Transcript Slide 1

Chapter 2
Polar Covalent Bonds: Acid and
Bases
Chapter 2 - Definitions
• Polar Covalent Bonds – electrons are even distributed
between two atoms in a molecule.
• Electronegativity- the attractiveness of an atom to an
electron in a bond.
• Dipole Moment – is the total net molecular polarity.
• Formal Charge – assigns specific charges to individuals
atoms inside molecule, particularly to atoms that have an
apparently “abnormal” number of bonds.
• Resonance Forms – describes the movement of
electrons that accounts for the electron densities of
molecules.
• Bronsted-Lowery acid – is an proton donator.
Chapter 2 - Definitions
• Bronsted-Lowery base – is a proton acceptor.
• Acidity constant, Ka – determines the strength of an
acid.
• Lewis acid - is a substance that accepts an electron
pair.
• Lewis base – are a substance that donates an electron
pair.
What is Electronegativity?
• Electronegativity is the desire of an atom to gain an
additional electron to fill its octet or the strength of
the atom to pull electrons.
• The higher number the stronger the
electronegativity.
– Largest electronegativity: F = 4.0, Cl = 3.5,
O = 3.5, N = 3.0, Br = 2.8, C = 2.5
Why is Electronegativity So
Important?
• The electronegativity of two atoms in a bond
determines what type of bond forms.
• There are basically 2 types of bonds with a 3rd
also considered a type of bond.
– Ionic
– Covalent
– Polar Covalent (between a Ionic and covalent
bond.)
Similar Electronegativities
1
2
H
3
Li
11
Na
19
K
37
Rb
55
Cs
87
Fr
He
13
12
Al
Mg
Ca
38
Sr
56
Ba
88
Ra
C
B
Be
20
6
5
4
21
Sc
39
Y
57
La
89
Ac
22
Ti
40
Zr
72
Hf
104
Rf
23
V
41
Nb
73
Ta
105
Db
24
Cr
25
Mn
42
43
Mo
Tc
74
75
W
106
Sg
Re
107
Bh
26
Fe
44
Ru
76
Os
108
Hs
27
28
29
Co
Ni
Cu
45
46
Rh
Pd
47
77
78
Ir
109
Mt
Pt
Ag
79
Au
30
Zn
48
Cd
80
Hg
31
Ga
49
In
81
Tl
7
N
15
14
Si
32
Ge
50
Sn
82
Pb
110
Ds
Same color denotes similar electronegativities and a covalent bond.
P
33
As
51
Sb
83
Bi
8
9
O
16
S
34
F
17
Cl
35
10
Ne
18
Ar
Br
36
52
53
Te
I
54
84
85
86
Se
Po
At
Kr
Xe
Rn
Some Examples
• Na = 0.9, Cl = 3.0 (Ionic)
• C = 2.5, C = 2.5 (Covalent)
• C = 2.5, O = 3.5 (Polar Covalent)
What is the Difference Between
Bonding?
• Imagine that atoms play tough of war with the
electrons that they share in a bond.
• There would be three possibilities.
– 1) One atom wins and takes the electron
(ionic bond)
– 2) Both atoms are even matched sharing the
electrons evenly. (covalent)
– 3) One atom is stronger than the other atoms
and has the electron over its side more than
the other atom. (polar covalent)
Ionic Bonds
• Ionic Bonds are where an
electron is donated to
another atom. This
creates two charged
species. These charged
atoms or molecules are
normally free in solution
but are held together
when solid by
electrostatic attractions.
gains an electron
Na + Cl
Na+ + Cl-
gives up an electron
Solution
+
Cl-
Na
-
Cl
+
Na
Solid -crystals
Na+
Cl-
Na+Cl-
+
Na+Cl- Na Cl-
Covalent Bond
• Describes the sharing of electrons between
two atoms. There are two different types of
covalent bonds.
– Nonpolar covalent bonds is defined as the even
distribution of electrons between 2 atoms.
– Polar covalent bonds is defined as the uneven
distribution of electrons between two atoms.
Nonpolar Covalent Bonds
• Are bonds that are formed
between atoms with similar
electronegativities.
• Example chains of carbon (C)
bonded to hydrogens (H)
(Hydrocarbons)
H
H
C
CH3
H 3C
C
H
H
Polar Covalent Bonds
• Describes the unequal sharing of electrons in a
covalent bond.
Electrons are around the oxygen 70%
of the time
R
O
C
R
Electrons are around the carbon 30%
of the time
• This makes the oxygen considered partially negative
because the electrons around it more. The carbon is
considered partially positive because the shared
electrons are mainly around the oxygen.
Polar Covalent Bonds and Polar
Molecules
• Polar covalent bonds can be found in both
individual bonds and in entire molecules.
– To calculate individual bond polarity you need
to use the electronegativities of the two
atoms.
– To calculate the dipole moment of a molecule
you need to determine the center of positive
and negative charges. If they are not the
same then there is an overall polarity of the
molecule (called a polar molecule).
Dipole Moment
– The dipole moment of a molecule describes the
region of the molecules where the electron
density is highest and lowest.
– The dipole moment maintains a vector from low
electron density to high electron density.
– If the electron density is equal across the
molecules then the molecules is nonpolar
molecule (evenly distributed)
Non-Polar Molecules
• Are molecules which either maintain
atoms with similar electronegativities or
molecules whose dipoles are even in all
directions.
CH4 (nonpolar)
CCl4 (Nonpolar)
CH2Cl2 (Polar)
Non-Polar Molecules
• These molecules are
molecules whose
dipoles are even
spread in all directions
or maintain similar
H
electronegativities.
H
Cl
C
C
H
H
Methane
Cl
Cl
Cl
Carbon
Tetrachloride
Questions
Cl
H
Cl
C
C
C
Cl
H
H
H
H
Cl
H
Dichloromethane
Methane
Pick the polar molecule(s). And its(their) vectors.
Cl
Cl
Carbon
Tetrachloride
Polar Covalent Molecules
• The electronegativity
of the oxygen and
nitrogen atoms are
different there by
causing a overall
dipole in the molecule
making the molecules H
polar.
**
O
**
Water
**
N
H
H
H
H
Ammonia
Properties of Covalent Molecules
• Remember (LIKE DISSOLVES LIKE)
• Polar Covalent Molecules – Water,
Methanol, Ethanol.
• Nonpolar Covalent Molecules – (oils,
hydrocarbons) Propane, decane, etc..
• Oils and water do not mix because one is
polar covalent and the other is nonpolar
covalent.
Formal Charges
• Formal charges - assigns specific charges to
individuals atoms inside molecule, particularly to
atoms that have an apparently “abnormal”
number of bonds.
• This is used when you see charge separation in
a molecules to indicate if the atom is positive or
negative.
H
H
O
S
H
H
H
H
Formal Charge
• Formal charge = the charge on each atom in a molecule.
Formal
charge
Formal
Charge
# of valence e(free atom)
# of valence e-
# of valence e(bound atom)
Half of
bonding e-
# of
nonbonding e-
Resonance
• Resonance is the movement of either free
electrons or p electrons to form other
possible structures.
• Because movement of electron occurs
frequently resonance structures try to
show how the electrons might move.
• Only the movement of double (p bonds) or
free electrons are found. Movement of the
sigma bonds would breakup the molecule
instead of create resonance structures.
Examples of Resonance
H
O
C C
H
O
H
H
O
C C
H
O
H
H
O
C C
H
O
H
1
2
2
1
O
H
H
H
H
O
Base
O
H
H
1
H
O
H
H
H
H
2
Rules of Resonance
• 1) Individual resonance forms are imaginary, not
real.
• 2) Resonance forms differ only in the placement
of their p or nonbonding electrons.
• 3) Different resonance forms of a substance
don’t have to be equivalent.
• 4) Resonance forms obey normal rules of
valency.
• 5) The resonance hybrid is more stable than any
individual resonance form.
Resonance Forms
H
O
C C
H
O
H
H
O
C C
H
O
H
1
H
O
C C
H
O
H
2
1) Number 1 and 2 are resonance forms
2) Although resonance forms many not be equivalent, 1 and 2 are
equal.
Question
H
O
C C
H
O
H
Why does this oxygen have a negative
charge?
This oxygen has a negative charge because it has
taken an electron from a hydrogen giving it a
negative electron and a charge of -1. This
completes its octet with 8 electrons
Bronsted-Lowry Acid Base
• Bronsted-Lowry Acid - is a substance that
donates a proton (H+). It keeps the electron and
becomes negative.
• Bronsted-Lowry Base – is a substance that
accepts a proton (H+). It gives its proton so it
becomes positive.
Examples of Bronsted Lowry Acids
and Bases
H Cl
+
1
Acid
H
O
H
Base
O
Cl+
H
H
H
Conjugate Base
Conjugate Acid
O
O
2
C
H3C
H +
O
Base
Acid
3
H
O
H
Acid
O H
+
H
N
H
H
Base
C
H3C
O
Conjugate Base
+
H
O
H
Conjugate Acid
H
N
O H
H
H +
H
Conjugate Base
Conjugate Acid
Conjugate Acid and Bases
H Cl
Acid
+
H
O
H
Base
O
Cl+
H
H
H
Conjugate Base
Conjugate Acid
• Conjugate Acids and bases are if you look at a reaction as a
reversibility, these product would be the acid base found on the
opposite side.
• If you look at the reaction above the acid an base are clearly
defined, however if you switch the reaction then these would be the
acid and base, as shown below.
O
+
Cl
H
H
H
Conjugate Base
Conjugate Acid
H Cl
Acid
+
H
O
H
Base
Acid Base Strength
• Acids differ in their ability to donate an (H+)
proton.
• Some acids break apart and donate their proton
well (~100%), while others acids only gives an
proton about 50 percent or less of the time.
• The ability to give protons donates the strength
of the acid.
How do you Calculate the Strength
of an Acid?
• The way that you normally determine the strength of an
acid is to use the equilibrium constant. In dilute solutions
this is rewritten as the bottom equation.
[A-] [H3O+]
Keq =
[HA] [H2O]
[in (moles/liters) M]
In Dilute Solutions
Ka = Keq [H2O] =
[A-] [H3O+]
[HA]
What is the Ka Range of Acids?
• The strongest acids has a range from 1015 and
weaker acids are about 10-60.
– This is a wide difference so the use of pKa is
used.
– P = -Log (Number)
– So the pKa of an acid whose Ka = 1015 equals
• -Log(1015)
• -(15)
• -15 is the pKa of a strong acid
Some Examples of pKa
Weakest
Strongest
•
•
•
•
•
•
•
CH3CH2OH
H2O
HCN
H2PO4CH3CO2H
HNO3
HCl
16.00
15.74
9.34
7.21
4.76
-1.3
-7.0
What is the pKb?
• Just the same way that the acid strength
can be determined the basic strength can
also be determined. This strength is
opposite of what the acidic strength would
be.
Strength of pKb
Strongest
Weakest
•
•
•
•
•
•
•
CH3CH2OHOCNHPO42CH3CO2NO3Cl-
Predicting Acid Base Reactions
• The predict if an acid base reaction will occur
you have to determine if the stronger acid and/or
stronger base are on the left side. If this is not
true then the reaction will not proceed.
O
O
C
H3C
O
Stronger Acid
H
+
O H
Stronger Base
O
C
H3C
O
Weaker Base
C
H
+
H3C
O
Weaker Base
O
H
Weaker Acid
O
+
O
H
H
Weaker Acid
C
H +
H3C
O
Stronger Acid
O H
Stronger Base
Organic Acids
• Rule of thumb: Organic Acids often are found
where oxygen's are. The more oxygen's the
stronger the acid.
• Some examples are the: alcohols and carboxylic
acids
H
H
H
O
Alcohol
H
H3C
O
C
O
Carboxylic Acid
H
Which is the stronger acid?
H
H
H
O
H
Alcohol
H3C
O
C
O
Carboxylic Acid
H
Organic Bases
• Rule of Thumb: Organic Bases can maintain both
an oxygen or a nitrogen. Nitrogen almost always
functions as a base however oxygen can
function as both an acid or base.
• When something has a plus charge it is a base.
H2N
Base
H3N
H
O
Base
H
H
O
H
H
Which Oxygen Functions as a
Base?
H
O
H
O
H
H
H
H
O
C
O
H
H
O
C
O
Lewis Acids and Bases
• Lewis Acids – is a substance that accepts an
electron pair.
• Lewis Bases – are a substance that donates an
electron pair.
• These are much broader and can often be used
in both organic and inorganic chemistry.
Lewis Acid
• Maintains an empty or vacant orbital. If you
think of H+ this is a Lewis acid because it has
given up its single electron to another atom and
has only a proton.
• Another example is Mg2+.
Lewis Base
• Maintains a filled orbital. If you think of Nitrogen
containing molecules they maintain 1 filled lone
pair of electrons to share making them a Lewis
base. Any atom with a filled lone pair of
electrons to share. Oxygen, Nitrogen, Sulfur.
• Examples are amines and sulfides a sulfur
containing molecule.
Which of these are considered
Lewis Bases?
H
**
H ** C H
H
** *
*
**
H ** N H
H
** *
*
Ammonia
Methane
** *
*
*
*
H ** O H
Water
Non-Covalent Interactions
• Just like ionic atoms and molecules pair up due
to charge, dipole interaction makes the
intermolecular forces occur.
• Intermolecular forces use the partial positive and
negative charges cause by dipoles to pair.
These pair partial positive to partial negative.
Types of Non-Covalent Interactions
• Hydrogen Bonding – the strongest of these
forces shows the attractive forces of a Hydrogen
bond atom to an electronegative atom of O and
N.
• Vander Walls Forces – weaker interactions of
non hydrogen bound atoms. For example Cl, Br,
etc…
• Dispersion Forces – forces other molecules
away because the electron distribution is
constantly changing in a non-uniform fashion.
Things to Know
•
•
•
•
•
•
•
•
•
Electronegativity
Bonding – Ionic, Covalent (nonpolar, polar)
Polar and Non-Polar Molecules
Resonance
Acids and Bases (Bronsted-Lowry)
Acid Strength (Ka, pKa)
Organic Acids and Bases
Know examples of Lewis Acids and Bases
Non-Covalent Interactions (#, how do they work)