Download Structure and Bonding

1/19/2011
John E. McMurry
http://www.cengage.com/chemistry/mcmurry
Chapter 1
Structure and Bonding
Richard Morrison • University of Georgia, Athens
Scientific Revolution
Scientific revolution leads to:
• Safer and more effective medicines
• Cures for genetic diseases
• Increased life span
• Improved quality of life
Scientific advances in medicine and biology require an
understanding of organic chemistry
Organic Chemistry
Organic and Biological Chemistry in Modern Medicine
• Coronary artery disease (CAD) is directly correlated with blood
cholesterol levels
• 75% of blood cholesterol (1000 mg each day) is biosynthesized
• Drugs known as Statins reduce the risk of CAD by lowering blood
cholesterol levels
• 3-Hydroxy-3-methylglutaryl coenzyme A (HMG-CoA) to mevalonate
conversion
• Crucial step in biosynthesis of cholesterol
• Atorvastatin (Lipitor) inactivates enzyme that catalyzes HMG-CoA to
vevalonate conversion
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Organic Chemistry
Carbon
Organic chemistry
• All organic compounds contain the element carbon
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4A element
Shares four electrons
Forms four strong covalent bonds
Bonds to other carbons to create chains and rings
Not all carbon compounds are derived from living
organisms
Over 99% of 37 million known compounds contain
carbon
Common Elements in Carbon Compounds
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1.1 Atomic Structure: The Nucleus
Structure of atom
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Nucleus
• Positively charged
• Made up of protons (positively charged) and neutrons (neutral)
• Small (10-14 to 10-15 m in diameter)
• Contains essentially all the mass of the atom
Electron cloud
• Negatively charged electrons in cloud around nucleus
• Atomic diameter is about 2 Angstroms (Å)
-10 m = 100 pm ; 1 pm = 10-12 m
• 1 Å = 10
• Atomic diameter in SI units is 2 10-10 m (200 picometers (pm))
Atomic Structure: The Nucleus
Atomic number (Z)
• Number of protons in the atom's nucleus
• All atoms of a given element have the same atomic
number
Mass number (A)
• Number of protons plus neutrons in the atom’s nucleus
Isotopes
• Atoms of the same element (same Z) that have different
Numbers of neutrons and therefore different mass
numbers (different A)
Atomic mass (atomic weight)
• Weighted average mass in atomic mass units (amu) of
an element’s naturally occurring isotopes
1.2 Atomic Structure: Orbitals
Quantum Mechanical Model
• Behavior of a specific electron in an atom described
by mathematical expression called a wave equation
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Wave equation is similar to mathematical expression
used to describe motion of waves in fluids
• Solution of wave equation is called a wave function
• Wave function is an orbital
• Orbital denoted by Greek letter psi,
• Plot of
2 describes volume of space around
nucleus that the electron is most likely to occupy
• Electron cloud has no sharp boundary
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Atomic Structure: Orbitals
Four different kinds of orbitals for electrons
• Denoted s, p, d, and f
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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
• four cloverleaf-shaped and one dumbbell-doughnut
Atomic Structure: Orbitals
Orbitals are occupied by zero, one, or two electrons and are
grouped in electron shells of increasing size and energy
Electron Shell
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A group of an atom’s electrons with the same principal quantum number
First shell contains one s orbital, denoted 1s, which holds only two
electrons
Second shell contains four orbitals, one s orbital (2s) and three p orbitals
(2p), which hold a total of eight electrons
Third shell contains nine orbitals, one s orbital (3s), three p orbitals (3p),
and five d orbitals (3d), which hold a total of 18 electrons
Atomic Structure: Orbitals
p Orbital
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In each shell, beginning with the second, there are three mutually
perpendicular p orbitals, px, py, and pz, of equal energy
Lobes of p orbitals are separated by region of zero electron density, a
node
Each lobe has a different algebraic sign, + and -, represented by
different colors
Algebraic signs are not charges
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1.3 Atomic Structure: Electron Configuration
Ground-state electron configuration
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Most stable, lowest-energy electron configuration of an atom
Three rules:
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Aufbau principle
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Lowest-energy orbitals fill first: 1s
2s
2p
3s
3p
4s
3d
Pauli Exclusion Principle
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Electron spin can have only two orientations, up and down
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Only two electrons can occupy an orbital, and they must be of
opposite spin to have unique wave equations
Atomic Structure: Electron Configuration
Hund's rule
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If two or more empty orbitals of equal energy are available,
electrons occupy each orbital with parallel spins until all
orbitals have one electron
1.4 Development of Chemical Bonding
Theory
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1858 - August Kekulé and Archibald Couper independently proposed
that carbon is tetravalent (always forms four bonds)
Emil Erlenmeyer proposed a carbon-carbon triple bond for acetylene
Alexander Crum Brown proposed a carbon-carbon double bond
1865 - Kekulé suggested that carbon chains can double back to form
rings of atoms
1874 - Jacobus van’t Hoff and Joseph Le Bel proposed four atoms to
which carbon is bonded sit at the corner of a regular tetrahedron
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Development of Chemical Bonding Theory
Atoms bond because the compound that results is more stable
and lower in energy than the separate atoms
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Energy is released from the chemical system when a bond
forms
Energy is consumed by the system when a bond breaks
Valence shell
• Outer most electron shell of an atom
• Eight electrons in valence shell (an electron octet) impart
special stability to noble-gas elements in 8A
• Ne (2 + 8); Ar (2 + 8 + 8); Kr (2 + 8 + 18 + 8)
• Chemistry of main group elements governed by their tendency
to take on electron configuration of the nearest noble gas
Development of Chemical Bonding Theory
Ionic compounds
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Some elements achieve an octet configuration by gaining or losing
electrons
Ions form when an electron is gained or lost from a neutral atom
Ions are charged because they have different numbers of protons and
electrons
Ions are held together by an electrostatic attraction, like in Na+ Cl-,
forming an ionic bond
Covalent compounds
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Covalent Bond
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Molecule
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Carbon achieves an octet configuration by sharing electrons
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Bond formed by sharing electrons between atoms
Neutral collection of atoms held together by covalent bonds
Development of Chemical Bonding Theory
Lewis structures (electron-dot structures)
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Dot representations of covalent bonds in molecules
Valence shell electrons of an atom are represented as dots
Kekulé structures (line-bond structures)
• Two-electron covalent bond is represented by a line
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Development of Chemical Bonding Theory
Number of covalent bonds depends on how many additional
valence electrons needed to reach noble-gas configuration
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H (1s) needs one more electron to attain (1s2) and forms one bond
N (2s22p3) needs three more electrons to attain (2s22p6) and forms
three bonds
Lone-pair electrons
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Valence-shell electron pairs not used for bonding
Worked Example 1.1
Predicting the Number of Bonds to Atoms in a
Molecule
How many hydrogen atoms does phosphorus bond to
in phosphine, PH??
1.5 The Nature of Chemical Bonds: Valence
Bond Theory
Valence bond theory
• Bonding theory that describes a
covalent bond as resulting from the
overlap of two atomic orbitals
• Electrons are paired in overlapping
orbitals and are attracted to nuclei of
both atoms, thus bonding the two
atoms together
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H–H bond results from the overlap of
two singly occupied hydrogen 1s
orbitals
H-H bond is cylindrically symmetrical
Bonds formed by head-on overlap of
two atomic orbitals along a line drawn
between the nuclei are sigma ( )
bonds
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The Nature of Chemical Bonds: Valence
Bond Theory
Bond strength
• H2 molecule has 436 kJ/mol less energy than the starting 2
H• atoms, the product is more stable than the reactant and
the H-H bond has a strength of 436 kJ/mol
• Conversely, the bond dissociation energy of H2 is 436
kJ/mol because it requires 436 kJ/mol of energy to break the
H2 bond
The Nature of Chemical Bonds: Valence
Bond Theory
There is an optimum distance between nuclei that leads to
maximum stability called the bond length
Bond length
• The distance between nuclei
at the minimum energy point
• Because a covalent bond is
dynamic, like a spring, the
characteristic bond length is
the equilibrium distance
between the nuclei of two
atoms that are bonded to
each other
1.6 sp3 Hybrid Orbitals and the Structure
of Methane
Carbon has four valence electrons (2s22p2) that form four
bonds
Methane CH4
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All four carbon-hydrogen bonds in methane are identical and are spatially
oriented toward the corners of a regular tetrahedron
sp3 hybrid orbitals
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A hybrid orbital derived from the combination of an s atomic orbital with
three p atomic orbitals
Linus Pauling (1931) showed mathematically how s orbitals and p
orbitals on an atom can combine, hybridize, to form four equivalent
atomic orbitals with tetrahedral orientation called sp3 hybrids
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sp3 Hybrid Orbitals and the Structure of
Methane
Two lobes of a p orbital have different algebraic signs (+ and -)
in corresponding wave functions
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When a p orbital hybridizes with an s orbital
• Positive p lobe adds to s orbital generating larger lobe of sp3 hybrid
• Negative p lobe subtracts from s orbital generating smaller lobe of
sp3 hybrid
• Resulting asymmetrical sp3 hybrid is strongly oriented in one
direction
• Larger lobe of sp3 hybrid can overlap more effectively with an orbital
from another atom forming much stronger bonds than s or p orbitals
sp3 Hybrid Orbitals and the Structure of
Methane
When each of the four identical sp3 hybrid orbitals of carbon
overlaps with the 1s orbital of a hydrogen atom, four
identical C-H bonds are formed and methane results
• Each C-H bond in methane has a strength of 436 kJ/mol and a length of
109 pm
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The four C-H bonds of methane have a specific geometry
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The angle formed between two adjacent bonds describes a
characteristic bond angle
1.7 sp3 Hybrid Orbitals and the Structure
of Ethane
Orbital hybridization accounts for the bonding together of carbon
atoms into chains and rings
Ethane C2H6
• Tetrahedral
• Bond angles are near 109.5º
• Carbon-carbon single bond
• Formed by
overlap of sp3
hybrids from each carbon
• The remaining sp3 hybrids of
each carbon overlap with 1s
orbitals of three hydrogen
atoms to form six carbonhydrogen bonds
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