Transcript Slide 1

1.1 Organic Chemistry
• The study of carbon-containing molecules and their
reactions
• What happens to a molecule during a reaction?
– A collision
– Bonds break/form
• WHAT IS A BOND?
• The BIG question: WHY do reactions occur?
– We will need at least 2 semesters of your time to answer
this question
– FOCUS ON THE ELECTRONS
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Klein, Organic Chemistry 1e
1.1 Organic Chemistry
• Why do we distinguish between organic and
inorganic compounds?
• Why are organic compounds important?
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1.2 Structural Theory
• In the mid 1800s, it was first suggested that
substances are defined by a specific arrangement of
atoms.
– Why is a compound’s formula NOT adequate to define it?
• What term do we use to describe different
substances with the same formula?
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1.2 Structural Theory
• Atoms that are most commonly bonded to carbon
include N, O, H, and halides (F, Cl, Br, I).
• With some exceptions, each element generally forms
a specific number of bonds with other atoms
• Practice with skill builder 1.1
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1.3 Covalent Bonding
• A covalent bond is a PAIR of electrons shared between
two atoms. For example…
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1.3 Covalent Bonding
• How do potential energy and stability relate?
• What forces keep the bond at the optimal length?
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1.3 Atomic Structure
• A review from General Chemistry
– Protons (+1) and neutrons (neutral) reside in the nucleus
– Electrons (-1) reside outside the nucleus. WHERE?
– Some electrons are close to the nucleus and others are far
away, WHY?
– Look at carbon for example. Which electrons are the valence
electrons?
– Why are valence electrons important?
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1.3 Counting Valence Electrons
• You can always calculate the number of valence electron
by analyzing the e- configuration. Look at phosphorus.
• Or, for Group A elements only, just look at the Group
number (Roman Numeral) on the periodic table
• Practice with skill builder 1.2
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1.3 Simple Lewis Structures
• For simple Lewis Structures…
1. Draw the individual atoms using dots to represent the
valence electrons.
2. Put the atoms together so they share PAIRS of electrons to
make complete octets. WHAT is an octet?
• Take NH3, for example…
• Practice with skill builder 1.3
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1.3 Simple Lewis Structures
• For simple Lewis Structures…
1. Draw the individual atoms using dots to represent the
valence electrons.
2. Put the atoms together so they share PAIRS of electrons to
make complete octets. WHAT is an octet?
• Try drawing the structure for C2H2
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1.4 Formal Charge
• What term do we use to describe atoms with an
unbalanced or FORMAL charge?
• How does formal charge affect the stability of an atom?
• Atoms in molecules (sharing electrons) can also have
unbalanced charge, which must be analyzed, because it
affects stability
• To calculate FORMAL charge for an atom, compare the
number of valence electrons that should be associated
with the atom to the number of valence electrons that
are actually associated with an atom
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Klein, Organic Chemistry 1e
1.4 Formal Charge
• Lets look at a specific formal charge example. Given the
Lewis structure, calculate the formal charge on each
atom.
or
• Carbon should have 4 valence e-s, because it is in group
IVA on the periodic table.
• Carbon actually has 8 valence e-s. It needs 8 for its
octet, but only 4 count towards its charge. WHY?
• The 4 it actually has balance out the 4 it should have, so
it does not have formal charge. Its neutral.
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1.4 Formal Charge
• Lets look at the formal charge of the oxygen atom.
or
• Oxygen should have 6 valence e-s, because it is in group
VIA on the periodic table.
• It actually has 8 valence e-s. It needs 8 for its octet, but
only 7 count towards its charge. WHY?
• If it actually has 7, but it should only have 6, what is its
formal charge?
• Practice with skill builder 1.4
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1.5 Polar Covalent Bonds
• Covalent bonds are electrons pairs that exist in an orbital
shared between two atoms. What do you think that
orbital looks like?
• Just like an atomic orbital, the electrons could be
anywhere within that orbital region.
• What factors determine which atom in the bond will
attract the shared electrons more?
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1.5 Polar Covalent Bonds
• Covalent bonds are either polar or nonpolar
– Nonpolar Covalent –bonded atoms share electrons evenly
– Polar Covalent – One of the atoms attracts electrons more than
the other
• Electronegativty - how strongly an atom attracts shared
electrons
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1.5 Polar Covalent Bonds
• Electrons tend to shift away from lower electronegativity
atoms to higher electronegativity atoms.
• The greater the difference in electronegativity, the more
polar the bond.
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1.5 Polar Covalent Bonds
• Can a bond have both covalent and ionic character?
• Practice with skill builder 1.5
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1.6 Atomic Orbitals
General Chemistry review
• In the 1920s, Quantum Mechanics was established as a
theory to explain the wave properties of electrons
• The solution to wave equations for electrons provides us
with visual pictures called orbitals
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1.6 Atomic Orbitals
General Chemistry review
• The type or orbital be identified by its shape
• An orbital is a region where there is a calculated 90%
probability of finding an electron. The remaining 10%
probability tapers off as you move away from the
nucleus
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1.6 Atomic Orbitals
• Electrons behave as both particles and waves. How can
they be BOTH? Maybe the theory is not yet complete
• The theory does match experimental data, and it has
predictive capability.
– Like a wave on a lake, an electron’s wavefunction can be (+),
(-), or ZERO.
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Klein, Organic Chemistry 1e
1.6 Atomic Orbitals
• Because they are generated
mathematically from wavefunctions,
orbital regions can also be (-), (+), or
ZERO
– The sign of the wave function has nothing
to do with electrical charge.
• In this p-orbital, there is a nodal plane.
The sign of the wavefunction will be
important when we look at orbital
overlapping in bonds.
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1.6 Atomic Orbitals
• Electrons are most stable (lowest in
energy) if they are in the 1s orbital?
• The 1s orbital is full once there are two
electrons in it. Why can’t it fit more?
• The 2s orbital is filled next. The 2s
orbital has a node. WHERE?
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1.6 Atomic Orbitals
• Once the 2s is full, electrons fill into the three
degenerate 2p orbitals
• Where are the nodes in each of the 2p orbitals?
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1.6 Atomic Orbitals
• Common elements and their electron configurations
• Practice with skill builder 1.6
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Klein, Organic Chemistry 1e
1.6 Atomic Orbitals
• What are the rules that govern our placement of
electrons ?
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1.7 Valence Bond Theory
• A bond occurs when atomic orbitals overlap.
Overlapping orbitals is like overlapping waves
• Only constructive interference results in a bond
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1.7 Valence Bond Theory
• The bond for a H2 molecule results from constructive
interference
• Where do the bonded electrons spend most of their
time?
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1.8 Molecular Orbital Theory
• Atomic orbital wavefunctions
overlap to form MOs that
extend over the entire
molecule.
• MOs are a more complete
analysis of bonds, because
they include both constructive
and destructive interference.
• The number of MOs created
must be equal to the number
of AOs that were used.
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H2 MOs
Klein, Organic Chemistry 1e
1.8 Molecular Orbital Theory
• Why is the antibonding orbital higher in energy?
• When the AOs overlap, why do the electrons go into the
bonding MO rather than the antibonding MO?
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1.8 Molecular Orbital Theory
• Imagine a He2 molecule. How would its MOs compare
to those for H2?
• How would the energy of the
He2 compare to 2 He?
• Why does Helium exist in its
atomic form rather than in
molecular form?
• In general, if a molecule has all of it MOs occupied, will
be stable or unstable?
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1.8 Molecular Orbital Theory
• Consider the MOs for CHBr3
– There are many areas of atomic orbital overlap
– Notice how the MOs extend over the entire molecule
• Each picture below represents ONE orbital.
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1.8 Molecular Orbital Theory
• How many electrons can fit into the areas represented in
(b)?
• Depending on the circumstances, we will use both MO
and valence bond theory to explain phenomena.
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1.9 Hybridized Atomic Orbitals
• Given the electron configuration for C and H, imagine
how their atomic orbitals might overlap
• Would such orbital overlap
yield methane?
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1.9 Hybridized Atomic Orbitals
• To make methane, the C atom must have 4 atomic
orbitals available for overlapping
• If an electron is excited from the 2s to the 2p, will that
make it suitable for making methane?
• If four H atoms were to come in and overlap with the 2s
and 2p orbitals, what geometry would the resulting
methane have?
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1.9 Hybridized Atomic Orbitals
• The carbon must undergo hybridization to form 4 equal
atomic orbitals
• The atomic orbitals must be equal in energy to form four
equal-energy symmetrical C-H bonds
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Klein, Organic Chemistry 1e
1.9 Hybridized Atomic Orbitals
• Should the shape of an sp3 orbital look more like an s or
more like p orbital?
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Klein, Organic Chemistry 1e
1.9 Hybridized Atomic Orbitals
• To make CH4, the 1s atomic orbitals of the H atoms will
overlap with the four sp3 hybrid atomic orbitals of C
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1.9 Hybridized Atomic Orbitals
• Draw a picture that shows the necessary atomic orbitals
and their overlap to form ethane (C2H6).
• Draw a picture that shows the necessary atomic orbitals
and their overlap to form water.
• Practice with conceptual checkpoint 1.19
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1.9 Hybridized Atomic Orbitals
• Consider ethene (ethylene).
• Each carbon in ethene must bond to three other atoms,
so only three hybridized atomic orbitals are needed
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1.9 Hybridized Atomic Orbitals
• An sp2 hybridized carbon will have three equal-energy
sp2 orbitals and one unhybridized p orbital
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Klein, Organic Chemistry 1e
1.9 Hybridized Atomic Orbitals
• The sp2 atomic orbitals overlap to form sigma (σ) bonds
• Sigma bonds provide maximum HEAD-ON overlap
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Klein, Organic Chemistry 1e
1.9 Hybridized Atomic Orbitals
• The unhybridized p orbitals in ethene form pi (π) bonds,
SIDE-BY-SIDE overlap
• Practice with conceptual checkpoint 1.20
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1.9 Hybridized Atomic Orbitals
• The unhybridized p orbitals in ethene form pi (π) bonds,
SIDE-BY-SIDE overlap
• MO theory provides a similar picture. Remember, red
and blue regions are all part of the same orbital
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1.9 Hybridized Atomic Orbitals
• Why is sp2 hybridization not appropriate for methane
(CH4)?
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1.9 Hybridized Atomic Orbitals
• Consider ethyne (acetylene).
• Each carbon in ethyne must bond to two other atoms, so
only two hybridized atomic orbitals are needed
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1.9 Hybridized Atomic Orbitals
• The sp atomic orbitals overlap HEAD-ON to form sigma
(σ) bonds while the unhybridized p orbitals overlap SIDEBY-SIDE to form pi bonds
• Practice with skill
builder 1.7
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1.9 Hybridized Atomic Orbitals
• Which should be stronger, a pi bond or a sigma bond?
WHY?
• Which should be stronger, an sp3 – sp3 sigma bond
overlap or an sp – sp sigma bond overlap?
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1.9 Hybridized Atomic Orbitals
• Explain the different strengths and lengths below.
• Practice with
conceptual
checkpoint
1.24
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Klein, Organic Chemistry 1e
1.10 Molecular Geometry
• Valence shell electron pair repulsion (VSEPR theory)
– Valence electrons (bonded and lone pairs) repel each other
• To determine molecular geometry…
1. Determine the Steric number
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1.10 Molecular Geometry
• Valence shell electron pair repulsion (VSEPR theory)
– Valence electrons (bonded and lone pairs) repel each other
• To determine molecular geometry…
2. Predict the hybridization of the central atom
•
•
•
If the Steric number is 4, then it is sp3
If the Steric number is 3, then it is sp2
If the Steric number is 2, then it is sp
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1.10
3
sp
Geometry
• For any sp3 hybridized atom, the 4 valence electron pairs
will form a tetrahedral electron group geometry
• Methane has 4
equal bonds,
so the bond
angels are
equal
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• How does the
lone pair of
ammonia
affect its
geometry?
1-51
• The bond
angels in
oxygen are
even smaller,
why?
Klein, Organic Chemistry 1e
1.10
3
sp
Geometry
• The molecular geometry is different from the electron
group geometry. HOW?
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1.10
2
sp
Geometry
• Calculate the Steric number for BF3
• Electron pairs that are located in sp2 hybridized orbitals
will form a trigonal planar electron group geometry
• What will be the molecular geometry?
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1.10
2
sp
Geometry
• How many electrons are in Boron’s unhybridized p
orbital?
• Does this geometry follow VSEPR theory?
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1.10
2
sp
Geometry
Analyze the steric number, hybridization, electron group
geometry and molecular geometry for this imine?
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Klein, Organic Chemistry 1e
1.10 sp Geometry
• Analyze the Steric number, the hybridization, the
electron group geometry, and the molecular geometry
for the following molecules
• BeH2
• CO2
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1.10 Geometry Summary
• Practice with
skill builder
1.8
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Klein, Organic Chemistry 1e
1.11 Molecular Polarity
• Electronegativity Differences cause induction
• Induction (shifting of electrons WITHIN their orbitals)
results in a dipole moment.
• Dipole moment = (the amount of partial charge) x (the
distance the δ+ and δ- are separated)
• Dipole moments are reported in units of debye (D)
• 1 debye = 10-18 esu ∙ cm
– An esu is a unit of charge. 1 e- has a charge of 4.80 x 10-10 esu
– cm are included in the unit, because the distance between the
centers of + and – charges affects the dipole
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1.11 Molecular Polarity
• Lets consider the dipole of CH3Cl
• Dipole moment (μ) = charge (e) x distance (d)
– Plug in the charge and distance
• μ = (1.056 x 10-10 esu) x (1.772 x 10-8 cm)
– Note that the amount of charge separation is
less than what it would be if it were a full
charge separation (4.80 x 10-10 esu)
• μ = 1.87 x 10-18 esu ∙ cm
– Convert to debye
• μ = 1.87 D
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1.11 Molecular Polarity
• What would the dipole moment be if CH3Cl were 100%
ionic?
• μ = charge (e) x distance (d)
– Plug in the charge and distance
• μ = (4.80 x 10-10 esu) x (1.772 x 10-8 cm)
– The full charge of an electron is plugged in
• μ = 8.51 x 10-18 esu ∙ cm = 8.51 D
• What % of the C-Cl bond is ionic?
• Is the C-Cl bond mostly ionic or mostly
covalent?
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1.11 Molecular Polarity
• Check out the polarity of come other common bonds
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1.11 Molecular Polarity
• Why is the C=O double bond so much more polar than
the C-O single bond?
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1.11 Molecular Polarity
• For molecules with multiple polar bonds, the dipole
moment is the vector sum of all of the individual bond
dipoles
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1.11 Molecular Polarity
• It is important to determine a molecule’s geometry
FIRST before analyzing its polarity
• If you have not drawn the molecule with the proper
geometry, it may cause you to aasses the polarity wrong
as well
• Would the dipole for water be different if it were linear
rather than angular?
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1.11 Molecular Polarity
• Practice with skill builder 1.9
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1.11 Molecular Polarity
• Explain why the dipole moment for pentane = 0 D
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Klein, Organic Chemistry 1e
1.12 Intermolecular Forces
• Many properties such as solubility, boiling point, density,
state of matter, melting point, etc. are affected by the
attractions BETWEEN molecules
• Neutral molecules (polar and nonpolar) are attracted to
one another through…
– Dipole-dipole interactions
– Hydrogen bonding
– Dispersion forces (a.k.a. London forces or fleeting dipoledipole forces)
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Klein, Organic Chemistry 1e
1.12 Dipole-Dipole
• Dipole-dipole forces result when polar molecules line up
their opposite charges.
• Note acetone’s permanent dipole results from the
difference in electronegativity between C and O
• The dipole-dipole attractions BETWEEN acetone
molecules affects acetone’s boiling and melting points.
HOW?
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1.12 Dipole-Dipole
• Why do isobutylene and acetone have such different MP
and BPs?
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1.12 Hydrogen Bonding
• Hydrogen bonds are an especially strong type of dipoledipole attraction
• Hydrogen bonds are strong because the partial + and –
charges are relatively large
• Why are the partial charges in the H-bonding examples
below relatively large?
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1.12 Hydrogen Bonding
• Only when a hydrogen shares electrons with a highly
electronegative atom (O, N, F, or Cl) will it carry a large
partial positive charge
• The large δ+ on the H atom can attract large δ– charges
on other molecules
• Even with the large partial charges, H-bonds are still
about 20 times weaker than covalent bonds
• Compounds with H atoms that are capable of forming Hbonds are called protic
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1.12 Hydrogen Bonding
• Which of the following solvents are protic (capable of Hbonding), and which are not?
• Acetic acid
• Diethyl ether
• Methylene chloride (CH2Cl2)
• Dimethyl sulfoxide
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1.12 Hydrogen Bonding
• Explain why the following isomers have different boiling
points
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Klein, Organic Chemistry 1e
1.12 London Dispersion Forces
• If two molecules are nonpolar (dipole = 0 D), will they
attract one another?
– YES! HOW?
• Nonpolar molecules normally have their electrons (-)
spread out evenly around the nuclei (+) completely
balancing the charge
• However, the electrons are in constant random motion
within their MOs
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1.12 London Dispersion Forces
• The constant random motion of the electrons in the
molecule will sometimes produce an electron
distribution that is NOT evenly balanced with the
positive charge of the nuclei
• Such uneven distribution produces a temporary dipole,
which can induce a temporary dipole in a neighboring
molecule
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Klein, Organic Chemistry 1e
1.12 London Dispersion Forces
• The result is a fleeting attraction between the two
molecules
• Such fleeting attractions are generally weak.
• But like any weak attraction, if there are enough of
them, they can add up to a lot
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1.12 London Dispersion Forces
• The greater the surface area of a molecule, the more
temporary dipole attractions are possible
• Consider the feet of Gecko. They have many flexible
hairs on their feet that maximize surface contact
• The resulting London dispersion forces are
strong enough to support the weight of the
Gecko
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Klein, Organic Chemistry 1e
1.12 London Dispersion Forces
• Explain why molecules with more mass generally have
higher boiling points
• Practice with skill builder 1.10
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1.12 London Dispersion Forces
• Explain why more highly branched molecules generally
have lower boiling points
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Klein, Organic Chemistry 1e
1.13 Solubility
• We use the principle, like-dissolves-like
• Polar compounds generally mix well with other polar
compounds
– If the compounds mixing are all capable of H-bonding and/or
strong dipole-dipole, then there is no reason why they
shouldn’t mix
• Nonpolar compounds generally mix well with other
nonpolar compounds
– If none of the compounds are capable of forming strong
attractions, then no strong attractions would have to be
broken to allow them to mix
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1.13 Solubility
• We know it is difficult to get a polar compound (like
water) to mix with a nonpolar compound (like oil)
– We can’t use just water to wash oil off our dirty cloths
• To remove nonpolar oils, grease, and dirt, we need soap
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1.13 Solubility
• Soap molecules organize into micelles in water, which
form a nonpolar interior to carry away dirt.
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1.13 Solubility
• Which attraction is generally stronger?
– The attraction between a permanent dipole and an induced
dipole
or
– The attraction between a temporary dipole and an induced
dipole
• Which attraction is generally stronger?
– The attraction between a polar molecule and a nonpolar
molecule
or
– The attraction between two nonpolar molecules?
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1.13 Solubility
• Why won’t a nonpolar compound readily dissolve in
water?
• Is it because the water molecules repel the nonpolar
molecules?
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