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Acid-Base Chapter 6
Bonding of Hydrogen
The chemistry of hydrogen depends on three electronic processes:
1.
2.
3.
The loss of an electron to give H+.
Acquisition of and electron to give H-.
The formation of a single covalent bond.
BUT, hydrogen has unique bonding features not described
well by the above generalizations.
These are the result of NO shielding of the nuclear charge
and allow for unique chemical activity.
1.
2.
3.
The hydrogen bond.
Hydrogen can bridge electron deficient compounds and
transition metal complexes.
Metal hydrides. It is possible to form numerous compounds with
metals.
The H-bond.
When hydrogen is bonded to another atom (F, O, N, Cl) such that the
bond is polar the H atom bears a partial positive charge and the more
electronegative atom bears a
negative charge.
This results in an interaction
between molecules
Details of these interactions
are disputed
but they are
generally accepted to be the
result of electrostatic forces.
In the case X-H--Y, the bond
between the X-H atoms is
found to elongate slightly.
BUT it remains a 2e- bond.
Very strong H-bonds result in
almost equal X-H and H-Y
spacings.
DNA and H-Bonding
Hydrogen-bonding pairs found in DNA. Note that
the A-T pair has only two H-bonds because of
the adenine molecule. (The lone pair of electrons
on the oxygen atom in thymine molecule could
act to form a third H-bond, but adenine does not
have the required donor). The arrangement of
donors and acceptors in thymine and guanine
precludes formation of H-bonding pairs between
these two compounds.
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Acid-Base Chapter 6
Agostic Interactions
This is a similar phenomenon to H-bonding.
This interaction occurs between hydrogen atoms bound to carbon
atoms of ligands and the transition metals to which they are bound.
This is an intramolecular bond.
(The word agostic is derived from the Greek word for "to hold on to oneself". )
The bonding arrangement is summarized as C-H--M
It is manifested by a significant lengthening of the C-H bond
Hydrides.
Hydrides are binary compounds of hydrogen and another element.
Although all hydrogen compounds could be termed as hydrides,
not all hydrogen containing compounds display hydridic
character.
Hydridic compounds are those that:
1.
2.
React as H- donors or,
Clearly contain anionic hydrogen.
Hence NaH is hydridic, methane and HCl are not.
This is parallel with the earlier statements…
1.
2.
3.
The loss of an electron to give H+.
Acquisition of and electron to give H-.
The formation of a single covalent bond.
Reactions of Ionic metal hydrides
1. All thermally decompose to give metal and
hydrogen.Only LiH is stable to its melting point
of 688oC.
Note that LiH is unreactive at moderate temperatures toward
oxygen and chlorine.
2. Generally ionic hydrides are highly reactive toward air
and water.
MH(s) + H2O
H2(g) + MOH(s)
MH2(s) + H2O
H2(g) + MOH2(s)
3. Ionic hydrides are powerful reducing agents and good
hydrogen-transfer agents.
NaH + B(OCH3)3
NaH + TiCl4
Na[HB(OCH3)3]
Ti0 +
4NaCl +2H2
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Acid-Base Chapter 6
Covalent hydrides
Covalent hydrides include:
Neutral binary XH4 compounds of Group 14, like methane.
Slightly basic binary XH3 compounds of Group 15, NH3 and PH3.
Weakly acidic or amphoteric, binary XH2 of Group 16, H2O and H2S.
Strongly acidic binary HX compounds of Group 17, HCl and HI.
Covalent hydrides of boron.
Hydridic, complex compounds of hydrogen.
Examples include LiAlH4 and NaBH4. These are powerful reducing agents
despite the covalent nature of the Al-H and B-H bonds.
1.
2.
3.
4.
5.
6.
Some interesting notes about LiAlH4 and NaBH4.
These two compounds are ionic in nature BUT they
possess tetrahedral anions containing covalent bonds to H
How are LiAlH4 and NaBH4 prepared?
Both these reactions are carried out in ether.
8LiH + Al2Cl6
2LiAlH4 + 6 LiCl
2 NaH + B2H6
2NaBH4
The anions are powerful hydrogen transfer agents.
2LiAlH4 + 2 SiCl4
I2 + 2 NaBH4
2SiH4 + 2 LiCl + Al2Cl6
B2H6 + 2NaI + H2
Some more details about covalent hydrides.
Covalent hydrides can be divided into three subcategories which rely
on the nature of the H atom.
1.
2.
3.
The H-atom is neutral.
The H-atom is positive.
The H-atom is negative.
These are a generalization of the previous statements.
By far the majority of covalent hydrides fall into the first category.
Given their low polarity these compounds are only held together by
weak intermolecular forces … termed dispersion forces.
This results in low boiling points.
SnH4 -52oC , PH3 -90oC
Carbon-based systems comprise the largest set of hydrides.
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Acid-Base Chapter 6
If the H-atom is positive.
This brings us to the case of H2O, HF, and NH3.
These molecules have VERY high
boiling points which arise as a result
of intermolecular interactions.
Water is an idea example of this.
The H can bridge two molecules
forming an intermolecular bond
with an adjacent molecule.
We can try to understand this using a molecular orbital picture for water.
The molecular orbital picture of water.
⇑⇓
⇑
⇑
2pz
2py
2px
⇑⇓
NB
⇑⇓
⇑⇓
⇑
x2
H
⇑⇓
2s
O
⇑⇓
Bringing two water molecules together.
NB
⇑⇓
⇑⇓
⇑⇓
⇑⇓
⇑⇓
⇑⇓
The total energy of
The system is LOWER
⇑⇓
⇑⇓
⇑⇓
Water 1
Water 2
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Acid-Base Chapter 6
H2 as a ligand
H2 can behave as a ligand and occupy a coordination space around a metal
center. This occurs for metals in low oxidation states.
Hydrogen can take on a “ side on” orientation with respect to the metal
resulting in weak donation of its electrons into an empty σ on the metal center
OR through acceptance of electrons from a filled orbital on the metal into the
σ* orbital of the H2 molecule. This weakens the H-H bond.
Unless the system is very carefully balanced a system will tend toward the
dihydride.
The bonding configuration of M-H2
Bonding in H2 complexes does not require any complicated explanation. In a
similar way to that seen in boranes a three-center two-electron bond is formed.
The H2 molecule thus acts a neutral two electron sigma donor. It is also
possible to understand the bonding as back-donation of electrons from a filled
metal orbital to the sigma-* orbital on the H2.
Both are shown schematically below.
Acids and Bases
A logical place to go from our study of
hydrogen compounds it to Acids and Bases.
In 1884 Arrhenius devised a theory to explain acid/base behavior.
His theory states that Acids contain protons
and bases contain hydroxide ions.
This neglects two issues that
result in flaws in the theory:
1.
2.
The solvent
The salt
This comes to light if we think about HCl in two different solvents.
It behaves very differently in Water and Benzene.
This implies the solvent has an affect on the chemistry.
Salts seem to contradict this rule.
Many including carbonates and phosphates are basic while,
ammonium and aluminum ions are acid.
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Acid-Base Chapter 6
In come Bronsted and Lowry
Bronsted and Lowry devise a theory that uses the concepts
that acids donate protons and bases accept protons.
In this theory the solvent becomes important
as it can self-ionize and have acid base
characteristics. AUTOIONIZATION
H2 O
H2 O
H3O+
OH-
Autoionization of water
H2 O
H2 O
H3O+
OH-
In this process water does two things; it accepts a proton and it
donates a proton. Hence, it behaves as both an acid and a base.
It is said to be amphoteric or amphiprotic.
This model relies on the existence of the H3O+ ion. The first crystal
structure of this of this ion was shown in 1924, one year after the
development of the theory. It was in the structure of perchloric
acid monohydrate.
The hydronium ion
H2 O
H2 O
+
H3O+
OH-
Although the hydronium ion is
typically shown as H3O+, there
are three waters of hydration Hbonded to it.
What effect does this have on its stability?
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Acid-Base Chapter 6
The importance of solvent.
Acid base behavior relies heavily on the solvent present.
The strongest acid in a given solvent is the protonated solvent.
This is illustrated well in the case of HF.
HF (aq) + H2O ⇔ H3O+(aq) + FHere water is acting as a base and the fluoride ion is the conjugate base of HF.
In a reaction with ammonia water is acting as an
acid to give the conjugate acid the ammonium ion.
H3O+(aq) + OH-(aq) ⇒ 2H2O
The reaction of a strong acid and strong base may be summarized as the
reaction between the hydronium ion and the hydroxide ion. This reaction is
viewed as going to completion because water dissociates to such a small
degree.
A note about aqueous solutions
In aqueous solution the strongest acid is the hydrodium
ion and the strongest base is the hydroxide ion.
SO WHAT!!
What does this mean? What impact does this have on reactions?
Generally, stronger acids and bases react upon addition to water
to produce either the hydronium ion or the hydroxide ion.
O2- + H2O ⇒ 2OH-(aq)
HClO4 + H2O ⇒ H3O+(aq) + ClO4-(aq)
are we as chemists, limited to water? NO!!
Solvents other than water..
Any solvent with an ionizable hydrogen
atom will work with a B-L acid-base system.
How about liquid ammonia? (rather cold : -33oC.)
NH3 (l) + NH3 (l) ⇔ NH4+(NH3) + NH2-(NH3)
Here is an example of an acid base reaction in ammonia.
NH4Cl(NH3) + NaNH2(NH3) ⇔ NaCl + 2NH3
Another COOL example is sodium metal
in ammonia. Sodium acts as a base and
the solution turns bright blue!
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Acid-Base Chapter 6
Equilibrium constants: Acids
We understand and quantify acid strength
in terms of how easy it is to transfer the H+.
We use water as a reference point.
The general dissociation process is outlined in the following reaction.
HA(aq) + H2O(l) ⇔ H3O+ + A-(aq)
The mathematical expression that describes this process is:
ka =
[H 3O+ ][A− ]
[HA]
To simplify things we use the value of pKa to describe the strength
of an acid. The more negative the value…the stronger the acid.
Equilibrium constants: Bases
We understand and quantify base strength
in terms of how easy it is to transfer the H+.
We use water as a reference point.
The general dissociation process is outlined in the following reaction.
A-(aq) + H2O(l) ⇔ OH- (aq) + HA(aq)
The mathematical expression that describes this process is:
kb =
[HA][OH − ]
[A− ]
To simplify things we use the value of pKb to describe the strength
of a base. Again the more negative the pKb the stronger the base.
Acid/Base Reactions: A Summary
Bronstead-Lowry
Lewis
Acids
Proton Donor
Electron Acceptor
Base
Proton Acceptor
Electron Donor
Neutralization reaction involving HCl and NaOH
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl-(aq) + H2O(l)
What is the NET ionic equation?
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Acid-Base Chapter 6
Relating Ka , Kb , and Kw
The autoionization of water can be related to the Ka and Kb of
acids and their conjugate bases in aqueous solutions.
kw =
[HA][OH − ] [H 3O+ ][A− ]
×
[A− ]
[HA]
= [H 3O][OH − ]
= 10−14
Applying the power scale again…
pkw = pk a + pk b
= 14
Trends in acid strength
Strong acids Ka> 1 (negative pKa values)
Examples of these acids are: hydrochloric, nitric, sulfuric, perchloric acids.
Weak acids Ka< 1 (positive pKa values)
Examples of these acids are: nitrous, hydrofluoric.
The majority of other inorganic acids are weak. That means there is
a significant amount of the molecular species existing in solution.
Strong Acids
Generally, all strong acids appear equally strong. This
means that nearly 100% of the species is ionized in
solution.
This underscores the importance of the solvent as a
“proton-vehicle” (remember the hydronium ion).
How would you evaluate the relative strength of acids?
Evaluating the strength of strong acids.
To qualitatively identify the strength of an acid we dissolve stronger
acids in a base weaker than water.
What does this tell you about the acidic strength of the “new” solvent
A good example is hydrofluoric acid.
If we dissolve perchloric acid in HF, what happens?
HClO4(HF) + HF(l) ⇔ H2F+ + ClO4-(HF)
As with water, the weaker acid (HF) acts as a proton acceptor (base)
for perchloric acid.
As a result of HF being a weaker base than water the equilibrium
doesn’t lie 100% “to the right”. We can do this for a series of acids
and evaluate their strength.
The strongest of common acids is Perchloric Acid.
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Acid-Base Chapter 6
Binary Acids
What is a binary acid?
Acids consisting of two simple ions (like hydrohalic acids)
HX(aq) + H2O(l) ⇔ H3O+ + X-(aq)
Acid
pKa
Bond Energy
(kJ mole-1)
565
HF
+3
HCl
-7
428
HBr
-9
362
HI
-10
295
Generally, the differences in strength can be related to the difference
between the strength of the H-X and O-H bonds. The O-H bond
energy is 459 kJ mole -1.
Any reaction tends toward the formation of a stronger bond.
How can this be related to acid strength?
Oxyacids
What is a oxyacid?
Ternary acids (contain more than one atomic species) that contain oxygen.
You MUST remember that for all common inorganic oxyacids the ionizable
hydrogen is bound to an oxygen.
Trends in strength.
If we study oxyacids of one element we see stronger acids when there are
more oxygen atoms…………….WHY?
Look at nitric and nitrous acids
pKa= -1.4 vs. pKa=3.3
How can YOU explain this?
Look at the structures.
Semiquantitative prediction of Oxyacid
strength.
We can predict pKa values from the formula (HO)nXOm
m
pKa
0
8
1
1
2
-1
3
-8
Note that ionizable H’s are attached to an oxygen!!
Can we understand the strength of
perchloric acid looking at its structure?
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Acid-Base Chapter 6
Polyprotic acids
What is a polyprotic acid?
Acids which have more than one ionizable hydrogen.
Note that it is increasingly difficult to remove successive
protons, hence the “strength” decreases.
Sulfuric acid is good example.
H2SO4 what’s its structure?
H2SO4 (aq) + H2O (l) ⇒ H3O+ + HSO4-
pKa= -2
HSO4- (aq) + H2O (l) ⇔ H3O+ + SO42-
pKa= 2
This brings an interesting point in regard to the crystallization
salts of the +2 and +3 ions with metal cations.
Sulfuric Acid
S(s)
O2
⇒
O2/ V2O5
SO2 ⇒
Contact
Process
H2 O
⇒
SO3
H2SO4
H2SO4
H2 O
H2S2O7
Fuming Sulfuric Acid
Sulfuric acid is the largest tonnage chemical.
It is NOT a strong oxidizing agent but does act
as a dehydrating agent very effectively.
Nitric Acid
Haber
Oswald
O2
(+4)
H2 O
(+5)
(+2)
N2(g) + 3H2 ⇒ 2NH3 ⇒ NO(g) ⇒ 3NO2 ⇒ 2HNO3 + NO
O2/Pt
Some useful notes:
Bottles of HNO3 have brown gas above the solution.
hv
Autoionization:
2HNO3 ⇒ 2NO2 + H2O + 1/2 O2
2HNO3 ⇔ H2NO3+ + NO3H2NO3+ ⇔ NO2 + H2O
This acid is a strong oxidizing agent because of the presence of the nitryl
cation NO2+. This cation is used in organic reactions for nitrating organics.
Aqua-Regia is a mixture of 3HCl/1HNO3 it contains free Cl2 and ClNO2. This
is a powerful oxidizing environment and will dissolve Au and Pt.
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Acid-Base Chapter 6
Acidic Metal Ions
Some metal ions are very acidic in aqueous solutions.
Examples are Aluminum and Iron. Why do these metals do this?
Both ions are relatively small and highly charged existing as
hexahydrates in water. [Al(OH2)6]3+ and [Fe(OH2)6]3+
[Fe(OH2)]3+(aq) + H2O(l) ⇔ H3O+(aq) + [Fe(OH2)5(OH)]2+(aq) pKa = 3.3
Be sure you can do it for Al3+!!
Mn+(aq) + H2O(l) ⇔ H3O+(aq) + M(n-1)+(aq)
pKa1 = Na+ (14) Mg2+ (12)
Al3+ (5)
neutral weakly acidic acidic
Fe2+ (20) Fe3+ (3)
Bronsted-Lowry Bases
The most important B-L base is hydroxide.
After this, it is ammonia. It reacts with water to give OH-.
NH3(aq) + H2O(l) ⇔ NH4+(aq) + OH-(aq)
Ammonia is useful as a glass cleaner as
it reacts with fat molecules to make water
soluble salts.
Other Bronsted-Lowry Bases
Other common bases are the conjugate bases of weak acids.
PO43-(aq) + H2O(l) ⇔ HPO42- (aq) + OH-(aq)
pKb= 1.35
S2-(aq) + H2O(l) ⇔ HS- (aq) + OH-(aq)
pKb= 2.04
F-(aq) + H2O(l) ⇔ HF (aq) + OH-(aq)
pKb= 10.35
Two of these are polyprotic acids…
what happens there?
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Acid-Base Chapter 6
Conjugate bases of polyprotic acids
HS-(aq) + H2O(l) ⇔ H2S (aq) + OH-(aq)
Can you predict the second step for the sulfide ion?
pKb = 6.96
Will its pKb be smaller, larger, or the same and why?
HPO42-(aq) + H2O(l) ⇔ H2PO4- (aq) + OH-(aq) pKb = 6.79
This is an interesting case as the hydrogen phosphate ion behaves as a BASE!
NOT AS AN ACID
H2PO4-(aq) + H2O(l) ⇔ H3PO4 (aq) + OH-(aq)
pKb = 11.88
What does the pKb of the third process tell you?
Conjugate bases of strong acids.
Generally, conjugate bases of strong
acids do not interact with water.
If in doubt think about a solution of
sodium chloride.
Lewis Acids and Bases.
Acid (LA) : electron pair acceptor
Base (LB): electron pair donor
Bronsted-Lowry acids and bases are a special case of this theory.
HA(aq) + H2O(l) ⇔ H3O+ + A-(aq)
Think about H+, OH- and NH3.
All metal ions are LA, and most ligands are LB.
Molecules with incomplete octets are LA.
B2H6 + 2N(CH3)3
2H3B⇐ N(CH3)3
Or how about BF3 and its reaction with NH3.
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Acid-Base Chapter 6
Lewis Acids and Bases.
Molecules or ions that expand their octet are LA.
SiF4 + 2FSi(sp3)
PF5 + FP(sp3)
⇒
⇒
SiF62- ; why not CF4?
Si(sp3d2)
PF6P(sp3d2)
Molecules or ions with complete octets that can alter their valence electrons
to accept electrons are LA.
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