells you the theory.
 Valence shell - outside electrons.
 Electron Pair repulsion - electron pairs try to get
as far away as possible.
 Can determine both the angles of bonds
 AND the shape of molecules
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VSEPR
 The bond angle is the angle formed by two
terminal atoms
 Both shared pair e- and lone pair e- repel
each other
– Lone pair e- repel more than shared e-
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# of eShared Lone
domains e- pairs pair e-
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Molecular
shape
Bond
Angle
Example
2
2
0
Linear
180o
BeH2
3
3
0
Trigonal
planar
120o
BH3
4
4
0
Tetrahedral
109.5o CH4
4
3
1
Trigonal
pyramidal
107.3o NH3
3 or 4
2
1 or 2
Bent
104.5o H2O
5
5
0
Trigonal
6
6
0
bipyramidal
Octahedral
90o/
120o
90o
http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html
PF5
SF6
Examples of how we get the
molecular shapes
H
H C H
H
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 Single bonds fill
all atoms.
 There are 4 pairs
of electrons
pushing away.
 The furthest they
can get away is
109.5º.
4 atoms bonded
 Basic shape is
tetrahedral.
 A pyramid with a
triangular base.
 Same basic shape
for everything
with 4 pairs.
H
H
109.5º
C
H
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H
3 bonded - 1 lone pair
 Still basic tetrahedral but you can’t see
the electron pair.
 Shape is called
trigonal pyramidal.
H N H H
H
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N
H
H
<109.5º
2 bonded - 2 lone pair
 Still basic tetrahedral but you can’t see
the 2 lone pair.
 Shape is called
bent.
H O
H
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O
H
H
<109.5º
3 atoms no lone pair
 The farthest you can the electron pair
apart is 120º.
 Shape is flat and called
trigonal planar.
 Will require 1 double bond
H
H
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C O
H
H
C
120º
O
2 atoms no lone pair
 With three atoms the farthest they can
get apart is 180º.
 Shape called linear.
 Will require 2 double bonds or one triple
bond
180º
O C O
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Try to determine the shapes of
these molecules:
 SiH4
 PF3
 HBr
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How do we get the shapes of these?
 C 2H 2
 CH3COOH
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Molecular Orbitals (MO)
 The overlap of atomic orbitals from
separate atoms makes molecular
orbitals
 Each molecular orbital has room for two
electrons
 Two types of MO
– Sigma ( σ ) between atoms
– Pi ( π ) above and below atoms
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Sigma bonding orbitals
 From s orbitals on separate atoms
+
+
s orbital s orbital
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+ +
+ +
Sigma bonding
molecular orbital
Sigma bonding orbitals
 From p orbitals on separate atoms


p orbital


p orbital


Sigma bonding
molecular orbital
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Pi bonding orbitals






 P orbitals on separate atoms
Pi bonding
molecular orbital
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Sigma & Pi Bonds
 Sigma bonds (s) occur from overlap of
orbitals between the atoms
 Pi bond (p bond) occur between p orbitals.
above and below atoms
 All single bonds are
s bonds
 Double bond is 1 s
and 1 p bond
 Triple bond is 1 s
and 2 p bonds
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Hybrid Orbitals
Combines bonding with geometry
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Hybridization
 The mixing of different atomic orbitals to form the same
type of hybrid orbitals.
 All the hybrid orbitals that form are identical.
 Each hybrid orbital contains one electron that it can share
with another atom.
 The number of atomic orbitals mixed to form the hybrid
orbital equals the total number of pairs of electrons
(double and triple bonds get treated as though they are
one pair of electrons)
 Lone pairs on the central atom also occupy hybrid orbitals.
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Types of Hybridization
 sp3 -1 s and 3 p orbitals mix to form 4 sp3
orbitals.
 EX: CH4, NH3, H2O
 sp2 -1 s and 2 p orbitals mix to form 3 sp2
orbitals leaving 1 regular p orbital.
 EX: BH3, AlCl3, C2H4
 sp -1 s and 1 p orbitals mix to form 2 sp
orbitals leaving 2 regular p orbitals.
 Ex: BeCl2, CO2, O2, N2
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sp3 hybridizaiton
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sp2 hybridization
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Where is the p orbital?
 Perpendicular
 The overlap of
orbitals makes a
sigma bond (s
bond)
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H
H
C
H
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C
H
sp hybridization
 when two things come off
 one s and one p hybridize
 linear
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sp hybridization
 end up with two lobes 180º
apart.
 p orbitals are at right
angles
 makes room for two p
bonds and two sigma
bonds.
 a triple bond or two double
bonds
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CO2
 C can make two s and two p
 O can make one s and one p
O
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C O
N2
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N2
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sp3d
 1 s, 3 p, and 1 d orbitals mix together
making 5 sp3d hybrid orbitals
 Ex:
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PCl5
sp3d2
 1 s, 3 p, and 2 d orbitals mix together
making 6 sp3d2 hybrid orbitals
 Ex:
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SF6
Molecular Polarity
How to show if the entire molecule
is polar or not.
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Molecular Polarity
 Molecules are either nonpolar or polar,
depending on the location and nature of
the covalent bonds.
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Nonpolar Molecules
 There is symmetry with regard to the
distribution of electrons.
 Determine the shape!
 If there is an electronegative atom on
one part of the molecule and one that
“balances” it on another part, then the
molecule is nonpolar. If not, it is a polar
molecule
 Ex: CH4 and CCl4 and CH4Cl
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Polar Molecules
 Molecules with a partially positive end and a
partially negative end
 Symmetry can not cancel out the effects of
the polar bonds. (There is no “balancing” of
electronegative atoms on another part of the
molecule)
Must determine shape first.
Examples: H2O
and NF3
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For each molecule, draw the Lewis structure,
predict the shape and bond angle, and
identify as polar or nonpolar.

Br2

H 2S

HCN

PF3

C2H2

CH2O

NH4+

MgO
 “Symmetrical” shapes are those without
lone pair on central atom
– Tetrahedral
– Trigonal planar
– Linear
 The molecule will be nonpolar if all the
atoms are the same or have low
differences in electronegativities
 Shapes with lone pair on central atom
are not symmetrical
 Can be polar even with the same atom
bonded to the central atom
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Is it a polar or nonpolar molecule?
 HF
 H2O
 NH3
 CBr4
 CO2
 CH3Cl
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Properties of Molecules
 Most have LOW melting & boiling points

tend to be gases and liquids at room temperature
 Ex: CO2, NH3, H2O
 Polar and Nonpolar molecules have a
little bit different properties due to the
partial charges.
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-
+
d+ d-
H-F
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Properties of Solid Molecules
 Two kinds of crystals:
– Molecular solids – molecules held
together by attractive forces
• Ex: BI3, Dry Ice, sugar
– Network solids- atoms held together by
bonds
• One big molecule (diamond, graphite)
• High melting & boiling points, brittle, extremely hard
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Graphite
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Diamond
Intermolecular Forces
What holds molecules to each
other
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Intermolecular Forces
 They are what make solid and liquid
molecular compounds possible.
 The weakest are called van der Waal’s
forces - there are two kinds
– Dispersion forces
– Dipole Interactions
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Dispersion Force
 Depends only on the number of
electrons in the molecule
 Bigger molecules more electrons
 More electrons stronger forces
• F2 is a gas
• Br2 is a liquid
• I2 is a solid
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Dispersion force
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d+
d-
d+
d-
H
H
H
H
Dipole interactions
 Occur when polar molecules are
attracted to each other.
 Slightly stronger than dispersion forces.
 Opposites attract but not completely
hooked like in ionic solids.
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Dipole interactions
 Occur when polar molecules are
attracted to each other.
 Slightly stronger than dispersion forces.
 Opposites attract but not completely
hooked like in ionic solids.
+
d
d
H F
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+
d
d
H F
+
+
-
-
+
-
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Hydrogen bonding
 Are the attractive force caused by
hydrogen bonded to F, O, or N.
 F, O, and N are very electronegative so
it is a very strong dipole.
 They are small, so molecules can get
close together
 The hydrogen partially share with the
lone pair in the molecule next to it.
 The strongest of the intermolecular
forces.
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Hydrogen Bonding
d+ dH O
+
Hd
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Hydrogen bonding
H O
H
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Video lesson
 Water, a polar molecule, on YouTube:
https://www.youtube.com/watch?v=iOOvX
0jmhJ4
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Review
Ionic and Covalent Compounds
 Practice Quiz and Graphics:
http://www.elmhurst.edu/~chm/vchemboo
k/145Areview.html
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Internet resources
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Molecular polarity:
http://www.elmhurst.edu/~chm/vchembook/210polarity.html
Polar covalent compounds:
http://www.elmhurst.edu/~chm/vchembook/152Apolar.html
Nonpolar covalent compounds:
http://www.elmhurst.edu/~chm/vchembook/150Anpcovalent.
html
Ionic compounds:
http://www.elmhurst.edu/~chm/vchembook/143Aioniccpds.ht
ml
Compare Ionic, Polar, and Nonpolar Bonds:
http://www.elmhurst.edu/~chm/vchembook/153Acompare.ht
ml