CEM 241

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CEM 241

• Chapter 1

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3 x 5 Cards

• Front Name
• Address
• Phone Number
• E-Mail
• Major
• Back Something about you.

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CEM 241 Organic Chemistry

• Instructor John Singer
• Office JM 232A
• Phone [phone] (office)
• [phone] (cell)
• (no calls after 1000 p.m. please)

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Atoms are composed of

• Protons
• positively charged
• mass 1.6726 X 10-27 kg
• Neutrons
• neutral
• mass 1.6750 X 10-27 kg
• Electrons
• negatively charged
• mass 9.1096 X 10-31 kg

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Atomic Number and Mass Number

• Atomic number (Z) number of protons in nucleus
• (this must also equal the number of electrons in
  neutral atom)
• Mass number (A) sum of number of protons
  neutrons in nucleus

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Schrödinger Equation

• Schrödinger combined the idea that an electron
  has wave properties with classical equations of
  wave motion to give a wave equation for the
  energy of an electron in an atom.
• Wave equation (Schrödinger equation) gives
  aseries of solutions called wave functions (? ).

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Wave Functions

• Only certain values of ? are allowed.
• Each ? corresponds to a certain energy.
• The probability of finding an electron at a
  particular point with respect to the nucleus
  isgiven by ? 2.
• Each energy state corresponds to an orbital.

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Figure 1.1 Probability distribution (? 2) for an
electron in a 1s orbital.
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A boundary surface encloses the regionwhere the
probability of finding an electronis highon the
order of 90-95
1s
2s
Figure 1.2 Boundary surfaces of a 1s orbitaland
a 2s orbital.
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Quantum Numbers

• Each orbital is characterized by a unique set
  of quantum numbers.
• The principal quantum number n is a wholenumber
  (integer) that specifies the shell and isrelated
  to the energy of the orbital.
• The angular momentum quantum number is usually
  designated by a letter (s, p, d, f, etc) and
  describes the shape of the orbital.

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s Orbitals

• s Orbitals are spherically symmetric.
• The energy of an s orbital increases with
  thenumber of nodal surfaces it has.
• A nodal surface is a region where the
  probabilityof finding an electron is zero.
• A 1s orbital has no nodes a 2s orbital has
  onea 3s orbital has two, etc.

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The Pauli Exclusion Principle

• No two electrons in the same atom can havethe
  same set of four quantum numbers.
• Two electrons can occupy the same orbitalonly
  when they have opposite spins.
• There is a maximum of two electrons per orbital.

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p Orbitals

• p Orbitals are shaped like dumbells.
• Are not possible for n 1.
• Are possible for n 2 and higher.
• There are three p orbitals for each value of n
  (when n is greater than 1).

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Ionic Bonding

• An ionic bond is the force of electrostaticattrac
  tion between oppositely charged ions

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Ionic Bonding

• Ionic bonds are common in inorganic
  chemistrybut rare in organic chemistry.
• Carbon shows less of a tendency to form
  cationsthan metals do, and less of a tendency to
  formanions than nonmetals.

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The Lewis Model of Chemical Bonding

• In 1916 G. N. Lewis proposed that atomscombine
  in order to achieve a more stableelectron
  configuration.
• Maximum stability results when an atomis
  isoelectronic with a noble gas.
• An electron pair that is shared between two
  atoms constitutes a covalent bond.

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Covalent Bonding in H2
Two hydrogen atoms, each with 1 electron,
can share those electrons in a covalent bond.

• Sharing the electron pair gives each hydrogen an
  electron configuration analogous to helium.

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Example
Combine carbon (4 valence electrons) andfour
fluorines (7 valence electrons each)
to write a Lewis structure for CF4.
The octet rule is satisfied for carbon and each
fluorine.
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Double and Triple Bonds
Ethylene
Acetylene
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Electronegativity
Electronegativity is a measure of the abilityof
an element to attract electrons toward itself
when bonded to another element.

• An electronegative element attracts electrons.
• An electropositive element releases electrons.

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Pauling Electronegativity Scale

• Electronegativity increases from left to rightin
  the periodic table.
• Electronegativity decreases going down a group.

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Bond Polarity

• Bonds will be polar when there are differences in
  electronegativity between the atoms involved.
• The greater the difference in electronegativity,
  the greater the polarity.

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Electrostatic Potential Maps

• Electrostatic potential maps show the
  chargedistribution within a molecule.

?
?-
Li
H
Red is negative chargeblue is positive.
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Formal Charge
Formal Charge is defined as the charge that an
Atom would have at the covalent limit.
Formal charge
group numberin periodic table
number ofbonds
number ofunshared electrons


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Nitric acid
Formal charge of H
..

• We will calculate the formal charge for each atom
  in this Lewis structure.

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Nitric acid
Formal charges


..

• A Lewis structure is not complete unless formal
  charges (if any) are shown.

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Condensed structural formulas

• Lewis structures in which many (or all) covalent
  bonds and electron pairs are omitted.

can be condensed to
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Bond-line formulas

• Omit atom symbols. Represent structure by
  showing bonds between carbons and atoms other
  than hydrogen.
• Atoms other than carbon and hydrogen are called
  heteroatoms.

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Bond-line formulas
is shown as
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Constitutional isomers

• Isomers are different compounds that have the
  same molecular formula.
• Constitutional isomers are isomers that differ
  in the order in which the atoms are connected.
• An older term for constitutional isomers is
  structural isomers.

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A Historical Note
NH4OCN
Urea
Ammonium cyanate

• In 1823 Friedrich Wöhler discovered that when
  ammonium cyanate was dissolved in hot water, it
  was converted to urea.
• Ammonium cyanate and urea are constitutional
  isomers of CH4N2O.
• Ammonium cyanate is inorganic. Urea is
  organic. Wöhler is credited with an important
  early contribution that helped overturn the
  theory of vitalism.

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Writing Lewis Structures

• 1. Sketch a skeleton.
• 2. Count Valence Electrons
• 3. Subtract bonding e- from valence e-.
• 4. Place remaining electrons as lone pairs to
  fulfill octet rule.
• 5. Use multiple bonds to complete octets.

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Resonance

• When more than one Lewis Structure can be drawn
  for a molecule that differs only by placement of
  electrons, the structure is said to possess
  resonance.

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Resonance Structures of Methyl Nitrite

• same atomic positions
• differ in electron positions

more stable Lewis structure
less stable Lewis structure
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Example

• Ozone (O3)
• Lewis structure of ozone shows one double bond
  and one single bond

Expect one short bond and one long
bond Reality bonds are of equal length (128 pm)
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Shapes of Molecules Valence Shell Electron Pair
Repulsions

• The most stable arrangement of groups attached
  to a central atom is the one that has the
  maximum separation of electron pairs(bonded or
  nonbonded).

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Dipole Moment

• A substance possesses a dipole moment if its
  centers of positive and negative charge do not
  coincide.
• ? e x d
• (expressed in Debye units)

not polar
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Figure 1.7
Resultant of thesetwo bond dipoles is
Resultant of thesetwo bond dipoles is
? 0 D
Carbon tetrachloride has no dipolemoment
because all of the individualbond dipoles cancel.
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Figure 1.7
Resultant of thesetwo bond dipoles is
Resultant of thesetwo bond dipoles is
? 1.62 D
The individual bond dipoles do notcancel in
dichloromethane it hasa dipole moment.
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Definitions

• Arrhenius
• An acid ionizes in water to give protons. A base
  ionizes in water to give hydroxide ions.
• Brønsted-Lowry
• An acid is a proton donor. A base is a proton
  acceptor.
• Lewis
• An acid is an electron pair acceptor. A base is
  an electron pair donor.

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Acid Strength is Measured by pKa
pKa log10Ka
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A Brønsted Acid-Base Reaction
A proton is transferred from the acid to the base.


.
.
B

H
A
H
A
B

base
acid
conjugate acid
conjugate base
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Equilibrium Constant for Proton Transfer
..

.
.

O

H
Br
H
O
..
H3OBr
Ka
HBr
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Dissociation Constants (pKa) of Acids
strong acids are stronger than hydronium ion
For a more detailed list click here for Table 1.7
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Example

• Which is the stronger base, ammonia (left) or
  pyridine (right)?
• Recall that the stronger the acid, the weaker the
  conjugate base.
• Therefore, the stronger base is the conjugate of
  the weaker acid.
• Look up the pKa values of the conjugate acids of
  ammonia and pyridine in Table 1.7.

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Example
H
H

weaker acid
N
H
pKa 9.3
H
pKa 5.2
stronger acid
Therefore, ammonia is a stronger base than
pyridine
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The Main Ways Structure Affects Acid Strength

• The strength of the bond to the atom from which
  the proton is lost.
• The electronegativity of the atom from which the
  proton is lost.
• Changes in electron delocalization on ionization.

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Bond Strength
Bond strength is controlling factor when
comparing acidity of hydrogen halides.
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Electronegativity
Electronegativity is controlling factor when
comparing acidity of protons bonded to atoms in
the same row of the periodic table.
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Electronegativity
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Electronegativity
The equilibrium becomes more favorable as A
becomes better able to bear a negative
charge. Another way of looking at it is that H
becomes more positive as the atom to which it is
attached becomes more electronegative.
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Bond strength versus Electronegativity
Bond strength is more important when comparing
acids in which the proton that is lost is bonded
to atoms in the same group of the periodic
table. Electronegativity is more important when
comparing acids in which the proton that is lost
is bonded to atoms in the same row of the
periodic table.
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Inductive Effect
The greater acidity of CF3CH2OH compared to
CH3CH2OH is an example of an inductive
effect. Inductive effects arise by polarization
of the electron distribution in the bonds between
atoms.
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Electron Delocalization
Ionization becomes more favorable if electron
delocalization increases in going from right to
left in the equation. Resonance is a convenient
way to show electron delocalization.
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Electron Delocalization
Ionization becomes more favorable if electron
delocalization increases in going from right to
left in the equation. Resonance is a convenient
way to show electron delocalization. (Example
Acetic Acid)
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Acid/Base Equilibria
The equilibrium in an acid-base reaction is
favorable if the stronger acid is on the left and
the weaker acid is on the right.
Stronger acid Stronger base
Weaker acid Weaker base
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Important Points
A strong acid is one that is stronger than
H3O.A weak acid is one that is weaker than
H3O. A strong base is one that is stronger than
HO.A weak base is one that is weaker than
HO. The strongest acid present in significant
quantities when a strong acid is dissolved in
water is H3O. The strongest acid present in
significant quantities when a weak acid is
dissolved in water is the weak acid itself.
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Example Two Neutral Molecules
F3B

Lewis acid
Lewis base
Product is a stable substance. It is a liquid
witha boiling point of 126C. Of the two
reactants,BF3 is a gas and CH3CH2OCH2CH3 with a
boiling point of 34C.
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Example Ion Neutral molecule
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The End