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Chapter 8
Covalent bonding
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Covalent BondingA metal and a nonmetal transfer
electrons– An ionic bond
Two metals just mix and don’t react– An alloy
What do two nonmetals do?– Neither one will give away an electron– So they share their valence electrons– This is a covalent bond
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Covalent bondingMakes molecules
– Specific atoms joined by sharing electrons
Two kinds of molecules:Molecular compound
– Sharing by different elementsDiatomic molecules
– Two of the same atom– O2 N2
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Diatomic elementsThere are 8 elements that always form
molecules
H2 , N2 , O2 , F2 , Cl2 , Br2 , I2 , and At2
Oxygen by itself means O2
The –ogens and the –ines
1 + 7 pattern on the periodic table
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1 and 7
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Molecular compoundsTend 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
C6H12O6
Formula doesn’t tell you about how atoms are arranged
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Polar BondsWhen the atoms in a bond are the
same, the electrons are shared equally.
This is a nonpolar covalent bond.
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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ElectronegativityA measure of how strongly the atoms
attract electrons in a bond.The bigger the electronegativity
difference the more polar the bond.Use table 12-3 Pg. 2850.0 - 0.4 Covalent nonpolar0.5 - 1.0 Covalent moderately polar1.0 -2.0 Covalent polar>2.0 Ionic
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9Chapter 9: Chemical Bonds 9
Electronegativity
Electronegativity (EN) is a measure of the ability of an atom to attract bonding electrons to itself
EOS
The greater the electronegativity of an atom in a molecule, the more strongly it attracts the electrons in a covalent bond
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10Chapter 9: Chemical Bonds 10
Pauling’s Electronegativities
EOS
Electronegativity Illustrated
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11Chapter 9: Chemical Bonds 11
Electronegativity Differenceand Bond Type
Two identical atoms have the same electronegativity and share a bonding electron pair equally. This is called a nonpolar covalent bond
Example: chlorine gas
EOS
All homonuclear diatomic molecules have nonpolar covalent bonds:
H2, N2, O2, F2, Cl2, Br2, I2
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12Chapter 9: Chemical Bonds 12
Electronegativity Differenceand Bond Type
In covalent bonds between atoms with somewhat larger electronegativity differences, electron pairs are shared unequally. This is called a polar covalent bond
Example: hydrogen chloride gas, HCl
EOS
The electrons are drawn closer to the atom of higher electronegativity, Cl
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Covalent BondingElectrons are shared by atoms.
These are two extremes.
In between are polar covalent bonds.
The electrons are not shared evenly.
One end is slightly positive, the other negative.
Indicated using small delta
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How to show a bond is polar Isn’t a whole charge just a partial charge
means a partially positive
means a partially negative
The Cl pulls harder on the electrons
The electrons spend more time near the Cl
H Cl
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H - F+ -
H - F
+-H - F+
-
H - F
+-
H - F +-
H - F+-
H - F
+-
H - F
+-
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16 © 2009, Prentice-Hall, Inc.
Polar Covalent Bonds
When two atoms share electrons unequally, a bond dipole results.
The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:
= Qr
It is measured in debyes (D).
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17Chapter 9: Chemical Bonds 17
Electronegativity Differenceand Bond Type
With still larger differences in electronegativity, electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds
Example: sodium chloride, NaCl
EOS
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18Chapter 9: Chemical Bonds 18
Electronegativity Differences
EOS
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19Chapter 9: Chemical Bonds 19
Electron Distributions and Covalent Bonds
Symmetric distribution
EOS
AsymmetricDistribution
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Electronegativity difference
Bond Type
Zero
Intermediate
Large
Covalent
Polar Covalent
Ionic
Co
valent C
haracter
decreases
Ion
ic Ch
aracter increases
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How does H2 form?The nuclei repel
++
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How does H2 form?
++
The nuclei repel
But they are attracted to electrons
They share the electrons
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23Chapter 9: Chemical Bonds 23
The Lewis Theory ofChemical Bonding: An
OverviewElectrons, particularly valence electrons, play a fundamental role in chemical bonding
EOS
In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases
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24Chapter 9: Chemical Bonds 24
Lewis Symbols
Valence electrons are shown by dots around the element symbol
EOS
Use rules of electron configurations when forming dot structures …e.g., electrons remain unpaired if possible
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Covalent bondsNonmetals hold onto their valence
electrons.
They can’t give away electrons to bond.
Still need noble gas configuration.
Get it by sharing valence electrons with each other.
By sharing both atoms get to count the electrons toward noble gas configuration.
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Covalent bonding 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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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F8 Valence electrons
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Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals
F F8 Valence electrons
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Lewis StructureShows how the valence electrons are
arranged.One dot for each valence electron.A stable compound has all its atoms with
a noble gas configuration.Hydrogen follows the duet rule.The rest follow the octet rule.Bonding pair is the one between the
symbols.
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30Chapter 9: Chemical Bonds 30
Multiple Covalent BondsThe covalent bond in which one pair of electrons is shared is called a single bond
e.g., H : Cl or H—Cl
Double bonds have two shared pairs of electrons
EOS
Triple bonds have three shared pairs of electrons
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RulesSum the valence electrons.
Use a pair to form a bond between each pair of atoms.
Arrange the rest to fulfill the octet rule (except for H and the duet).
H2O
A line can be used instead of a pair.
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A useful equation(happy-have) / 2 = bonds
CO2 C is central atom
POCl3 P is central atom
SO42-
S is central atom
SO32-
S is central atom
PO43-
P is central atom
SCl2 S is central atom
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Exceptions to the octetBH3
Be and B often do not achieve octetHave less than an octet, for electron
deficient molecules.
SF6
Third row and larger elements can exceed the octet
Use 3d orbitals?
I3-
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Exceptions to the octetWhen we must exceed the octet, extra
electrons go on central atom.
(Happy – have)/2 won’t work
ClF3
XeO3
ICl4-
BeCl2
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35Chapter 9: Chemical Bonds 35
Writing Lewis Structures
Hydrogen atoms are terminal atoms (bonded to only one other atom)
EOS
The central atom of a structure usually has the lowest electronegativity and the terminal atoms (except H) generally have higher electronegativities
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Single Covalent BondA sharing of two valence electrons.
Only nonmetals and Hydrogen.
Different from an ionic bond because they actually form molecules.
Two specific atoms are joined.
In an ionic solid you can’t tell which atom the electrons moved from or to.
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How to show how they formed It’s like a jigsaw puzzle.
I have to tell you what the final formula is.
You put the pieces together to end up with the right formula.
For example- show how water is formed with covalent bonds.
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Water
H
O
Each hydrogen has 1 valence electron
and wants 1 more
The oxygen has 6 valence electrons
and wants 2 more
They share to make each other “happy”
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WaterPut the pieces together
The first hydrogen is happy
The oxygen still wants one more
H O
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WaterThe second hydrogen attaches
Every atom has full energy levels
H OH
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Multiple BondsSometimes atoms share more than one
pair of valence electrons.
A double bond is when atoms share two pair (4) of electrons.
A triple bond is when atoms share three pair (6) of electrons.
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Carbon dioxideCO2 - Carbon is central
atom ( I have to tell you)
Carbon has 4 valence electrons
Wants 4 more
Oxygen has 6 valence electrons
Wants 2 moreO
C
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Carbon dioxideAttaching 1 oxygen leaves the oxygen 1
short and the carbon 3 short
OC
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Carbon dioxide Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2 short
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more
OCO
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the
bond
OCO
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the
bond
OCO8 valence electrons
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the
bond
OCO8 valence electrons
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Carbon dioxide The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in the
bond
OCO
8 valence electrons
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How to draw themTo figure out if you need multiple bondsAdd up all the valence electrons.Count up the total number of electrons to
make all atoms happy.Subtract.Divide by 2Tells you how many bonds - draw them.Fill in the rest of the valence electrons to
fill atoms up.
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ExamplesNH3
N - has 5 valence electrons wants 8
H - has 1 valence electrons wants 2
NH3 has 5+3(1) = 8
NH3 wants 8+3(2) = 14
(14-8)/2= 3 bonds
4 atoms with 3 bonds
N
H
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N HHH
ExamplesDraw in the bonds
All 8 electrons are accounted for
Everything is full
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ExamplesHCN C is central atom
N - has 5 valence electrons wants 8
C - has 4 valence electrons wants 8
H - has 1 valence electrons wants 2
HCN has 5+4+1 = 10
HCN wants 8+8+2 = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds -will require multiple bonds - not to H
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HCNPut in single bonds
Need 2 more bonds
Must go between C and N
NH C
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HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add
NH C
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HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet
NH C
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Where do bonds go?Think of how many electrons they are
away from noble gas.
H should form 1 bond- always
O should form 2 bonds – if possible
N should form 3 bonds – if possible
C should form 4 bonds– if possible
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PracticeDraw electron dot diagrams for the
following.
PCl3
H2O2
CH2O
C3H6
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Another way of indicating bonds
Often use a line to indicate a bond
Called a structural formula
Each line is 2 valence electrons
H HO =H HO
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Structural Examples
H C N
C OH
H
C has 8 electrons because each line is 2 electrons
Ditto for N
Ditto for C here
Ditto for O
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Coordinate Covalent BondWhen one atom donates both electrons
in a covalent bond.
Carbon monoxide
CO
OC
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Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OC
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Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OCOC
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How do we know ifHave to draw the diagram and see what
happens.
Often happens with polyatomic ions
If an element has the wrong number of bonds
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Polyatomic ionsGroups of atoms held by covalent
bonds, with a charge
Can’t build directly, use (happy-have)/2
Have number will be different
Surround with [ ], and write charge
NH42+
ClO21-
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ResonanceWhen more than one dot diagram with
the same connections is possible.Choice for double bond
NO2-
Which one is it?Does it go back and forth?Double bonds are shorter than single
In NO2- all the bonds are the same
length
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Formal ChargeFor molecules and polyatomic ions that
exceed the octet there are several different structures.
Use charges on atoms to help decide which.
Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.
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Formal ChargeThe difference between the number of valence electrons on the free atom and that assigned in the molecule or ion.
We count half the electrons in each bond as “belonging” to the atom.
SO4-2
Molecules try to achieve as low a formal charge as possible.
Negative formal charges should be on electronegative elements.
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Formal Charge
EOS
Formal charge is the difference between the number of valence electrons in a free (uncombined) atom and the number of electrons assigned to that atom when bonded to other atoms in a Lewis structure
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Formal Charge
Usually, the most plausible Lewis structure is one with no formal charges
When formal charges are required, they should be as small as possible and negative formal charges should appear on the most electronegative atoms
EOS
Adjacent atoms in a structure should not carry formal charges of the same sign
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ExamplesXeO3
NO43-
SO2Cl2
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Resonance It is a mixture of both, like a mule.
CO32-
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Bond Dissociation EnergyThe energy required to break a bond
C - H + 393 kJ C + H
Double bonds have larger bond dissociation energies than single
Triple even larger
– C-C 347 kJ
– C=C 657 kJ
– C≡C 908 kJ
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Bond Dissociation EnergyThe larger the bond energy, the harder
it is to break
Large bond energies make chemicals less reactive.