Lecture 3 Polar and non-polar covalent bonds Dr. A.K.M. Shafiqul Islam 21.07.08.
Chapter 2 Polar Covalent Bonds; Acids and Bases Part I Organic Chemistry.
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Transcript of Chapter 2 Polar Covalent Bonds; Acids and Bases Part I Organic Chemistry.
Chapter 2Polar Covalent Bonds; Acids and Bases
Part I
Organic Chemistry
Chapter Objectives
Take an in-depth look at polarity of molecules
Use formal charges to designate the distribution of electrons
Represent molecules with resonance structures by ‘pushing’ electrons
Examine the acid-base behavior of molecules
Predict acid-base reactions from pKa values
Electronegativity
electronegativity – a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound
The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond or a polar or nonpolar covalent bond.
Bond Formation
Ionic bonding involves the loss of an electron due to a large difference in electronegativity (EN>2.0)
Covalent bonding involves the sharing of electrons Equal sharing: non-polar bond (EN<.5) Unequal sharing: polar bond (.4<EN<2.1)
Polarity
If one side is more electronegative, it tends to have a partial negative charge (δ-) [electron-rich]
The other side tends to have a partial positive charge (δ+)[electron-poor]
The δ- and δ+ difference along a bond is called a dipole moment
δ-δ+
Electrostatic Potential Maps
Red – electron rich (δ-) Blue – electron poor (δ+)
Electrostatic Potential Maps
C CH
H H
HRed – electron rich (δ-) Blue – electron poor (δ+)
Electrostatic Potential Maps
Red – electron rich (δ-) Blue – electron poor (δ+)
Electrostatic Potential Maps
Red – electron rich (δ-) Blue – electron poor (δ+)
You describe it…What molecule do you
think it is? Take a guess…
Dipole Moments
2.1
2.1
2.1
2.1
2.5
3.52.5
3.5
acetic acid (ethanoic acid)
overall dipole moment = 1.70 D
Dipole Moments
acetic acid (ethanoic acid)
overall dipole moment = 1.70 D
Dipole Moment Calculations
Section 2.2 dipole moment (μ – Greek mu) – the
magnitude of the charge (Q) at either end of the molecular dipole times the distance (r) between the charges
measured in debyes (D) μ = Q x r Just be familiar with magnitude of values
Dipole Moment Values
Inductive Effect
2.1
2.1
2.1
2.12.5
3.52.5
3.5
acetic acid (ethanoic acid)
inductive effect – the shifting of electrons in a σ (sigma) bond in response to the electronegativity of nearby atoms.
Inductive Effect
Why would HCN allow the H+ to be released (proton donor – acid), thus categorizing HCN as an acid, when CH4 is not usually categorized as an acid?
You Try It.
Draw the complete Lewis Structure for the alcohol, methanol (methyl alcohol). Show the direction of its dipole moment. (μ =1.70)
You Try It.
Determine if the following molecules are polar or non polar. Show any dipoles.
(a) (b) (c)O O
OH
You Try It.
Draw a Lewis Structure of each of the following molecules and predict whether each has a dipole moment. If you expect a dipole moment, draw it in the correct direction.
(a) C2HF (b) CCl4 (c) CH3CHO
Formal Charges (Section 2.3)
formal charges – these charges don’t imply the presence of actual ionic charges …instead they give insight into the distribution of electrons
calculating the formal charges of each atom in a molecule will help you determine the best, most favorable structure (lowest energy)
General Rules of Stability
Lewis structures that approximate the actual molecule most closely are those that have:
maximum number of covalent bonds minimum separation of unlike charges formal charges of zero are ideal placement of any negative charges on the
most electronegative atom (or any positive charge on the most electropositive atom) Ex. Oxygen would rather 1- then 1+
DMSO (dimethyl sulfoxide)
Formal Charges
formal charge is calculated in the following manner:
If it violates HONC 1234, then it will have a formal charge on it.
1FC= # of valence electrons - # of non-bonding electrons+ # of bonding electrons
2
Formal Charges
Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there
is one less electron
CH4 H3O+ NH3BH3
Formal Charges
Give the formal charges for any atom on each of the following compounds Recall, having an overall + charge means that there
is one less electron
CH3NO2 H2C=N=N O3
[H2CNH2]+ (draw all resonance structures)(1 very likely, 1 less likely, 1 very
unlikely)
Resonance Structures
formaldehyde
Resonance Structures
Some compounds are not adequately represented by a single Lewis structure as we saw in the previous example.
When two or more structures are possible, the molecule will show characteristics of each structure.
Resonance Structures
Draw resonance structures for NO3-
The “real” structure is a resonance hybrid Each oxygen has a partial negative charge
N
O
OO
_ _
N
O
OO
_
N
O
OO
Resonance Structures
The “real” structure is said to have its electrons delocalized and is represented by a dotted bond
Resonance Structures
In some cases, one resonance form is more stable than another
(one accommodates formal charges better)
Resonance Structures
When drawing resonance structures, follow these rules:
1. Individual resonance forms are imaginary, not real2. Resonance forms differ ONLY in the placement of
their pi or non-bonding electrons3. Different resonance forms of a substance don’t
have to be equivalent4. All resonance forms must be valid Lewis
structures and obey normal rules for valency5. The resonance hybrid is more stable than any
individual resonance form
Resonance Structures
When drawing resonance structures, follow these rules:
1. 2. Resonance forms differ ONLY in the placement of
their pi or non-bonding electrons
Resonance Structures
When drawing resonance structures, follow these rules:
1. 2.
3. Different resonance forms of a substance don’t have to be equivalent
Resonance Structures
When drawing resonance structures, follow these rules:
1. Individual resonance forms are imaginary, not real2. Resonance forms differ ONLY in the placement of
their pi or non-bonding electrons3. Different resonance forms of a substance don’t
have to be equivalent4. All resonance forms must be valid Lewis
structures and obey normal rules for valency5. The resonance hybrid is more stable than any
individual resonance form
General Trends
+ C
- C
+ N/O
- N/O
Radicals
radical - (free radical) a neutral substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot (·) leaving it with an odd number of electrons.
Radicals are highly reactive! (octet rule) Radicals can form from stable molecules and
can also react with each other.
Radical Resonance
Resonance forms for radicals will depend upon three-atom groupings that contain a multiple bond next to a p-orbital.
Pentadienyl Radical
Pentadienyl Radical
Pentadienyl Radical
You try it.
Show all the resonance forms for the straight chained C7H9
. radical (in line angle).
.
Chapter 2Polar Covalent Bonds; Acids and Bases
Part II
Organic Chemistry
Define and describe acids and bases based on the Brønsted-Lowry and Lewis definitionsUse the curved-arrow formalism to show movement of electrons between Lewis acids and basesDetermine conjugate acid-base pairsPredict strength of acids and bases based on size, electronegativity, resonance stabilization, hybridization, and induction Predict reactions using pKa values
Section 2.7-2.11 Objectives
Why Study Acids/Bases?
At a deeper level, acid/base strength allows us to predict reactivity Compounds tend to react in such a way that
they become more stable (in the long run) Compounds considered “strong” are called that
(technically) because they dissociate completely, but (practically) also because they tend to react quickly. (This occurs because of LOW stability.)
Compounds considered “weak” tend not to react quickly or completely because they are stable (“happy” where they are )
Everything wants to be at the
lowest possible energy.
(most stable)
Stability
Acids & Bases
Definitions of acids and bases: Brønsted-Lowry definition
Acids donate protons (H+) (Proton donor) Hint for recognizing the acid – look for Hs!
Bases accept protons (Proton acceptor) Lewis definition
Acids accept electrons (electrophile)Bases donate electrons (nucleophile)
Hint for recognizing the base – look for electrons!Either a lone pair or pi bonded electrons
Morphine
Acid Reactions
So, if something loses a hydrogen, it has acted as an acid. It then has the capability to accept a proton. Therefore, what is it called at this point?
A BASE! Acids will donate a proton to become a
conjugate base. Bases will accept a proton to become a
conjugate acid.
General Acid-Base Reaction
When writing reactions, we show the movement (called an “attack”) of electrons with an arrow. Full headed arrow – both electrons Half headed arrow – one electron
Acid-Base Reactions
Brønsted-Lowry Theory
What is the acid, base, conjugate acid, and conjugate base?
acid base conjugate conjugate
acid base
Brønsted-Lowry Theory
base acid conjugate conjugate
acid base
What is the acid, base, conjugate acid, and conjugate base?
Dual Personality
amphoteric – a substance, that depending on the circumstances, can act like an acid or a base (like water!)
What makes hydrogen sulfate ion amphoteric?
- +4 2 4
- + -24 4
HSO + H H SO
HSO H + SO
Conjugate Acid-Base Pairs
+ -2 3HCl(aq) + H O(l) H O (aq) + Cl (aq)
acid conjugate acid
base conjugate base
+ -3 2 4NH (aq) + H O(l) NH (aq) + OH (aq)
acidbase conjugate acid
conjugate base
Conjugate Acid-Base Pairs
The stronger the acid, the weaker the conjugate base.
The weaker the acid, the stronger the conjugate base.
The stronger the base, the weaker the conjugate acid.
The weaker the base, the stronger the conjugate acid.
You Try It
Write the acid-base reaction between
CH3CH2OH and NaNH2
Write the acid-base reaction between
CH3COOH and NaOCH3
Write the acid-base reaction between CH3CH2OH and HCl
You Try It
What is the conjugate base of the following acids?
1. CH3COOH
2. CH3CH2NH3+
3.
Messing With Stability
If I take something that’s stable and change it by taking something away from it, what happens? It becomes unstable. Is this good or bad?
Ex. CH3OH a weak reagent Pretty stable (How do I know this?) If I remove an “H” – CH3O-
Not stable at all a strong reagent
Acid-Base Strength
Up to this point, the terms we’ve been using to describe acid-base strength have been very relative
Actual numerical values exist Recall the Ka value
The Equilibrium Expression(Law of Mass Action)
x y
eq n m
nA + mB xC + yD
[C] [D] K =
[A] [B]
The relationship between the concentration of products and reactants at equilibrium can be expressed by K
Acid Dissociation Constant
What if it is a reaction for the dissociation of an acid?
aK - +2 3
+ -3
a
HA + H O A + H O
[H O ][A ]K =
[HA]
What does the size of Ka mean?
High Ka = strong acid
Low Ka = weak acidacid dissociation constant
Acid Strength
IN CHEMICAL REACTIONS, THE ARROW USUALLY FAVORS THE PRODUCTION OF A WEAKER ACID AND BASE!!!
Why? What favors a weak acid over a strong one?
Weak acids and bases are more STABLE. If they weren’t stable, they would react to become stable…that’s why they are weak!
2 3 a
3 2 3 3 a
K
K
HCl H O H O Cl HIGH
CH OH H O H O CH O LOW
Strong acid
Weak acid Strong base
Weak base
What kind of values do you expect?
Ka vs. pKa
Acids with a greater Ka value are stronger than acids with a smaller Ka value
Problem with Ka relatively inconvenient because Ka values are usually on a negative power of ten Example: 1.0 x 10-4
To make things easier, the value pKa is used:
loga apK K
Calculating pKa
Determine the pKa of Hydrofluoric acid: Ka = 3.5 x 10-4
Phosphoric Acid: Ka = 7.5 x 10-3
~Which of these acids is stronger? H3PO4
pKa of HF:
pKa of H3PO4: What do you notice about pKa value compared to
acid strength?
3.52.1
The smaller the pKa, the stronger the acid.
pKa and Acid Strength
The smaller the pKa, the stronger the acid.
Reactivity
Do all acids react with all bases?
NO!!!!!!!!!!!!
How do we know when an acid will react with a particular base?
pKa values
You Try It
You Try It
Will the following reaction occur?
pKa = 49 pKa = 16
- -4 3 3 3CH + CH O CH + CH OH
You Try It
Write the products of the reaction and determine if it will occur.
You Try It
Predicting Acid/Base Strength
Use the pKa values if they are handy. Otherwise… 5 major factors exist which affect acid strength
Electronegativity Size Resonance stabilization (delocalization) Hybridization Induction
Predicting Acid/Base Strength
Electronegativity
Electronegativity
Which will give up a hydrogen ion (proton) more readily?CH4, NH3, H2O, HF
HF is most electronegative therefore the HF bond is shared unequally and easier to break
THE MORE ELECTRONEGATIVE THE CONJUGATE BASE, THE STRONGER THE ACID
Size
Size
Which is most reactive?HF, HCl, HBr, HI Recall:
If the negative charge is spread out more, it is a more stable conjugate base…therefore…
Resonance Stabilization
Again…if the charge on the conjugate base is spread out more than it is a more stable conjugate base. Therefore, the original acid is a stronger acid.
…so, how does resonance help this?
The negative charge of a conjugate base may be delocalized over several atoms thus making it more stable.
Resonance Stabilization
The negative charge of a conjugate base may be delocalized over several atoms thus making it more stable.
You Try It Which is the strongest acid?CH3CH2OH, CH3COOH, CH3SO3H
Hybridization
H3C-CH3 < H2C=CH2 < HC≡CH
worst acid best acid
sp3 sp2 sp
25%-s 33%-s 50%-s The more percent “s” in character, the closer the
electrons are to the nucleus, therefore the more polarized the structure…the H becomes more positive due to the pull of e- towards the C – makes a better acid
Hybridization
Inductive Effects – e- Withdrawing
Electronegative elements “take away” electron density from a negative charge:
Stabilityincreases
Inductive Effects – e- Withdrawing
Inductive Effects – e- Donating
hyperconjugation - Donation of a pair of bonding electrons into an unfilled or partially filled orbital
Inductive Effects – e- Donating
Which is the most stable conjugate base?
O– O – O
–
somewhat destabilizing very destabilizing!
Lewis Acids & Bases
Acid/base reactions can take place with or without a proton
Lewis bases are species that are able to donate a pair of electrons - Called nucleophiles (lover of nuclei)
Lewis acids are species that can accept this same pair of electrons – Called electrophiles (lover of electrons)
Drawn using curved arrow formalism (movement of electrons represented with arrows)
Lewis Acids & Bases
Strong nucleophiles are usually very strong Brønsted-Lowry bases (HIGH pKa) (Unstable)
Section 2.13 Noncovalent Interactions
Intermolecular Forces of Attraction dipole-dipole interactions (polar molecules) hydrogen bonding (polar molecules with F, O, or N
bonded to a H) London forces or dispersion forces (All molecules
have this but it is the only force present in nonpolar molecules)
hydrophilic – water loving (attracted to water) hydrophobic – water fearing (not attracted to water)