Transcript Lecture - Ch 1

Chapter 1
Structure and
Bonding
Suggested Problems 1-17,23-4,30,36,43,56,57
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What is Organic Chemistry?
• Organic compounds: from living organisms
(with a vital force)
• Inorganic compounds: from minerals
(without a vital force)
organic chemistry = compounds that contain carbon
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Carbon Compounds Are Ubiquitous
• Living things are
made of organic
chemicals
• Proteins – hair,
cell structures
• Nucleic Acids genetic make-up
• Carbohydrates &
fats - energy
• Medicines
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What Makes Carbon So Special?
• Atoms to the left of carbon give up electrons.
• Atoms to the right of carbon accept electrons.
• Carbon shares electrons.
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The Structure of an Atom
Protons are positively charged.
Neutrons have no charge.
Electrons are negatively charged.
Atomic number = # of protons
Atomic number of carbon = 6
Neutral carbon has six protons
and six electrons.
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Isotopes
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The Distribution of Electrons in an Atom
• The first shell is closest to the nucleus.
• The closer the atomic orbital is to the nucleus,
the lower its energy.
• Within a shell, s < p.
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Relative Energies of the Atomic Orbitals
We will not be utilizing d orbitals in this course.
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Shapes of Atomic Orbitals
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-Aufbau principle: An electron goes into the atomic orbital with
the lowest energy.
-Pauli exclusion principle: No more than two electrons can be
in an atomic orbital.
-Hund’s rule: An electron goes into an empty degenerate
orbital rather than pairing up.
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Atoms on the Left Side of the
Periodic Table Lose an Electron
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Atoms on the Right Side of the
Periodic Table Gain an Electron
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An Ionic Bond is Formed by the Attraction
Between Ions of Opposite Charge
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Covalent Bonds are Formed by
Sharing Electrons
Nonpolar covalent bond = bonded atoms are the same
Polar covalent bond = bonded atoms are different
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Bond Polarity Depends on
the Difference in Electronegativity
2.0
0.5
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Polar Covalent Bonds
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Dipole Moments
the greater the difference in electronegativity,
the greater the dipole moment, and the more polar the bond
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Electrostatic Potential Maps
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Formal Charge
Formal Charge = # of valence electrons –
(# of bonds + # of lone-pair electrons)
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How Many Bonds Does
an Atom Form?
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Neutral Carbon Forms Four Bonds
if carbon does not form four bonds, it has a charge
(or it is a radical)
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Neutral Nitrogen Forms Three Bonds
Nitrogen has one lone pair.
If nitrogen does not form three bonds, it is charged.
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Neutral Oxygen Forms Two Bonds
Oxygen has two lone pairs.
If oxygen does not form two bonds, it is charged.
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Hydrogen and the Halogens
Form One Bond
A halogen has three lone pairs.
if hydrogen or halogen does not form one bond, it has a charge
(or it is a radical)
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The Number of Bonds and Lone Pairs
• Halogen = 3 lone pairs
• Oxygen = 2 lone pairs
• Nitrogen = 1 lone pair
You might want to know this
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Lewis Structures
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How to Draw a Lewis Structure
NO3–
Determine the total number of valence electrons (5 + 6 + 6 + 6 = 23).
Because they are negatively charged, add another electron = 24.
Avoid O–O bonds.
Check for formal charges.
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Kekulé Structures and
Condensed Structures
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What is an Atomic Orbital?
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The Lobes of a p Orbital Have
Opposite Phases
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The Three p Orbitals
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The Bonding in H2
• A 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
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 (s) bond
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The Bonding in H2
Molecular Orbital Theory:
• Covalent bonds result when atomic orbitals combine to form
molecular orbitals.
• In analogy to an atomic orbital, a molecular orbital describes
the volume of space around a molecule where an electron is
likely to be found.
-- An atomic orbital surrounds an atom; a molecular orbital
surrounds a molecule.
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Waves Can Reinforce Each Other;
Waves Can Cancel Each Other
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Atomic Orbitals Combine
to Form Molecular Orbitals
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Side-to-Side Overlap of In-Phase
p Orbitals Forms a π Bond
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The Bonding in Methane
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In Order to Form Four Bonds,
Carbon Must Promote an Electron
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Four Orbitals are Mixed to Form
Four Hybrid Orbitals
An sp3 orbital has a large lobe and a small lobe.
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The Carbon in Methane is sp3
Carbon is tetrahedral.
The tetrahedral bond angle is 109.5°.
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The Bonding in Ethane
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The Bonding in Ethane
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End-On Overlap of Orbitals
Forms a (σ) Bond
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Bonding in Ethene
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An sp2 Carbon Has Three sp2 Orbitals
and One p Orbital
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The Carbons in Ethene are sp2
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Bonding in Ethyne
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The Two sp Orbitals Point in Opposite Directions
The Two p Orbitals are Perpendicular
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The Carbons in Ethyne are sp
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Methyl Anion is sp3
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Nitrogen Has Three Unpaired Valence Electrons
and Forms Three Bonds in NH3
Nitrogen does not have to promote an electron.
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The Bonds in Ammonia (NH3)
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The Ammonium Ion (+NH4)
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Oxygen Has Two Unpaired Valence
Electrons and Forms Two Bonds in H2O
Oxygen does not have to promote an electron.
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The Bonds in Water (H2O)
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Summary of Hybridization
orbitals used in bond formation determine the bond angle
You might want to know this
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Single Bond: 1 σ + 0 π
Double bond: 1 σ + 1 π
Triple Bond: 1 σ + 2 π
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Hybridization of C, N, and O
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Bond Strength and Bond Length
The shorter the bond, the stronger it is.
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s Character
The more s character, the shorter and stronger the bond.
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The More s Character in the Orbital,
the Shorter and Stronger the Bond
The more s character, the shorter and stronger the bond.
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The More s Character in the Orbital,
the Greater the Bond Angle
The more s character, the greater the bond angle.
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Hybridization, Bond Angle,
Bond Length, Bond Strength
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Summary
• The shorter the bond, the stronger it is.
• The greater the electron density in the region of orbital
overlap, the stronger the bond.
• The more s character, the shorter and stronger the bond.
• The more s character, the larger the bond angle.
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A π Bond is Weaker Than a σ Bond
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