Transcript CHEM 210 Ch06x

Organic Chemistry
PRINCIPLES AND MECHANISMS
Chapter 6: Lecture PowerPoint
The Proton Transfer Reaction
6.1 An Introduction to Reaction Mechanisms:
The Proton Transfer Reaction
and Curved Arrow Notation
• A proton transfer reaction (or a Brønsted–Lowry acid–base
reaction) is one in which a Brønsted–Lowry base reacts with a
Brønsted–Lowry acid.
• Brønsted–Lowry bases accept protons (H+), and Brønsted–
Lowry acids donate the protons.
• A proton transfer reaction consists of a single elementary
step, in which the bonds break and form simultaneously.
Curved Arrow Notation
• This notation (also called arrow pushing) keeps account
of the valence electrons involved in the reaction
mechanism.
Curved Arrow Notation
continued…
• Bond breaking and bond formation involve only valence
electrons.
• Curved arrows show the appropriate movement of those
electrons.
6.2 Chemical Equilibrium and the
Equilibrium Constant, Keq
• A proton transfer reaction between HCl and HO⁻ will take
place readily (i.e., strong acid/strong base reaction).
• Essentially no proton transfer will occur between HO⁻
and NH3 (strong base/weak acid reaction).
• A reaction’s tendency to form products is described by its
equilibrium constant, Keq.
Proton Transfer and Keq
• For a proton transfer reaction between HA (the acid) and
B⁻ (the base), the equilibrium constant expression is
written as:
Acid Strengths: Ka and pKa
• The strength of an acid can be thought of as its ability to
donate a proton.
• Acid strengths can be obtained experimentally.
• The acidity constant, Ka, eliminates H2O from the equation.
• When comparing two acids, the one with the larger Ka value is
the stronger acid.
Acid Strengths: Ka and pKa
continued…
• Chemist often work with pKa instead of Ka due to Ka’s
immensely large range of values.
pKa = -log(Ka)
• As the value of Ka increases, pKa decreases.
• Compounds with low pKa values are more acidic than
compounds with high pKa values.
• Each difference of 1 in pKa values represents a factor of
10 difference in acid strength.
How to use a pKA table:
- The lower the pKA, the stronger the acid
- The higher the pKA the weaker the acid
- The lower the pKA, the weaker the CB
- The higher the pKA, the stronger the CB
- An acid can protonate any base derived
from an acid of higher pKA
- i.e. HI can protonate NH3 to NH4+
but
- H2O cannot protonate Cl- to HCl
- The CB of any acid can deprotonate any acid
of lower pKA :
- i.e. HO- can deprotonate NH4+ to NH3
but
- HO- cannot deprotonate NH3 to NH2-
Graphically for protonation:
Right up and back – yes
Right down and back – NO!
Graphically for deprotonation:
Left down and back – yes
Left up and back – NO!
Predicting the Outcome of a Proton
Transfer Reaction Using pKa Values
• Proton transfer reactions favor the side opposite the
stronger acid.
The Leveling Effect
• Even if a solvent does not act as an acid or a base in an
intended proton transfer reaction, it can limit the
existence of certain acids and bases in solution.
• These restrictions are the result of what is called the
leveling effect.
The Leveling Effect and Other Solvents
• The strongest acid that can exist in solution to any
appreciable concentration is the protonated solvent.
• The strongest base that can exist in solution is the
deprotonated solvent.
• In water, no acid stronger than H3O+ and no base
stronger than HO⁻ can exist to any appreciable extent.
Le Châtelier’s Principle and pH
• Le Châtelier’s principle states that if a reaction at
equilibrium experiences a change in reaction conditions
(e.g., concentrations, temperature, pressure, or volume),
then the equilibrium will shift to counteract that change.
Le Châtelier’s Principle and pH
continued…
• In water, increasing the concentration of H3O+ will cause the
equilibrium to shift toward reactants.
• At the new equilibrium, there will be more HA and less A⁻, so
the percent dissociation of HA will decrease.
• Additionally, decreasing the concentration of H3O+ will cause
the equilibrium to shift toward products, which results in an
increase in the percent dissociation of HA.
• The concentration of H3O+ in solution is reported as
pH = log[H3O+].
Henderson–Hasselbalch Equation
• The Henderson-Hasselbalch equation is a common
general chemistry equation used for dealing with buffers.
• At 50% dissociation of the acid, the equilibrium
concentrations of A⁻ (aq) and HA(aq) are equal.
• This makes the second term in Equation 6-16 equal to
zero, in which case pH = pKa.
Percent Dissociation of an Acid
with pKa = 5 as a Function of pH
6.3 Thermodynamics
and Gibbs Free Energy
• Equilibrium constant is also related to the standard Gibbs
free energy difference (DGorxn) between reactants and
products.
• The naught (o) signifies standard conditions (298 K,
1 atm, and the concentrations of all solutions are
1 mol/L)
6.3 Thermodynamics
and Gibbs Free Energy
continued…
• Because the gas constant (R) and temperature (T) are
both positive, a negative value for DGorxn results when
Keq > 1.
• Conversely, a positive value for DGorxn means that
Keq < 1.
• When Keq < 1 reactants are favored over products and
when Keq > 1, products are favored.
Enthalpy and Entropy
• The DGorxn can be expressed in terms of the reaction’s
enthalpy change (DHorxn) and entropy change (DSorxn).
• The DHorxn term is the standard enthalpy difference between
the reactants and products.
• At constant pressure, DHorxn equals the heat absorbed or
released by the reaction.
• If DHorxn is “+”, then the reaction absorbs heat (endothermic).
• If DHorxn is “-”, then the reaction releases heat (exothermic).
Enthalpy and Entropy
continued…
• DSorxn is the standard entropy difference between the
reactants and products.
• DSorxn is often thought of as a “measure of disorder.”
• DSorxn for a proton transfer reaction is usually very small.
The Reaction Free Energy Diagram
• In a reaction free-energy diagram, Gibbs free energy is plotted as a
function of the reaction coordinate.
• A reaction coordinate is a variable that corresponds to the changes
in geometry, on a molecular level, as reactants are converted into
products.
6.4 Strategies for Success:
Functional Groups and Acidity
• One way to estimate the
unknown pKa value of a
compound is to observe any
structural similarities to a
compound that is already
tabulated.
• The acidity (pKa) of a compound is
mostly governed by the
functional group on which the
acidic proton is found.
• Acidity is relatively independent
of the molecule’s carbon
skeleton, as observed by the
structures at the right.
Factors That Affect pKa
• Nearby structural features, such as a highly electronegative
substituent or adjacent double bond, can alter the acidity
significantly.
• Trichloroacetic acid (Cl3CCO2H; pKa = 0.77) is a stronger acid
than acetic acid (H3CCO2H; pKa = 4.75).
• Phenol is 106 times stronger than ethanol!
6.5 Relative Strengths of Charged and Uncharged
Acids: The Reactivity of Charged Species
• There are trends that dictate the strength of an acid.
• A proton is significantly more acidic when it is attached
to a positively charged atom than when that atom is
uncharged.
• The pKa of H3O+ is -1.7, whereas that of H2O is 15.7.
• The pKa of H4N+ is 9.4, whereas that of H3N is 36.
6.5 Relative Strengths of Charged and Uncharged
Acids: The Reactivity of Charged Species
continued…
• The stronger acidity of positively charged acids is a
reflection of lower stability of charged species than
similar uncharged species.
The Reactivity of Charged Species
• H3O+ and HO⁻ products, which are charged, are less stable
than the uncharged reactants by 80 kJ/mol.
6.6 Relative Acidities of Protons
on Atoms with Like Charges
• The farther to the right an atom appears in the periodic table,
the more acidic the protons that are attached to it.
pKa: CH4 > NH3 > H2O > HF
• Negative charge generated by loss of H+ is better stabilized on
the more electronegative atom.
6.6 Relative Acidities of Protons
on Atoms with Like Charges
continued…
Protons on Different Atoms in the Same
Column of the Periodic Table
• The farther down a column an atom
appears in the periodic table, the more
acidic the protons that are attached to it.
pKa: HF > HCl
• The stability of an anion increases when
the negative charge is on an atom farther
down a column of the periodic table.
• The negative charge is more distributed
over the larger atoms, which are found as
you go down the column (i.e., principal
quantum number increases).
Hybridization of the Atom to
Which the Proton Is Attached
• Ethane, ethene, and ethyne are all hydrocarbons, yet
differ in hybridization of their carbon atoms (sp3, sp2, and
sp, respectively).
• These three hydrocarbons are also uncharged acids.
• Their differences in acidity must be caused by differences
in the stability of their conjugate bases.
• In other words, HCΞC⁻ is more stable than H2C=CH⁻,
which is more stable than H3C-CH2⁻.
• The C atom bearing the negative charge in HCΞC⁻ is sp
hybridized (highest effective electronegativity = greatest
stability).
Effects from Adjacent Double
and Triple Bonds: Resonance Effects
• Ethanoic acid and ethanol are both uncharged acids yet
ethanoic acid is a considerably stronger acid.
• Ethanoic acid has a C=O bond adjacent to the acidic OH
group, and is more acidic than ethanol (CH3CH2OH) by about
11 pKa units.
• One of the main reasons ethanoic acid is so much more
acidic than ethanol has to do with resonance effects.
Effects from Adjacent Double
and Triple Bonds: Resonance Effects
continued…
• Ethanoic acid’s conjugate base is highly stabilized by
resonance, whereas ethanol’s negative charge is localized on
the O atom.
Resonance Stabilization
• The more delocalized a negative charge, the more stable the
conjugate base becomes.
• Resonance stabilization generally increases as the number of
atoms over which a charge is delocalized increases.
Resonance Stabilization:
H2SO4 vs Ethanoic Acid
• Sulfuric acid offers a greater number of electronegative
atoms to delocalize its negative charge (3 total) as
compared to ethanoic acid (2 total).
• This additional delocalized site helps explain why sulfuric
acid has a pKa = -9, while ethanoic acid’s pKa = 4.75!
Effects from Nearby Atoms:
Inductive Effects
• The presence of nearby electronegative atoms increases
the stability of the conjugate base.
• 2-Chloroethanol (ClCH2CH2OH; pKa = 12.9) is more acidic
than ethanol (CH3CH2OH; pKa = 16).
• The O atom bears less of a negative charge in ClCH2CH2O⁻
than it does in CH3CH2O⁻.
Inductive Effects
• Cl is electron withdrawing in comparison to H.
• The electron density on the left-most C atom of 2chloroethanol is shifted toward the Cl atom.
• To compensate for this shift, it would draw electron
density away from other atoms in which it is bonded.
• This effect is repeated down the chain until the electron
density is pulled from the O atom. This is called
induction.
Inductive Effects
continued…
Inductive Effects
continued…
Anion and Cation Stabilization
• Most atoms are more electronegative than H so their
presence is usually electron withdrawing.
• Some atoms are actually electron donating, when compared
to H (example: the Si atom).
• Alkyl groups are also electron donating.
Anion and Cation Stabilization
continued…
• Electron-donating groups stabilize nearby cations.
• The inductive effects of alkyl groups stabilizes the nearby
positive charge.
• Carbocations (C+) are increasingly stabilized by the
addition of more alkyl groups.
• The alkyl groups help diminish the positive charge by
inductive effects.
Carbocations and Their Stability
• Carbocations are distinguished by their degree of alkyl
substitution.
• Carbocation stability increases in the following order:
Methyl < 1o < 2o < 3o
6.7 Strategies for Success: Ranking Acid
and Base Strengths—The Relative Importance
of Effects on Charge
• To predict the relative stabilities of two species, the
following questions must be asked in the following order.
1. Do the species have different charges?
A charged species is usually more reactive than one that bears no
formal charge.
2. Do the charges appear on different atoms?
The size, electronegativity, and hybridization of the atom on which
the charge is located affects stability.
6.7 Strategies for Success: Ranking Acid and Base
Strengths—The Relative Importance of Effects on Charge
continued…
3. Are the charges delocalized differently through resonance?
All else being equal, stability increases as the number of atoms over
which a charge is shared increases.
4. Are there differences in inductive effects?
Electron-withdrawing groups stabilize nearby negative charges but
destabilize nearby positive charges. Electron-donating groups
stabilize nearby positive charges but destabilize nearby negative
charges.
Acid Strength
• Because pKa is defined in terms of each acid’s reaction with water, we
begin by writing out their reactions.
Free Energy Diagrams for
the Proton Transfer Reactions
• The acid in Equation 6-24c
(red) is charged, whereas
those in Equation 6-24a, 624b, and 6-24d are
uncharged.
• In Figure 6-13, therefore,
the reactants in Equation 624a, 6-24b, and 6-24d are
lower in energy than the
reactants in Equation 6-24c.
• Because the acids in
Equation 6-24a, 6-24b, and
6-24d are all uncharged, the
remaining three tiebreaking questions cannot
be used to distinguish their
relative stabilities; their
stabilities are assumed to
be similar.
Free Energy Diagrams for
the Proton Transfer Reactions
continued…
Free Energy Diagrams for
the Proton Transfer Reactions
continued…
• Moving to the product side, notice first that H3O+ can be
ignored in each case, so the tie-breaking questions
should be applied just to the conjugate bases.
• Acid strength therefore increases in the order:
A<B<C<D
Conjugate Base Strengths
• If you know the order of acid strengths, then it is
straightforward to determine the order of base strengths.
6.8 Strategies for Success: Determining
Relative Stabilities of Resonance Structures
• Although the resonance hybrid is
an average of all the resonance
contributors, not all resonance
contributors are weighted
equally.
• The resonance hybrid looks most
like the lowest energy (most
stable) resonance structure.
• The second resonance structure
on the right is less stable for
having formal charges. Therefore,
that structure makes a smaller
contribution to the resonance
hybrid.
Case #1 Charged verses Uncharged
Resonance Form
6.8 Strategies for Success: Determining
Relative Stabilities of Resonance Structures
continued…
• The main difference in the
following example is where the
formal charge is localized.
• Negative charge on the more
electronegative atom is more
stable.
• Therefore, the resonance
structure on the left having the
“˗” charge on the carbon atom
(more electropositive atom), is
less stable.
• Resonance structure on the left
contributes less to the resonance
hybrid.
Case #2 Negative charge on
different atoms.
6.8 Strategies for Success: Determining
Relative Stabilities of Resonance Structures
continued…
• The resonance structure which
possesses the “+” charge on
the more substituted C atom is
the more stable one.
• Again, the greater the degree
of alkyl substitution, the more
stable the carbocation.
Case #3 Positive charge on different
carbon atoms.
6.10 The Structure of Amino Acids
in Solution as a Function of pH
• Amino acids have both the amino
and the carboxyl group.
• This form never dominates in
aqueous solution.
• The carboxyl group is weakly
acidic and the amine group is
weakly basic.
• The form that the amino acid
takes depends upon the pH of the
solution.
The Zwitterion
• Under highly acidic (pH <
2) conditions, the weakly
basic N atom is
protonated. The resulting
species, which bears an
overall charge of +1, is
shown on the left in the
following equation.
• When the pH of the solution has risen
significantly above 2, the OH proton is
lost and the dominant form is the middle
species (this makes a zwitterion).
• When the pH of the solution is
significantly above 9 or 10 (the pKa of the
protonated amine), the anion on the right
dominates.
6.11 Electrophoresis
and Isoelectric Focusing
• Electrophoresis, one of
the most common ways
of separating a mixture of
amino acids or proteins.
• A mixture is spotted on a
gel or strip of paper that
has been buffered to a
specific pH.
• A high voltage (50–1,000
V) is applied by the
anode and cathode ends.
The Isoelectric Point (pI)
• Any ion that has a net positive charge at that pH will
migrate toward the negatively charged cathode.
• Any ion having a net negative charge will migrate toward
the positively charged anode.
• If the net charge is zero, the species will not move.
• An amino acid’s isoelectric pH, or isoelectric point (pI), is
the pH at which the substance has a charge of zero.
• The pI is unique for each of the 20 amino acids (pI = 6.07
for glycine; pI = 2.98 for aspartic acid; pI = 9.74 for
asparagine).
Summary and Conclusions
• Proton transfer reactions occurs when a Brønsted acid
donates an H+ to a Brønsted base.
• Curved arrow notation shows how the valence electrons
are involved in the proton transfer reaction.
• The equilibrium constant reflects the tendency to form
products.
• The acidity constant, Ka, reflects an acid’s tendency to
donate a proton. Electronegativity, charge, hybridization,
and resonance all affect a compound’s acid strength.
Summary and Conclusions
continued…
• A solvent’s leveling effect dictates the maximum strength
of an acid or a base that can exist in solution.
• The percent dissociation of an acid increases as the pH
of the solution increases.
• Functional groups mostly govern a compound’s pKa.
• If DGorxn is negative, the reaction tends to be
spontaneous. If DGorxn is positive, it tends to be
nonspontaneous.
Summary and Conclusions
continued…
• For two ions in which the formal charge is on an atom of
the same element, hybridization governs stability.
• Resonance effects can stabilize a charged species.
• Inductive effects can affect the stability of a charged
species by shifting electron density through covalent
bonds.