the basic – bonding and molecular structure. organic chemistry and life
TRANSCRIPT
The Basic – Bonding and Molecular Structure
Organic Chemistry and Life
1.1 The Development of Organic Chemistry as ScienceOrganic compounds: compounds that
could be obtained from living organismsThe scientific study of the structure,
properties, composition, reactions, and preparation (by synthesis or by other means) of chemical compounds that contain carbon
Inorganic compounds: those came from non-living sourcesOccur as a salts
Atomic Orbitals
Atomic OrbitalsS - orbital p- orbital
Electrons ConfigurationShow
Orbital Diagram and Electron ConfigurationElectrons
configuration: H, He, C, Mg
Aufbau Principle: fill lowest orbital first to full capacity, then next
1.2 The structural Theory of Organic ChemistryAtoms in organic compounds can form a fixed
number of bonds using their valence electrons
C
Carbon atoms are tetravalent
O
Oxygen atomsare divalent
H Cl
Hydrogen and halogenare monovalent
1.2 The structural Theory of Organic ChemistryA carbon atom can use one or more of its
valence electrons to form bonds to other carbon atoms
C C C C C C
single bond double bond triple bond
1.3 Isomers: The Importance of Structural FormulasConstitutional isomers – non identical
compounds with same molecular formulaDo not necessary share similar properties
H C C O H
H
H
H
H
ethyl alcohol
H C O C
H
H
H
H
H
Dimethyl ether
1.4 Ionic BondsOccurs in ionic compound
Results from transferring electron
Created a strong attraction among the closely pack compound
Covalent Bonding Formation of a covalent Bond
Two atoms come close together, and electrostatic interactions begin to develop
Two nuclei repel each other; electrons repel each other Each nucleus attracts to electrons; electrons attract
both nuclei Attractive forces > repulsive forces; then
covalent bond is formed
Electronegativity Electronegativity
(EN): the ability of an atom in a molecule to attract the shared electron in a bond
Metallic elements – low electronegativities
Halogens and other elements in upper right-hand corner of periodic table – high electronegativity
PolarityPolar covalent bonds
– the bonding electrons are attracted somewhat more strongly by one atom in a bondElectrons are not
completely transferredMore electronegative
atom: δ- . (δ represents the partial negative charge formed)
Less electronegative atom: δ+
Lewis Structuresrepresents how an atom’s valence electrons
are distributed in a molecule Show the bonding involves (the maximum
bonds can be made)Try to achieve the noble gas configuration
RulesDuet Rule: sharing of 2 electrons
E.g H2 H : H
Octet Rule: sharing of 8 electronsCarbon, oxygen, nitrogen and fluorine always
obey this rule in a stable moleculeE.g F2, O2
Bonding pair: two of which are shared with other atoms
Lone pair or nonbonding pair: those that are not used for bonding
1 Lewis Structures of Molecules with Multiple BondsUse 6N + 2 Rule
N = number of atoms other than HydrogenIf
Total valence – (6N + 2) = 2 1 double bond
Total valance e- - (6N + 2) = 4 two double bonds or 1 triple bond
ExamplesWrite the Lewis structure of CH3F, ClO3
-, F2
1.7 Formal ChargesDifference between the number of outer-shell
electrons “owned” by a neutral free atom and the same atom in a compound
Formal charge = group # - unshared e- - share e-12
ExamplesDetermine the formal charge for each atom
in the following moleculesNH4
+
NO2-
CO32-
ResonanceWhenever a molecule or ion can be
represented by two or more Lewis structures that differ only in the position of the electronsNone of these resonance structures will be a
correct representation for the molecule or ionThe actual molecule or ion will be better
represented by a hybrid or hypothetical structures
Represented by a double headed arrows ( )
O
C
O O
23
-
23
-23
-
hybrid
O
C
O O
O
C
O O
O
C
O O
contributing resonance structures
Resonance - stabilizationThe more covalent bonds a structure has, the
more stable it is
H2C CH
CH
CH2 H2C CH
CH
CH2
H2C CH
CH
CH2
Resonance-stabilizationStructure in which all the atoms have a
complete valence shell of electrons are especially H2C O CH3 H2C O CH3
Here this carbonatom has only sixelectrons
Here the carbon atomhas eight electrons
Resonance stabilizationCharge separation decrease stabilizationResonance contributors with negative charge
on highly electronegative atoms are stable ones with negative charge on less or nonelectronegative atoms
H2C CH
Cl H2C CH
Cl
1.9 Quantum Mechanisms and Atomic Structure Schröndinger’s quantum mechanical model of
atomic structure is frame in the form of a wave equation; describe the motion of ordinary waves in fluids.
i. Wave functions or orbitals (Greek, psi , the mathematical tool that quantum mechanic uses to describe any physical system
ii. 2 gives the probability of finding an electron within a given region in space
iii. Contains information about an electron’s position in 3-D space
defines a volume of space around the nucleus where there is a high probability of finding an electron
say nothing about the electron’s path or movement
11.2 Electromagnetic Radiation Radiation energy – has
wavelike properties Frequency (υ, Greek nu) –
the number of peaks (maxima) that pass by a fixed point per unit time (s-1 or Hz)
Wavelength (λ, Greek lambda) – the length from one wave maximum to the next
Amplitude – the height measured from the middle point between peak and trough (maximum and minimum)
Intensity of radiant energy is proportional to amplitude
Constructive interfence of wave Destructive interference of wave
1.10 Atomic orbitalHeisenberg Uncertainty
Principle – both the position (Δx) and the momentum (Δmv) of an electron cannot be known beyond a certain level of precision
1. (Δx) (Δmv) > h 4π
2. Cannot know both the position and the momentum of an electron with a high degree of certainty
Molecular OrbitalsTwo types of atomic of atomic orbitals are
combined as they come close to each otherHybridization: blending combination of atomic orbitals
to form new orbitalCarbon has three possible molecular orbitals
sp3 sp2 sp
Orbitals repsonsible for creating the covalent bonds2 special names for covalent bonds of organic
moleculesSigma (σ) bond Pi (π) bond
Created when “head on” overlap occurs of orbitals
Created when “side on” overlap occurs of orbitals
sp3 orbitals responsible for creating all “single bonds” of all organic molecules alkanes
s sp3 sp3sp3
take pure atomic orbitalsand mix them
gives 4 molecular orbitals called sp3 orbitalsall equal in size, energy and shape
p sp3
Examples
H C
H
H
H
C
H
H
H
H
sp3C - sH
C C
H
H
H
H
H
H
sp2 molecular orbitals
All sp2 molecular orbitals responsible for creating all double bonds in organic molecules
alkenes
s sp2sp2 unhybridized orbital
p
hybridization
take pure s + 2p"mix them"
gives 3 molecular orbitals called sp2and lef t over p atomic orbital (unhybridized)sp2 orbitals equal in size, energy and shape
sp2p
C C
H
H
H
H
C C
H
H H
H
sp2C - sp2
C
sp2C - sH
C C
H
H
H
H
sp2C - sp2
C
framework
framework
1.13B – Cis –Trans IsomerismWhich of the following alkene can exist as cis-
trans isomers? Write their structure
H2C CHCH2CH3a.
b.H2C C(CH3)2
c. CH3CH CHCH3d. CH3CH2CH CHCl
sp molecular orbitals
All sp orbitals responsible for creating all triple bonds of organic molecules
alkynes
s sp sp unhybridized orbitalp
hybridization
p
take 1s + 1pmix them
give 2 molecular orbitals called sp and 2 unhybridized p atomic orbitals
H C C H
C CH H
spC - spC
spC - sH
H C C H
spC - spC
ExamplesDraw a bonding picture for the following
molecule, showing all π, σ – bonds using σ-framework and π-framework
C CH
H
H
C C H
Molecular OrbitalsTwo types:
Bonding molecular orbitals Contains both
electrons in the lowest energry state or ground state
Formed by intereaction of orbitals with same phase signs
Increases the propability
Molecular orbitalsAntimolecular
orbitals Contains no
electrons in the ground state
Formed by intereaction of orbitals with opposite phase signs
Result with nodes
Molecular orbitals
Shape of MoleculesVSEPR Theory
Valence shell electron pair repulsionBond angles and geometry
Steric number = # bond to - # lone pairs central atom to central atom
Rules: 1- Carbon will always be the central atom 2 – Double bond; triple bonds will count
as 1 bond
H C
H
H
H
109.5o
tetrahedral
C CC
HH
OH
H
H
H
120o 120o
trigonal planar
C CH H
180o
linear
O
H H
V-shaped or bent
109.5o
Molecular shapesVSEPR method can be used to predict the
shapes of molecules containing multiple bondsAssume that all electrons of a multiple bond
act as one unit
ExamplesUse VSEPR theory to predict the geometry of
each of the following molecules and ionsSiF4
BeF2
1.17 Representation of Structural formulasStructual formula for propyl alcohol
O H
H
H
H
H
H
H
H
Dash formulaor structural formula
CH3CH2CH2OH
condensed formula
OH
Bond-line formulaor stick f igure
Dash StructureAtoms are joined by single bonds can rotate
relatively freely with respect to one another
C
C
C
H
O
H
HH H
H H
H
C
C
C
O
H
H
H
H H
H H
H
C
C
C
OH
HH
H H
H H
H
12
3
Equivalent dash formula for propyl alcohol
C C C O
H
H
H H
H
H
H
C C C H
H
H
H H H
H
H
O
H
constitutional isomers
Condensed Structural FormulasAll hydrogen atoms are written immediately
after the carbon that they’re attached
H C C C H
H
H
H
Cl
H
H
H C C C C
H
H
H
H
H
H
H
H
O H
CH3CH2CH3
Bond-Line FormulasHydrogen and carbon atoms will not appear
in the formulaEach end of the line represents carbon atom
C C C O
H
H
H H
H
H
H
H
OH
H C C C H
H
H
H
Cl
H
H
Cl
bond-line
structural
ExamplesFor each of the following, write a bond line
formula(CH3)2NCH2CH3
H2C
H2C CH2
a. b.
c.
H3C
HC
CH
H2C
Cl
d.
CH3CHC CH
F