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Transcript of 10/12/2014Dr Seemal Jelani 1. 10/12/2014Dr Seemal Jelani 2.
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04/11/23Dr Seemal Jelani 1
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When two or more atoms of the same or similar electronegativities react, a complete transfer of electrons does not occur
In these instances the atoms achieve noble gas configurations by sharing electrons
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Covalent bonds Form by sharing of electrons
between atoms of similar electronegativities to achieve the configuration of a noble gas.
Molecules are composed of atoms joined exclusively or predominantly by covalent bonds
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Molecules may be represented by electron-dot formulas or, more conveniently, by bond formulas where each pair of electrons shared by two atoms is represented by a line.
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Types of electron
• Electrons of an atom are of two types:
• Core electrons and Valence electrons
• Only the valence electrons are shown in Lewis dot-line structures
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• The number of valence electrons is equal to the group number of the element for the representative elements
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How to distribute electrons For atoms the first four dots are
displayed around the four “sides” of the symbol for the atom.
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• If there are more than four electrons, the dots are paired with those already present until an octet is achieved
• Ionic compounds are produced by complete transfer of an electron from one atom to another
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• Covalent compounds are produced by sharing of one or more pairs of electrons by two atoms.
•The valence capacity of an atom is the atom’s ability to form bonds with other atoms.•The more bonds the higher the valence.
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The valence of an atom is not fixed, but some atoms have typical valences which are most common:Carbon: valence of 4Nitrogen: valence of 3 Oxygen: valence of 2 Fluorine: valence of 1
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Covalent bonding and Lewis structures• Covalent bonds are formed from
sharing of electrons by two atoms• Molecules possess only covalent
bonds• The bedrock rule for writing Lewis
structures for the first full row of the periodic table is the octet rule for C, N, O and F
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• C, N, O and F atoms are always surrounded by eight valence electrons
• For hydrogen atoms, the doublet rule is applied: H atoms are surrounded by two valence electrons.
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The Lewis Model of Chemical BondingThe Lewis Model of Chemical Bonding
In 1916 G. N. Lewis proposed that atomscombine in order to achieve a more stableelectron configuration.
Maximum stability results when an atomis isoelectronic with a noble gas.
An electron pair that is shared between two atoms constitutes a covalent bond.
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Covalent Bonding in H2
Covalent Bonding in H2
Sharing the electron pair gives each hydrogen an electron configuration analogous to helium.
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HH .. HH..
Two hydrogen atoms, each with 1 electron,,
can share those electrons in a covalent bond..
HH :: HH
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Covalent Bonding in F2Covalent Bonding in F2
Sharing the electron pair gives each fluorine an electron configuration analogous to neon.
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Two fluorine atoms, each with 7 valence Two fluorine atoms, each with 7 valence electronselectrons
can share those electrons in a covalent bond.can share those electrons in a covalent bond.
....
....FF.. FF..:: ::
....
....
FF :: FF:: ::........
....
....
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The Octet RuleThe Octet Rule
The octet rule is the most useful in cases involving covalent bonds to C, N, O, and F.
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FF :: FF:: ::........
....
....
In forming compounds, atoms gain, lose, In forming compounds, atoms gain, lose, or share electrons to give a stable or share electrons to give a stable electron configuration characterized by 8 electron configuration characterized by 8 valence electrons.valence electrons.
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ExampleExampleExampleExample
CC ........
FF::..........
Combine carbon (4 valence electrons) Combine carbon (4 valence electrons) and four fluorines (7 valence electrons and four fluorines (7 valence electrons each)each)
to write a Lewis structure for CFto write a Lewis structure for CF44..
:: FF::........CC
:: FF::........
:: FF::........:: FF::
....
....
The octet rule is satisfied for carbon and The octet rule is satisfied for carbon and each fluorine.each fluorine.
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ExampleExampleExampleExample
It is common practice to represent a covalentIt is common practice to represent a covalentbond by a line. We can rewritebond by a line. We can rewrite
:: FF::........CC
:: FF::........
:: FF::........:: FF::
....
....
....
CCFF
FF
FF
FF
....
............:: ::
:: ::
:: ::
....
asas
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Double Bonds and Triple Bonds
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Inorganic examplesInorganic examplesInorganic examplesInorganic examples
CC:: :: ::OO....
::OO....
:: :: CC ::OO....
OO....
::
:: :: ::NN::CC::HH ::NNCCHH
Carbon dioxideCarbon dioxide
Hydrogen cyanideHydrogen cyanide
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Organic examplesOrganic examplesOrganic examplesOrganic examples
EthyleneEthylene
AcetyleneAcetylene:: :: ::CC::CC::HH HH CCCCHH HH
CC:: ::CC....
HH :: ::....
HHHHHH
CC CC
HH HH
HHHH
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Formal Charges Revision
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Formal charge is the charge calculated for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs.
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Nitric acidNitric acid
We will calculate the formal charge for each atom in this Lewis structure.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of HFormal charge of H
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Nitric acidNitric acid
Hydrogen shares 2 electrons with oxygen.Assign 1 electron to H and 1 to O.A neutral hydrogen atom has 1 electron.Therefore, the formal charge of H in nitric
acid is 0.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of HFormal charge of H
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Nitric acidNitric acid
Oxygen has 4 electrons in covalent bonds.Assign 2 of these 4 electrons to O.Oxygen has 2 unshared pairs. Assign all 4
of these electrons to O.Therefore, the total number of electrons
assigned to O is 2 + 4 = 6.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of OFormal charge of O
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Nitric acidNitric acid
Electron count of O is 6. A neutral oxygen has 6 electrons. Therefore, the formal charge of O is
0.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of OFormal charge of O
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Nitric acidNitric acid
Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons).
A neutral oxygen has 6 electrons. Therefore, the formal charge of O is 0.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of OFormal charge of O
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Nitric acidNitric acid
Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons).
A neutral oxygen has 6 electrons. Therefore, the formal charge of O is -1.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of OFormal charge of O
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Nitric acidNitric acid
Electron count of N is 4 (half of 8 electrons in covalent bonds).
A neutral nitrogen has 5 electrons. Therefore, the formal charge of N is
+1.04/11/23Dr Seemal Jelani 31
.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal charge of NFormal charge of N
––
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Nitric acidNitric acid
A Lewis structure is not complete unless formal charges (if any) are shown.
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.... ::
....HH OO
OO
OO
NN
::
::....
....
Formal chargesFormal charges
––
++
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Formal ChargeFormal ChargeFormal ChargeFormal Charge
Formal charge = Formal charge =
group numbergroup numberin periodic tablein periodic table
number ofnumber ofbondsbonds
number ofnumber ofunshared electronsunshared electrons
–– ––
An arithmetic formula for calculating formal charge.An arithmetic formula for calculating formal charge.
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"Electron counts""Electron counts" and formal and formal charges in NHcharges in NH44
+ + and BFand BF44--
11
44
NN
HH
HH HH
HH
++77
44
....
BBFF
FF
FF
FF
....
............:: ::
:: ::
:: ::
....
––
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Lewis structures in which many (or all) covalent bonds and electron pairs are omitted.
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HH
OO
CC CC CC
HH HH HH
HH
HHHH :: ::
HH
can be condensed to:can be condensed to:
CHCH33CHCHCHCH33
OHOH
(CH(CH33))22CHOHCHOHoror
Condensed structural formulas
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Bond-line formulas
Omit atom symbols. Represent structure by showing bonds between carbons and atoms other than hydrogen.
Atoms other than carbon and hydrogen are called heteroatoms.
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CHCH33CHCH22CHCH22CHCH3 3 is shown asis shown as
CHCH33CHCH22CHCH22CHCH22OHOH is shown asis shown as
OHOH
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Bond-line formulas
Omit atom symbols. Represent structure by showing bonds between carbons and atoms other than hydrogen.
Atoms other than carbon and hydrogen are called heteroatoms.
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HH ClClCC
CC
HH22CC
HH22CC
CHCH22
CHCH22
HHHH
is shown asis shown as
ClCl
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Constitutional Isomers
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Constitutional isomers
Isomers are different compounds that have the same molecular formula.
Constitutional isomers are isomers that differ in the order in which the atoms are connected.
An older term for constitutional isomers is “structural isomers.”
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A Historical Note
In 1823 Friedrich Wöhler discovered that when ammonium cyanate was dissolved in hot water, it was converted to urea.
Ammonium cyanate and urea are constitutional isomers of CH4N2O.
Ammonium cyanate is “inorganic.” Urea is “organic.” Wöhler is credited with an important early contribution that helped overturn the theory of “vitalism.”
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NHNH44OCNOCNAmmonium cyanateAmmonium cyanate
HH22NCNHNCNH22
OO
UreaUrea
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NitromethaneNitromethane Methyl nitriteMethyl nitrite
.... ::
HH CC
OO
OO
NN
::
::....
––
++
HH
HH
Examples of constitutional isomers
Both have the molecular formula CH3NO2 but the atoms are connected in a different order.
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....CC OO NN OOHH
HH
HH
....::.... ....
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Resonance
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Resonance
two or more acceptable octet Lewis structures
may be written for certain compounds (or ions)
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How to Write Lewis StructuresHow to Write Lewis Structures
If an atom lacks an octet, use electron pairs on an adjacent atom to form a double or triple bond.
Example:Nitrogen has only 6 electrons in the structure shown.
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....CC OO NN OOHH
HH
HH
........ ::
.... ....
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How to Write Lewis StructuresHow to Write Lewis Structures Step
If an atom lacks an octet, use electron pairs on an adjacent atom to form a double or triple bond.
Example:All the atoms have octets in this Lewis structure.
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........
CC OO NN OOHH
HH
HH
....::....
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How to Write Lewis StructuresHow to Write Lewis Structures Step
Calculate formal charges.
Example:None of the atoms possess a formal charge in this Lewis structure.
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........
CC OO NN OOHH
HH
HH
....::....
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How to Write Lewis Structures How to Write Lewis Structures Step
Calculate formal charges.
Example:This structure has formal charges; is less stable Lewis structure.
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........
CC OO NN OOHH
HH
HH
.... ::....++ ––
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Resonance Structures of Methyl Nitrite
same atomic positions
differ in electron positions
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more stable more stable Lewis Lewis
structurestructure
less stable less stable Lewis Lewis
structurestructure
........
CC OO NN OOHH
HH
HH
.... ::....++ ––
........
CC OO NN OOHH
HH
HH
....::....
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Resonance Structures of Methyl Nitrite
same atomic positions
differ in electron positions
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more stable more stable Lewis Lewis
structurestructure
less stable less stable Lewis Lewis
structurestructure
........
CC OO NN OOHH
HH
HH
.... ::....++ ––
........
CC OO NN OOHH
HH
HH
....::....
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Why Write Resonance Structures?
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Electrons in molecules are often delocalized
between two or more atoms.
Electrons in a single Lewis structure are assigned to specific atoms
A single Lewis structure is insufficient to show electron
delocalization.
Composite of resonance forms more accurately
represents electron distribution.
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ExampleOzone (O3)
Lewis structure of ozone shows one double bond and one single bond
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Expect: one short bond and one Expect: one short bond and one long bondlong bond
Reality: bonds are of equal length Reality: bonds are of equal length (128 pm)(128 pm)
OO OO••••
OO••••
••••••••••••••••––++
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ExampleOzone (O3)
Lewis structure of ozone shows one double bond and one single bond
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Resonance:Resonance:
OO OO••••
OO••••
••••••••••••••••––++
OO OO••••
OO••••
••••••••••••••••––++
OO OOOO••••
••••••••••••••••
–– ++
••••
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3.7The Shapes of Some Simple Molecules
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Methane
tetrahedral geometryH—C—H angle = 109.5°
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Methane
tetrahedral geometryeach H—C—H angle =
109.5°
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Valence Shell Electron Pair Repulsions
The most stable arrangement of groups attached to a central atom is the one that has the maximum separation of electron pairs(bonded or nonbonded).
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Water
bent geometryH—O—H angle = 105°
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but notice the tetrahedral arrangement but notice the tetrahedral arrangement of electron pairsof electron pairs
OO
HH
....
HH
::
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Ammonia
trigonal pyramidal geometryH—N—H angle = 107°
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but notice the tetrahedral arrangement but notice the tetrahedral arrangement of electron pairsof electron pairs
NNHH
HH
HH
::
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Boron Trifluoride
F—B—F angle = 120°trigonal planar geometry
allows for maximum separation
of three electron pairs
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Multiple Bonds
Four-electron double bonds and six-electron triple bonds are considered to be similar to a two-electron single bond in terms of their spatialrequirements.
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Formaldehyde: CH2=O
H—C—H and H—C—Oangles are close to 120°
trigonal planar geometry
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CC OO
HH
HH
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Figure 1.12: Carbon Dioxide
O—C—O angle = 180°linear geometry
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OO CC OO
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3.7:Polar Covalent Bonds and Electronegativity
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ElectronegativityElectronegativityElectronegativityElectronegativity
Electronegativity is a measure of an element to attract electrons toward itself when bonded to another element.
Electronegativity is a measure of an element to attract electrons toward itself when bonded to another element.An electronegative element attracts electrons.An electropositive element releases electrons.
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Pauling Electronegativity ScalePauling Electronegativity ScalePauling Electronegativity ScalePauling Electronegativity Scale
1.0
Na
0.9
Li Be B C N O F
1.5
Mg
1.2
2.0
Al
1.5
2.5
Si
1.8
3.0
P
2.1
3.5
S
2.5
4.0
Cl
3.0
Electronegativity increases from left to rightin the periodic table.
Electronegativity decreases going down a group.
Electronegativity increases from left to rightin the periodic table.
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GeneralizationGeneralization
The greater the difference in electronegativitybetween two bonded atoms; the more polar the
bond.
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nonpolar bonds connect atoms ofnonpolar bonds connect atoms ofthe same electronegativitythe same electronegativity
H—HH—H ::NN NN::FF::........FF::
....
....
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GeneralizationGeneralization
The greater the difference in electronegativitybetween two bonded atoms; the more polar the
bond.
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polar bonds connect atoms ofpolar bonds connect atoms ofdifferent electronegativitydifferent electronegativity
::OO CC
FF::........HH
OO........HH
HH
OO::.... ....
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3.7Molecular Dipole Moments
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++——
not polarnot polar
Dipole Moment
A substance possesses a dipole moment if its centers of positive and negative charge
do not coincide. = e x d
(expressed in Debye units)
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——++
polarpolar
Dipole Moment
A substance possesses a dipole moment if its centers of positive and negative charge
do not coincide. = e x d
(expressed in Debye units)
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Molecular Dipole Moments
molecule must have polar bonds
necessary, but not sufficient
need to know molecular shape
because individual bond dipoles can cancel
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OO CC OO++-- --
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OO CC OO
Carbon dioxide has no dipole moment; Carbon dioxide has no dipole moment; = 0 D = 0 D
Molecular Dipole Moments
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= 1.62 D= 1.62 D = 0 D= 0 D
Carbon tetrachlorideCarbon tetrachloride DichloromethaneDichloromethane
Comparison of Dipole Moments
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Resultant of theseResultant of thesetwo bond dipoles istwo bond dipoles is
= 0 D= 0 D
Carbon tetrachloride has no dipoleCarbon tetrachloride has no dipolemoment because all of the individualmoment because all of the individualbond dipoles cancel.bond dipoles cancel.
Resultant of theseResultant of thesetwo bond dipoles istwo bond dipoles is
Carbon tetrachloride
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Resultant of theseResultant of thesetwo bond dipoles istwo bond dipoles is
= 1.62 D= 1.62 D
Resultant of theseResultant of thesetwo bond dipoles istwo bond dipoles is
The individual bond dipoles do notThe individual bond dipoles do notcancel in dichloromethane; it hascancel in dichloromethane; it hasa dipole moment.a dipole moment.
Dichloromethane
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