Transcript 2. Polar Covalent Bonds: Acids and Bases

Organic Chemistry
M. R. Naimi-Jamal
Faculty of Chemistry
Iran University of Science & Technology
Chapter 1. Continue
Polar Covalent Bonds
Acids and Bases
Based on: McMurry’s Fundamental of Organic Chemistry, 4th
edition, Chapter 1
Polar Covalent Bonds: Electronegativity
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Covalent bonds can have ionic character
These are polar covalent bonds
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Bonding electrons attracted more strongly by
one atom than by the other
Electron distribution between atoms in not
symmetrical
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Bond Polarity and Electronegativity
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Electronegativity (EN): intrinsic ability of an
atom to attract the shared electrons in a covalent
bond
Differences in EN produce bond polarity
Arbitrary scale. As shown in next figure,
electronegativities are based on an arbitrary scale
F is most electronegative (EN = 4.0), Cs is least
(EN = 0.7)
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The Periodic Table and Electronegativity
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Bond Polarity and Electronegativity
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Metals on left side of periodic table attract
electrons weakly: lower electronegativities
Halogens and other reactive nonmetals on right
side of periodic table attract electrons strongly:
higher electronegativities
Electronegativity of C = 2.5
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Bond Polarity and Inductive Effect
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Nonpolar Covalent Bonds: atoms with similar
electronegativities
Polar Covalent Bonds: Difference in EN of atoms < 2
Ionic Bonds: Difference in electronegativities > 2
(approximately).
Other factors (solvation, lattice energy, etc) are
important in ionic character.
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Bond Polarity and Inductive Effect
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Bonding electrons are pulled toward the more
electronegative atom in the bond
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Electropositive atom acquires partial positive
charge, +
Electronegative atom acquires partial negative
charge, -
Inductive effect: shifting of electrons in a bond
in response to the electronegativities of nearby
atoms
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Electrostatic Potential Maps
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Electrostatic potential maps show calculated charge
distributions
Colors indicate electron-rich (red) and electron-poor
(blue) regions
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Polar Covalent Bonds: Dipole Moments
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Molecules as a whole are often polar, from vector
summation of individual bond polarities and lonepair contributions
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Polar Covalent Bonds: Dipole Moments
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Dipole moment - Net molecular polarity, due to
difference in summed charges
 - magnitude of charge Q at end of molecular
dipole times distance r between charges
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 = Q  r, in debyes (D)
1 D = 3.34  1030 coulomb meter
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Dipole Moments in Water and Ammonia
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Large dipole moments
 Electronegativities of O and N > H
 Both O and N have lone-pair electrons oriented
away from all nuclei
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Practice: Suggest an explanation
Ammonia (NH3) has a dipole moment of 1.46 D,
while the dipole moment of NF3 is only 0.24 D.
Why?
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Absence of Dipole Moments
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In symmetrical molecules, the dipole moments of
each bond has one in the opposite direction
The effects of the local dipoles cancel each other
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Cis- & trans-1,2-dichloroethylenes:
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Formal Charges
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Sometimes it is necessary to have structures with
formal charges on individual atoms
We compare the bonding of the atom in the
molecule to the valence electron structure
If the atom has one more electron in the molecule, it
is shown with a “-” charge
If the atom has one less electron, it is shown with a
“+” charge
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Formal Charges
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Nitromethane:
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Formal Charges
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Resonance
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Some molecules have structures that cannot be
shown with a single Lewis representation
In these cases we draw Lewis structures that
contribute to the final structure but which differ in
the position of the  bond(s) or lone pair(s)
Such a structure is delocalized and is represented by
resonance forms
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Resonance
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The resonance forms are connected by a doubleheaded arrow
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Resonance Hybrids
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A structure with resonance forms does not alternate
between the forms
Instead, it is a hybrid of the two resonance forms, so
the structure is called a resonance hybrid
For example, benzene (C6H6) has two resonance
forms with alternating double and single bonds
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In the resonance hybrid, the actual structure, all
of the C-C bonds are equivalent, midway between
double and single bonds
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Resonance Hybrids
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Resonance Hybrids
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Rules for Resonance Forms
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Individual resonance forms are imaginary - the
real structure is a hybrid (only by knowing the
contributors can you visualize the actual structure)
Resonance forms differ only in the placement of
their  or nonbonding electrons
Different resonance forms of a substance don’t
have to be equivalent
Resonance forms must be valid Lewis structures:
the octet rule usually applies
The resonance hybrid is more stable than any
individual resonance form would be
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Curved Arrows and Resonance Forms
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We can imagine that electrons move in pairs to
convert from one resonance form to another
A curved arrow shows that a pair of electrons
moves from the atom or bond at the tail of the
arrow to the atom or bond at the head of the
arrow
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Curved Arrows and Resonance Forms
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Drawing Resonance Forms
Any three-atom grouping with a multiple bond has
two resonance forms
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Different Atoms in Resonance Forms
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Sometimes resonance forms involve different
atom types as well as locations
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The resulting resonance hybrid has properties
associated with both types of contributors
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The types may contribute unequally
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Resonance in the acetone enolate
The “enolate” derived from acetone is a good illustration,
with delocalization between carbon and oxygen.
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2,4-Pentanedione
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The anion derived from 2,4-pentanedione
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Lone pair of electrons and a formal negative
charge on the central carbon atom, next to a
C=O bond on the left and on the right
Three resonance structures result
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Practice: Draw three resonance forms:
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Solution:
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Practice: Draw three resonance forms:
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Solution:
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Solution:
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Acids and Bases:
The Brønsted–Lowry Definition
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Brønsted–Lowry theory defines acids and bases
by their role in reactions that transfer protons
(H+) between donors and acceptors.
“proton” is a synonym for H+ - loss of an electron
from H leaving the bare nucleus - a proton.
Protons are always covalently bonded to another
atom.
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Brønsted Acids and Bases
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“Brønsted-Lowry” is usually shortened to “Brønsted”
A Brønsted acid is a substance that donates a
hydrogen ion, or “proton” (H+): a proton donor
A Brønsted base is a substance that accepts the H+:
a proton acceptor
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The Reaction of HCl with H2O
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When HCl gas dissolves in water, a Brønsted acid–
base reaction occurs
HCl donates a proton to water molecule, yielding
hydronium ion (H3O+) and Cl
The reverse is also a Brønsted acid–base reaction
of the conjugate acid and conjugate base
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The Reaction of HCl with H2O
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Quantitative Measures of Acid Strength
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The equilibrium constant (Keq) for the reaction of an
acid (HA) with water to form hydronium ion and the
conjugate base (A-) is a measure related to the
strength of the acid
Stronger acids have larger Keq
Note that brackets [ ] indicate concentration, moles
per liter, M.
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Ka – the Acidity Constant
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The concentration of water as a solvent does not change
significantly when it is protonated in dilute solution.
The acidity constant, Ka for HA equals Keq times 55.6 M
(leaving [water] out of the expression)
Ka ranges from 1015 for the strongest acids to very small
values (10-60) for the weakest
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Ka – the Acidity Constant
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Acid and Base Strength
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The ability of a Brønsted acid to donate a proton
to is sometimes referred to as the strength of the
acid.
The strength of the acid can only be measured
with respect to the Brønsted base that receives
the proton
Water is used as a common base for the purpose
of creating a scale of Brønsted acid strength
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pKa – the Acid Strength Scale
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pKa = -log Ka (in the same way that
pH = -log [H+]
The free energy in an equilibrium is related to
–log of Keq (DG = -RT log Keq)
A larger value of pKa indicates a stronger acid
and is proportional to the energy difference
between products and reactants
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pKa – the Acid Strength Scale
The pKa of water is 15.74
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pKa values for some acids:
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Predicting Acid–Base Reactions from
pKa Values
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pKa values are related as logarithms to equilibrium
constants
The difference in two pKa values is the log of the
ratio of equilibrium constants, and can be used to
calculate the extent of transfer
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Predicting Acid–Base Reactions
from pKa Values
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Predicting Acid–Base Reactions
from pKa Values
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Organic Acids and Organic Bases
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The reaction patterns of organic compounds often
are acid-base combinations
The transfer of a proton from a strong Brønsted
acid to a Brønsted base, for example, is a very
fast process and will always occur along with
other reactions
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Organic Acids
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Those that lose a proton from O–H, such as
methanol (CH3OH) and acetic acid (CH3COOH)
Those that lose a proton from C–H, usually
from a carbon atom next to a C=O double
bond (O=C– C– H)
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Organic Acids
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Organic Acids:
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Carboxylic Acids:
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Conjugate Bases:
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Organic Bases
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Have an atom with a lone pair of electrons that
can bond to H+
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Nitrogen-containing compounds derived from
ammonia are the most common organic bases
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Oxygen-containing compounds can react as bases
when with a strong acid or as acids with strong
bases
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Organic Bases
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Acids and Bases: The Lewis Definition
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Lewis acids are electron pair acceptors;
Lewis bases are electron pair donors
The Lewis definition leads to a general description of
many reaction patterns but there is no quantitatve
scale of strengths as in the Brønsted definition of pKa
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Lewis Acids and the Curved Arrow
Formalism
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The Lewis definition of acidity includes metal cations,
such as Mg2+
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They accept a pair of electrons when they form a
bond to a base
Group 3A elements, such as BF3 and AlCl3, are Lewis
acids because they have unfilled valence orbitals and
can accept electron pairs from Lewis bases
Transition-metal compounds, such as TiCl4, FeCl3,
ZnCl2, and SnCl4, are Lewis acids
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Lewis Acids and the Curved Arrow
Formalism
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Organic compounds that undergo addition
reactions with Lewis bases (discussed later) are
called electrophiles and therefore Lewis Acids
The combination of a Lewis acid and a Lewis base
can shown with a curved arrow from base to acid
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Illustration of Curved Arrows in Following
Lewis Acid-Base Reactions
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Some Lewis Acids:
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Lewis Bases
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Lewis bases can accept protons as well as other
Lewis acids, therefore the definition encompasses
that for Brønsted bases
Most oxygen- and nitrogen-containing organic
compounds are Lewis bases because they have
lone pairs of electrons
Some compounds can act as either acids or bases,
depending on the reaction
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Some Lewis Bases
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Lewis Acid-Base Reactions
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Imidazole:
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Drawing Chemical Structures
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Condensed structures: C-H and C-C and single bonds
aren't shown but understood
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If C has 3 H’s bonded to it, write CH3
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If C has 2 H’s bonded to it, write CH2; and so on.
Horizontal bonds between carbons aren't shown in
condensed structures - the CH3, CH2, and CH units
are assumed to be connected horizontally by single
bonds, but vertical bonds are added for clarity
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2-methylbutane Structures
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Skeletal Structures
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Minimum amount of visual information but
unambiguous
C’s not shown, but assumed to be at each
intersection of two lines (bonds) and at end of
each line
H’s bonded to C’s aren't shown – whatever
number is needed will be there to fill out the
four bonds to each carbon.
All atoms other than C and H are shown
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Practice: How many H’s on each carbon?
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Solution:
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Problem: How many H’s on each carbon?
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Molecular Models
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We often need to visualize the shape or
connections of a molecule in three dimensions
Molecular models are three dimensional objects,
on a human scale, that represent the aspects of
interest of the molecule’s structure (computer
models also are possible)
Drawings on paper and screens are limited in
what they can present to you
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Molecular Models
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Framework models (ball-and-stick)
are essential for seeing the
relationships within and between
molecules – you should own a set
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Space-filling models are better for
examining the crowding within a
molecule
Framework
Space-filling
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Summary
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Organic molecules often have polar covalent bonds
as a result of unsymmetrical electron sharing caused
by differences in the electronegativity of atoms
The polarity of a molecule is measured by its dipole
moment, .
(+) and () indicate formal charges on atoms in
molecules to keep track of valence electrons around
an atom
Some substances must be shown as a resonance
hybrid of two or more resonance forms that differ
by the location of electrons.
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Summary
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A Brønsted(–Lowry) acid donates a proton
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A Brønsted(–Lowry) base accepts a proton
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The strength Brønsted acid is related to the -1 times
the logarithm of the acidity constant, pKa. Weaker
acids have higher pKa’s
A Lewis acid has an empty orbital that can accept an
electron pair
A Lewis base can donate an unshared electron pair
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Summary
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In condensed structures C-C and C-H are implied
Skeletal structures show bonds and not C or H (C is
shown as a junction of two lines) – other atoms are
shown
Molecular models are useful for representing
structures for study
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Chapter 1-2, Questions
42, 43, 44, 46, 47,
50, 54, 55, 57
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