LU 1 chapter1.pdf C

John E. McMurry
www.cengage.com/chemistry/mcmurry
Paul D. Adams • University of Arkansas
Chapter 1
Structure and Bonding

 Organic chemistry is study of carbon compounds.
 Why is it so special?
 90% of more than 30 million chemical compounds contain carbon.
 Examination of carbon in periodic chart answers some of these questions.
 Carbon is group 4A element, it can share 4 valence electrons and form 4
covalent bonds.
Origins of Organic Chemistry

 Review ideas from general chemistry: atoms, bonds,
molecular geometry
Why This Chapter?

 Structure of an atom: small diameter (2 × 10-10 m = 200 pm)
 Nucleus very dense
 protons (positively charged)
 neutrons (neutral)
 small (10-15 m)
 Electrons
 negatively charged
 located in space remindful of a cloud (10-10 m) around nucleus
1.1 Atomic Structure
[ångström (Å) is 10-10 m = 100 pm]

 The atomic number (Z): number of protons in nucleus
 The mass number (A): number of protons plus neutrons
 All atoms of same element have the same Z value
 Isotopes: atoms of the same element with different
numbers of neutrons and thus different A
 The atomic mass (atomic weight) of an element is
weighted average mass in atomic mass units (amu) of
an element’s naturally occurring isotopes.
 Carbon:
Atomic Number
and Atomic Mass
12C
6
AC
Z
13C
6
(98.9 12.000) (1.1 13.000)
12.011
100
× + ×
=

 Four different kinds of orbitals for electrons based on
those derived for a hydrogen atom
 Denoted s, p, d, and f
 s and p orbitals most important in organic and biological
chemistry
 s orbitals: spherical, nucleus at center
 p orbitals: dumbbell-shaped, nucleus at middle
 d orbitals: elongated dumbbell-shaped, nucleus at
center
Shapes of Atomic Orbitals for
Electrons

 Orbitals are grouped in shells of increasing size and
energy
 Different shells contain different numbers and kinds of
orbitals
 Each orbital can be occupied by two electrons
Orbitals and Shells
(Continued)

 First shell contains one s orbital, denoted 1s, holds only
two electrons
 Second shell contains one s orbital (2s) and three p
orbitals (2p), eight electrons
 Third shell contains an s orbital (3s), three p orbitals
(3p), and five d orbitals (3d), 18 electrons
Orbitals and Shells
(Continued)

 In each shell there are three perpendicular p orbitals,
px, py, and pz, of equal energy
 Lobes of a p orbital are separated by region of zero
electron density, a node
P-Orbitals

 Ground-state electron configuration (i.e., lowest
energy arrangement) of atom
 lists orbitals occupied by its electrons.
 Rules:
 1. Lowest-energy orbitals fill first: 1s → 2s → 2p → 3s →
3p → 4s → 3d (Aufbau (“build-up”) principle)
1.3 Atomic Structure: Electron
Configurations

 Rules:
 2. Electrons act as if they were spinning around an axis.
Electron spin can have only two orientations, up ↑ and
down ↓. Only two electrons can occupy an orbital, and
they must be of opposite spin (Pauli exclusion
principle) to have unique wave equations
Configurations

 Rules:
 3. If two or more empty orbitals of equal energy are
available, electrons occupy each with spins parallel until
all orbitals have one electron (Hund's rule).
 To the chalkboard (p-orbital filling as example)
Configurations

 Atoms form bonds because the resulting compound is
more stable than the separate atoms
 Ionic bonds in salts form by electron transfers
 Organic compounds have covalent bonds from sharing
electrons (G. N. Lewis, 1916)
1.4 Development of Chemical
Bonding Theory

 Lewis structures (electron dot) show valence electrons of an atom as dots
 Hydrogen has one dot, representing its 1s electron
 Carbon has four dots (2s2 2p2) due to 4 e- in valence shell
 Kekulé structures (line-bond structures) have a line drawn between two atoms
indicating a 2 e- covalent bond.
 Stable molecule results at completed shell, octet (eight dots) for main-group
atoms (two for hydrogen)
Development of Chemical
Bonding Theory

 Atoms with one, two, or three valence electrons form
one, two, or three bonds.
Bonding Theory
valence e- valence e-

 Atoms with four or more valence electrons form as
many bonds as electrons needed to fill the s and p
levels of their valence shells to reach a stable octet.
 Carbon has four valence electrons (2s2 2p2), forming
four bonds (CH4).
Bonding Theory
valence e-

 Nitrogen has five valence electrons (2s2 2p3) but forms
only three bonds (NH3).
Bonding Theory
valence e-

 Oxygen has six valence electrons (2s2 2p4) but forms
two bonds (H2O)
Bonding Theory
valence e-

Bonding Theory

 Valence electrons not used in bonding are called nonbonding
electrons, or lone-pair electrons
 Nitrogen atom in ammonia (NH3)
 Shares six valence electrons in three covalent bonds
and remaining two valence electrons are nonbonding
lone pair
Non-Bonding Electrons

 Covalent bond forms when two atoms approach each other
closely so that a singly occupied orbital on one atom overlaps a
singly occupied orbital on the other atom
 Two models to describe covalent bonding.
 Valence bond theory
 Molecular orbital theory
1.5 Describing Chemical Bonds:
Valence Bond Theory

Valence Bond Theory:
 Electrons are paired in the overlapping orbitals and are attracted to
nuclei of both atoms
 H–H bond results from the overlap of two singly occupied
hydrogen 1s orbitals
 H-H bond is cylindrically symmetrical, sigma (σ) bond
Valence Bond Theory
cylindrically symmetrical

 Reaction 2 H· → H2 releases 436 kJ/mol
 i.e., product has 436 kJ/mol less energy than two atoms:
H–H has bond strength of 436 kJ/mol
Bond Energy

 Distance between
nuclei that leads to
maximum stability
 If too close, they
repel because both
are positively
charged
 If too far apart,
bonding is weak
Bond Energy

 Kekulé and Couper independently observed that
carbon always has four bonds
 van't Hoff and Le Bel proposed that the four bonds of
carbon have specific spatial directions
 Atoms surround carbon as corners of a tetrahedron
Describing Chemical Bonding
Theory

 Carbon has 4 valence electrons (2s2 2p2)
 In CH4, all C–H bonds are identical (tetrahedral)
 sp3 hybrid orbitals: an s orbital and three p orbitals
combine: form four equivalent, unsymmetrical,
tetrahedral orbitals (s + ppp = sp3)
1.6 sp3 Orbitals and the
Structure of Methane
Linus Pauling (1931): his picture near men’s bathroom across from elevators

 sp3 orbitals on C overlap with 1s orbitals on 4 H atoms
to form four identical C-H bonds
 Each C–H bond has a strength of 439 kJ/mol and
length of 109 pm
 Bond angle: each H–C–H is 109.5°: the tetrahedral
angle.
The Structure of Methane

1.7 sp3 Orbital-based Structure
of Hexane

 Some Representations of Ethylene are given
 To explain planar geometry and trigonal shape about C’s in
ethylene
Structure of Ethylene
C C
H
H
H
H
H2C CH2 C C
H
H
H
H

 sp2 hybrid orbitals: 2s orbital combines with two 2p
orbitals, giving 3 orbitals (s + pp = sp2). This results in a
double bond.
 sp2 orbitals are in a plane with120° angles
 Remaining p orbital is perpendicular to the plane

 Two sp2-hybridized orbitals overlap to form a σ bond
 p orbitals overlap side-to-side to formation a pi (π)
bond
 sp2–sp2 σ bond and 2p–2p π bond result in sharing four
electrons and formation of C-C double bond
Bonds From sp2 Hybrid
Orbitals

 Electrons in the σ bond are centered between nuclei
 Electrons in the π bond occupy regions are on either
side of a line between nuclei
Bonds From sp2 Hybrid
Orbitals

 H atoms form σ bonds with four sp2 orbitals
 H–C–H and H–C–C bond angles of about 120°
 C–C double bond in ethylene shorter and stronger than
single bond in ethane
 Ethylene C=C bond length 134 pm (C–C 154 pm)

 Sharing of six electrons forms C ≡C
 Two sp orbitals form σ bonds with hydrogens
1.9 sp Orbitals and the
Structure of Acetylene

 C-C a triple bond sharing six electrons
 Carbon 2s orbital hybridizes with a single p orbital
giving two sp hybrids
 two p orbitals remain unchanged
 sp orbitals are linear, 180° apart on x-axis
 Two p orbitals are perpendicular on the y-axis and the
z-axis
sp Orbitals and the Structure of
Acetylene

 Two sp hybrid orbitals from each C form sp–sp σ bond
 pz orbitals from each C form a pz–pz π bond by
sideways overlap and py orbitals overlap similarly
Orbitals of Acetylene

 Elements other than C can have hybridized orbitals
 H–N–H bond angle in ammonia (NH3) 107.3°
 C-N-H bond angle is 110.3 °
 N’s orbitals (sppp) hybridize to form four sp3 orbitals
 One sp3 orbital is occupied by two nonbonding
electrons, and three sp3 orbitals have one electron
each, forming bonds to H and CH3.
1.10 Hybridization of Nitrogen
and Oxygen

 A molecular orbital (MO): where electrons are most likely
to be found (specific energy and general shape) in a
molecule
 Additive combination (bonding) MO is lower in energy
 Subtractive combination (antibonding) MO is higher energy
Molecular Orbital Theory

 The π bonding MO is from combining p orbital lobes
with the same algebraic sign
 The π antibonding MO is from combining lobes with
opposite signs
 Only bonding MO is occupied
Molecular Orbitals in
Ethylene

 Drawing every bond in organic molecule can
become tedious.
 Several shorthand methods have been
developed to write structures.
 Condensed structures don’t have C-H or C-C
single bonds shown. They are understood.
e.g.
1.12 Drawing Structures
(Avoided in this Class)

General Rules:
1) Carbon atoms aren’t usually shown
Drawing Skeletal Structures
(Commonly Used)

General Rules:
2) Instead a carbon atom is assumed to be at each
intersection of two lines (bonds) and at the end of
each line.
(Commonly Used)

General Rules:
3) Hydrogen atoms bonded to carbon aren’t shown.
(Commonly Used)

General Rules:
4) Atoms other than carbon and hydrogen ARE shown
(Commonly Used)

 Organic chemistry – chemistry of carbon compounds
 Atom: charged nucleus containing positively charged protons and
netrually charged neutrons surrounded by negatively charged
electrons
 Electronic structure of an atom described by wave equation
 Electrons occupy orbitals around the nucleus.
 Different orbitals have different energy levels and
different shapes
 s orbitals are spherical, p orbitals are dumbbell-shaped
 Covalent bonds - electron pair is shared between atoms
 Valence bond theory - electron sharing occurs by overlap of two
atomic orbitals
 Molecular orbital (MO) theory - bonds result from combination of
atomic orbitals to give molecular orbitals, which belong to the
entire molecule
Summary

 Sigma (σ) bonds - Circular cross-section and are formed by head-
on interaction
 Pi (π) bonds - “dumbbell” shape from sideways interaction of p
orbitals
 Carbon uses hybrid orbitals to form bonds in organic molecules.
 In single bonds with tetrahedral geometry, carbon has four sp3
hybrid orbitals
 In double bonds with planar geometry, carbon uses three
equivalent sp2 hybrid orbitals and one unhybridized p orbital
 Carbon uses two equivalent sp hybrid orbitals to form a triple
bond with linear geometry, with two unhybridized p orbitals
 Atoms such as nitrogen and oxygen hybridize to form strong,
oriented bonds
 The nitrogen atom in ammonia and the oxygen atom in water
are sp3-hybridized
Summary (Continued)

Draw an electron-dot structure for acetonitrile, C2H3N, which
contains a carbon-nitrogen triple bond. How many electrons
does the nitrogen atom have in its outer shell ? How many
are bonding, and how many are non-bonding?
Let’s Work a Problem

To address this question, we must realize that the nitrogen
will contain 8 electrons in its outer shell. Six will be used in
the C-N triple bond (shaded box), and two are non-bonding
Answer

LU 1 chapter1.pdf C

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