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Chapter 1
Carbon Compounds and Chemical
Bonds
Organic Chemistry
– The chemistry of the compounds of carbon
History-Unofficially, Organic is one of the oldest sciences
-Officially, it is one of the youngest
History
• Vitalism- the belief that the intervention of a
“vital force” was necessary to the synthesis of
organic molecules.
• Friedrich Wöhler, 1828, “Father of Organic
Chemistry”
O
H
H
N
H
O
N
C
C
H
Ammonium Cyanate
heat
H 2N
NH 2
Urea
On your own
• Review section 1.2 to make sure you
understand terms such as:
– Compounds, elements, atoms, isotopes, electron
shells, valence shell, valence electrons, etc.
Structural Theory
• 1800’s by Kekulé, Couper, and Butlera
• Two central premises
– 1) The atoms in organic compounds can form a
fixed number of bonds
• Valence- the measure of an atoms ability to form
bonds
– 2) A carbon atom can use one or more of its
valences to form bonds to other carbon atoms
Isomers
• Isomers- Different compounds with the same
molecular formula.
ex.
• Constitutional Isomers: Isomers that differ in
their connectivity, that is, in the sequence in
which their atoms are bonded together
Two types of Bonding
• Ionic Bonds- Formed by the transfer of one or
more electrons from one atom to another to
create ions
• Covalent Bonds- a bond that forms when
atoms share electrons with one another.
• Octet Rule- the tendency for atoms to achieve
an electron configuration where its valence
shell contains eight electrons.
Ionic Bonds
• Occur between atoms of widely different
electronegativity.
• Typically, that means between metals and
non-metals
Ex.
Electronegativity
• Electronegativity- a measure of the ability of
an atom to attract electrons.
• Periodic Table Trends
– Increases as you move left to right
– Increases as you move bottom to top
• Important Order of Electronegativity:
F > O > Cl,N > Br > C,S > H,P
Why Do Ions Form?
• Electronically identical to Nobel Gas
• Stabilize each other
• Note:
– Electrons dictate reactivity, not charges!!
Properties of Ionic Compounds
• Very strong
• High MP
• Sometimes dissolve in Polar Solvents
• Usually called Salts
Covalent Bonds
• Occur when two atoms have the same or
similar electronegativity
• “Share” electrons instead of complete transfer
• Typically called molecules
• Molecules can be represented by electron dot
formulas, or dash formulas, where a dash
represents a pair of shared electrons.
• Examples:
Lewis Structures
• Lewis Structures are electron dot formulas
where only the valence electrons are shown
• Multiple bonds are represented by multiple
lines.
• Ex.
• Note: Ions may also contain covalent bonds!
• Ex.
Writing Lewis Structures
• Assemble the molecule or ion showing only
valence electrons
• Strive to give each atom an octet, except
Hydrogen which receives only 2
• The number of valence electrons for an atom
is equal to the group number
• If the structure is an ion, we add or subtract
electrons to give appropriate charge
Rules:
1) Find total number of valence electrons
2) Use pairs of electrons to form bonds
between the atoms. (Note: remember to
consider the typical valence of each atom)
3) Add remaining electrons as pairs to give each
atom an octet.
Exceptions to Octet Rule
• 1st period elements- only two electrons
• 2nd period elements- Boron is stable with six
valence electrons
• 3rd period and beyond- Elements have access
to d orbitals which allows them to
accommodate more than 8 electrons
Formal Charges
• Formal Charge- the charge associated with the
electronic difference between the atomic
state and bonded state of an atom
• The sum of the formal charges on each
element equals the total charge for the
molecule or ion.
• To calculate formal charges, simply compare
the number of electrons an atom has in the
bonded state to the number it has in the
atomic state.
Counting Valence Electrons in the
Bonded State
• Non-shared electrons, or lone pairs, belong
solely to the atom that posses them
• Shared pairs of electrons are split giving half
the shared electrons to each atom sharing
them.
• Ex.
Common Formal Charge States
Resonance Structures
• Lewis structures incorrectly create an artificial
location for electrons
• Consider the carbonate ion, CO32-
Resonance Structures, cont
• We can account for the experimental data by
showing how the structures can be converted to
the others using curved arrows.
• (arrow speech!)
• The overall structure, or true structure, is a
mixture of the individual structures, and is called
the resonance hybrid.
• Note: Individual resonance structures only exist
on paper!!
Rules for Drawing Resonance
Structures
1) Resonance Structures only exist on paper!
2) You are only allowed to move electrons.
3) All structures must be proper Lewis structures
4) The energy of the actual molecule will always be
lower than a single contributing structure. This
is called Resonance Stabilization.
Rules for Drawing Resonance
Structures, cont
5) Equivalent resonance structures make equal
contributions to the overall structure.
6) The more stable a structure, the more it will
contribute to the overall structure:
a) The more covalent bonds a structure
has, the more stable it is.
Rules for Drawing Resonance
Structures, cont
b) Structures in which all atoms have octets are
more stable
c) Opposite charge separation decreases
stability
d) Structures with negative charges on highly
electronegative atoms are more stable than those
with negative charges on less electronegative atoms
Section 1.9 Quantum Mechanics and
Atomic Structure
•
•
•
•
WHAT YOU NEED TO KNOW:
Electrons in atoms and molecules have both particle
and wave characteristics
Wave properties are used to predict the shape of
orbitals
Orbitals- a region of space where the probability of
finding an electron is very large
The volumes we use to express the orbital shapes
represents where the electron would be 90-95% of
the time.
Atomic Orbital shape
Energy of orbitals
1s < 2s < 2px = 2py = 2pz < 3s
• Degenerate orbitals- atomic orbitals of equal
energy, like the three 2p orbitals
• Using these relative energies, we can derive
the electron configuration
Electron Configuration
1) Aufbau Principle: Orbitals are filled so that those of
lowest energy are filled first.
2) Pauli Exclusion Principle: Only two electrons are
allowed in each orbital and must have opposite spins
3) Hund’s Rule: When multiple orbitals of equal
energy are present, each orbital receives one electron
before any pairs are created.
Molecular Orbitals
• Atomic Orbitals combine to form molecular
orbitals which are used for bonding.
• The number of Molecular orbitals must equal
the number of Atomic orbital used.
• In General Chemistry, we concentrated on the
new Molecular orbitals formed from bonding
• In organic, we use this theory to explain the
combination of atomic orbitals on a single
atom, called Hybridization, to form new
orbitals that are then used for bonding.
• Consider the simplest organic molecule,
methane, CH4
Methane
• To account for what we see, we must mix the
2s orbital with all three 2p orbitals.
• This forms four new, equal orbitals called sp3
hybridized orbitals!
Sigma Bonds
• When hybridized orbitals overlap, a sigma
bond is formed
• Sigma Bond- term used to describe bonds in
which the greatest density of electrons lies
between the two nuclei.
• Sigma bonds have cylindrical symmetry along
the bond axis.
• As a result, there is “free” rotation about
sigma bonds.
sp2 Hybridization
• When two carbons share two pairs of electrons, the
result is a carbon-carbon double bond.
• Hydrocarbons whose molecules contain a carboncarbon double bonds are called alkenes.
• Pi bond- created by the overlap of p orbitals above
and below the sigma bond framework.
Restricted Rotation
• For maximum overlap, the p orbitals must be
parallel
• As a result, ~264 kJ/mol of energy is needed to
break the p orbital overlap and allow rotation
to occur
• Only about 13-26 kJ/mol is needed to rotate
around a sigma bond.
Cis/Trans Isomers
• Because of the restricted rotation, a new form
of isomers is created call Cis/Trans Isomers.
• Ex.
• These compounds are not superposable.
• They are not constitutional isomers because
the connectivity is the same
Stereoisomers
• Stereoisomers- isomers that differ only in the
arrangement of their atoms in space.
• For a double bond to qualify for Cis/Trans
Isomers, both carbons of the double bond
must be bonded to two different groups.
• Ex.
sp hybridization
• When two carbons share three pairs of
electrons, a triple bond is formed.
• Hydrocarbons in which two carbons share
three pairs of electrons are called alkynes.
• The triple bond consists of 2 pi bonds and 1
sigma bond.
Bond Lengths
• The shortest C-H bonds are achieved with
hybridized orbitals having the most s
character.
%s
sp <
50%
sp2
33%
<
sp3
25%
VSEPR Theory
• Valence Shell Electron Pair Repulsion Theory is
used to predict the geometry of atoms and
the resulting shape of molecules.
• It consists of four points:
1) We consider molecules (or ions) in which the
central atom is covalently bonded to two or
more groups.
VSEPR cont
2) We consider all of the valence electron pairs of the
central atom
-Shared electrons are called bonding pairs
-Unshared electrons are called non-bonding or
unshared pairs
3) Because electrons repel each other, they stay as far
away from one another as possible
Nonbonding repulsion > bonding
VSEPR cont
4) We arrive at the geometry of an atom by
considering both bonding and nonbonding
electrons but the shape of the molecule by
referring to the position of nuclei.
See table 1.3 on page 47.
Structural Formulas
• Dash Formula- must show all atoms, bonds,
and lone pairs.
• Condensed Formula- must show all atoms but
may or may not show bonds
• Bond Line Formula- show bonds and all atoms
except Carbon and Hydrogen
3-D designations
• Wedge bond- used to show bond coming out
of plane towards the viewer
• Dash bond- used to show bond going behind
the plane away from the viewer.
• These can be used with any type of structural
formula.
Applications of Basic Principles
• Page 47, section 1.17
• These sections are available at the end of
most chapters. You should review these as
they associate basic theories to topics in
organic chemistry.
Other Tools
• There are also other great tools at the end of
the chapters such as Summaries and Review
Tools, Practice Problems, Concept Maps, and
Synthetic Connections.