Transcript CHAPTER 10

Unit 6
Reactions in
Aqueous Solutions I:
Acids, Bases & Salts
Properties of Aqueous Solutions of Acids
and Bases
Aqueous acidic solutions have the following
properties:
1.
2.
They have a sour taste.
They change the colors of many indicators.
–
–
Acids turn blue litmus to red.
Acids turn bromothymol blue from blue to yellow.
Bromothymol blue is yellow in acidic
solution and blue in basic solution
3.
They react with metals to generate hydrogen, H2(g).
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Properties of Aqueous Solutions of Acids
and Bases
4. They react with metal oxides and hydroxides to
form salts and water.

HCl (aq) +
CaO(s)  CaCl2 (aq) + H2O (l)
5. They react with salts of weaker acids to form the
weaker acid and the salt of the stronger acid.

3HCl(aq) + Na3PO4 (aq) 
H3PO4 (aq) + 3NaCl (aq)
6. Acidic aqueous solutions conduct electricity.
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Properties of Aqueous Solutions of Acids
and Bases
Aqueous basic solutions have the following
properties:
1. They have a bitter taste.
2. They have a slippery feeling.
3. They change the colors of many indicators
–
–
Bases turn red litmus to blue.
Bases turn bromothymol blue from yellow to blue.
4. They react with acids to form salts and water.
5. Aqueous basic solutions conduct electricity.
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The Arrhenius Theory
• Acids are substances that contain hydrogen and
produce H+ in aqueous solutions.
• Two examples of substances that behave as
Arrhenius acids:
HCl (aq)  H 2 O (  )  H 3O aq   Cl-aq 
 H O   HCO 
HCO 2 H (aq)  H 2 O (  ) 
3 aq 
2(aq)
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The Arrhenius Theory
• Bases are substances that contain the
hydroxyl group (–OH) and produce
hydroxide ions (OH-) in aqueous solutions.
• Two examples of substances that behave
as Arrhenius bases:

aq 
NaOH  Na  OHaq 
2


Ca(OH)
Ca  2 OH
2
aq 
-
aq 
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The Arrhenius Theory
• Although Arrhenius described H+ ions in water as
bare ions (protons) they really exist as hydronium
ions, H3O+
• It is this hydronium ion that gives aqueous
solutions of acids the characteristic acidic
properties
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The BrØnsted-Lowry Theory
• An acid is a proton donor (H+).
• A base is a proton acceptor.
• Two examples to illustrate this concept:
HBr  H 2 O  H 3O   Br acid
NH3
base
base


 H 2 O  NH4  OH
acid
• In the Arrhenius definition, NH3 would not be classified as a
base.
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The BrØnsted-Lowry Theory
• Acid-base reactions are the transfer of a proton
from an acid to a base.

Note that coordinate covalent bonds are often
made in these acid-base reactions.
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The BrØnsted-Lowry Theory
•
An important part of BrØnsted-Lowry acid-base
theory is the idea of conjugate acid-base pairs.
– Two species that differ by a proton are called acid-base
conjugate pairs.
•
•
Conjugate Base is what the acid becomes when it has lost an H+ ion
Conjugate Acid is what the base becomes what it has accepted an H+10ion
The BrØnsted-Lowry Theory
Example: HNO3 + H2O  H3O+ + NO31. Identify the reactant acid and base.
You do it!
2. Find the species that differs from the acid by a
proton, that is the conjugate base.
You do it!
3. Find the species that differs from the base by a
proton, that is the conjugate acid.
You do it!
•
•
HNO3 is the acid, conjugate base is NO3H2O is the base, conjugate acid is H3O+
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The BrØnsted-Lowry Theory
•
The major differences between Arrhenius
and Brønsted-Lowry theories.
1. The reaction does not have to occur in an
aqueous solution.
2. Bases are not required to be hydroxides.
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The BrØnsted-Lowry Theory
• An important concept in BrØnsted-Lowry theory
involves the relative strengths of acid-base pairs.
• Weak acids have strong conjugate bases.
• Weak bases have strong conjugate acids.
• The weaker the acid or base, the stronger the
conjugate partner.
• The reason why a weak acid is weak is because
the conjugate base is so strong it reforms the
original acid.
• Similarly for weak bases.
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The BrØnsted-Lowry Theory
NH3


 H 2O  NH4  OH
• The 2-way arrows implies a reversible reaction
and hence indicates that ammonia is a weak base
– Since NH3 is a weak base, NH4+ must be a strong acid.
• NH4+ gives up H+ to reform NH3.
• Compare that to
NaOH  Na+ (aq) + OH-(aq)
– Na+ must be a weak acid or it would recombine to form
NaOH
• Remember NaOH ionizes 100%.
NaOH is a strong base.
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The BrØnsted-Lowry Theory
• Amines are weak bases that behave
similar to ammonia.
• The functional group for amines is an -NH2
group attached to other organic groups.


NH3  H 2O  NH4  OH


CH3 NH2  H 2O  CH3 NH3  OH
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The Autoionization of Water
• Water can be either an acid or base in Bronsted-Lowry
theory.
– Reaction with ammonia it acts as an acid:
– Reaction with HF it acts as a base:
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The Autoionization of Water
• Whether water acts as an acid or base depends on
the other species present
• Consequently, water can react with itself.
– This reaction is called autoionization.
• One water molecule acts as a base and the other
as an acid.


H 2O  H 2O  H 3O  OH
base1
acid 2
acid 1
base 2
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The Autoionization of Water
• Water does not do this extensively.
[H3O+] = [OH-]  1.0 x 10-7 M
• Autoionization is the basis of the pH scale
• Water is said to be amphiprotic
– It can both donate and accept protons
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Amphoterism
• Other species can behave as both acids and
bases.
– are called amphoteric.
– (amiphiprotic behaviour describes the cases in
which substances exhibit amphoterism by
accepting or donating a proton).
• Examples: some insoluble metal hydroxides
– Zn and Al hydroxides
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Amphoterism
• Zn(OH)2 behaves as a base in presence of strong acids.
– Reacts with nitric acid to form a normal salt (contains
no ionizable H atoms or OH groups)
• Molecular equation
Zn(OH) 2  2 HNO3  Zn(NO 3 ) 2  2 H 2O
• Total ionic equation

Zn(OH) 2  2H  2 NO  Zn
• Net ionic equation
3

2
Zn(OH) 2  2H  Zn
2
 2 NO  2 H 2O
3
 2 H 2O
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Amphoterism
• Zn(OH)2 behaves as an acid in presence of
strong bases.
• Molecular equation
Zn(OH)2 + 2KOH K2Zn(OH)4
Zn(OH)2 is insoluble until it reacts with KOH
• Net ionic equation
Zn(OH) 2  2 OH  Zn(OH)
-
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Strengths of Acids
• The easy of ionization of binary acids depends on:
– Ease of breaking H-X bonds
– The stability of the resulting ions
• For binary acids, acid strength increases with
decreasing H-X bond strength.
• For example, the hydrohalic binary acids
• Bond strength has this periodic trend.
HF >> HCl > HBr > HI
• Acid strength has the reverse trend.
HF << HCl < HBr < HI
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Strengths of Acids
• The same trend applies to the VIA hydrides.
• Their bond strength has this trend.
H2O >> H2S > H2Se > H2Te
• The acid strength is the reverse trend.
H2O << H2S < H2Se < H2Te
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Strengths of Acids
• In dilute aqueous solutions, HCl, HBr and HI are
completely ionized and all show the same apparent
strength
• Water is sufficiently basic to mask the differences in
acid strength of the hydrohalic acids.
– Referred to as the leveling effect
• The strongest acid that can exist in water is H3O+.
– Acids that are stronger than H3O+ merely react with water
to produce H3O+.
– Consequently all strong soluble acids have the same
strength in water
HI + H2O  H3O+ + Iessentially 100%
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Strengths of Acids
• HBr, which should be a weaker acid, has the
same strength in water as HI.
HBr + H2O  H3O+ + Bressentially 100%
• Acid strength differences for strong acids
can only be distinguished in nonaqueous
solutions like acetic acid, acetone, etc.
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• it is possible to construct a relative ranking of acid and base strengths
(and their conjugate partners.)
Strengths of Acids
• The strongest acid that can exist in aqueous solution is
H3O+.
HCl + H2O  H3O+ + Cl-
– HCl is strong enough that it forces water to accept H+.
– All acids stronger than H3O+ react completely with water to form
H3O+ and their conjugate base partner.
• The strongest base that can exist in aqueous solution is
OH-.
NH2- + H2O  NH3 + OH-
– NH2- is strong enough to remove H+ from water.
– Bases stronger than OH- react completely with water to form OHand their conjugate base partner.
• The reason that stronger acids and bases cannot exist in
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water is that water is amphiprotic.
Strengths of Acids
• Acids containing 3 or more elements
• Ternary acids are hydroxides of nonmetals
that produce H3O+ in water.
– Consist of H, O, and a nonmetal.
• HClO4
H3PO4
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Strengths of Acids
• HClO4
H3PO4
O-H bonds must broken for these compounds to be acidic
Note: the acidic hydrogens are bonded to the O atoms ( and
the metal as depicted by the molecular formula)
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Strengths of Acids
• For ternary acids, acid strength also increases
with decreasing H-X bond strength.
– Strong ternary acids have weaker H-O bonds than weak
ternary acids.
• For example, compare acid strengths:
HNO2<HNO3
H2SO3< H2SO4
• This implies that the H-O bond strength is:
HNO2 > HNO3
H2SO3 > H2SO4
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Strengths of Acids
• Ternary acid strength usually increases with:
1. an increasing number of O atoms on the
central atom and
2. an increasing oxidation state of central atom.
• Effectively, these are the same phenomenon.
– Every additional O atom increases the oxidation
state of the central atom by 2.
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Strengths of Acids
• For ternary acids having the same central atom:
the highest oxidation state of the central atom is usually strongest
acid.
• For example, look at the strength of the Cl ternary acids.
HClO < HClO2 < HClO3 < HClO4
weakest
strongest
Cl oxidation states
+1
+3
+5
+7
(perchloric acid)
(Hypochlorous acid)
(chlorous acid)
(chloric acid)
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Strengths of Acids
• Other examples:
H2SO3
<
(sulfurous acid)
H2SO4
(sulfuric acid)
(stronger acids are on the right)
HNO2
(nitrous acid)
< HNO3
(nitric acid)
• Monoprotic acids have only one ionizable H e.g. HCl
• Diprotic acids have 2 ionizable H atoms e.g. H2SO4
– Polyprotic acids have more than 1 ionizable H atom e.g. H3PO4
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Acid-Base Reactions in Aqueous
Solutions
•
There are four acid-base reaction
combinations that are possible:
1.
2.
3.
4.
•
Strong acids – strong bases
Weak acids – strong bases
Strong acids – weak bases
Weak acids – weak bases
Let us look at one example of each acidbase reaction.
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Acid-Base Reactions in
Aqueous Solutions
1. Strong acids - strong bases
–
•
(a) forming soluble salts
This is one example of several possibilities
hydrobromic acid + calcium hydroxide
•
The molecular equation is:
You do it!
2 HBr(aq) + Ca(OH)2(aq)  CaBr2(aq) + 2 H2O()
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Acid-Base Reactions in
Aqueous Solutions
•
The total ionic equation is:
You do it!
2H+(aq) + 2Br-(aq) + Ca2+(aq) + 2OH-(aq) Ca2+(aq) + 2Br-(aq) +
2H2O()
•
The net ionic equation is:
You do it!
2H+ (aq) + 2OH- (aq)  2H2O()
or
H+ (aq) + OH-( aq)  H2O()
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Acid-Base Reactions in
Aqueous Solutions
1. Strong acids-strong bases
– (b) forming insoluble salts
•
There is only one reaction of this type:
sulfuric acid + barium hydroxide
•
The molecular equation is:
H2SO4(aq) + Ba(OH)2(aq)  BaSO4(s)+ 2H2O()
•
The net ionic equation is:
2H+(aq) + SO42-(aq) + Ba2+(aq) + 2OH-(aq)  BaSO4(s) + 2H2O()
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Acid-Base Reactions in
Aqueous Solutions
2. Weak acids - strong bases
– forming soluble salts
•
This is one example of many possibilities:
nitrous acid + sodium hydroxide
•
The molecular equation is:
HNO2(aq) + NaOH(aq)  NaNO2(aq) + H2O()
•
The net ionic equation is:
HNO2(aq) + OH-(aq)  NO2-(aq) + H2O()
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Acid-Base Reactions in
Aqueous Solutions
•
Reminder – there are 3 types of substances that are
written as ionized in total and net ionic equations.
1. Strong acids
2. Strong bases
3. Strongly water soluble salts
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Acid-Base Reactions in
Aqueous Solutions
3. Strong acids - weak bases
–
forming soluble salts
•
This is one example of many.
nitric acid + ammonia
•
The molecular equation is:
HNO3(aq) + NH3(aq)  NH4NO3(aq)
•
The total ionic equation is:
H+(aq) + NO3-(aq) + NH3(aq) NH4+(aq) + NO3-(aq)
•
The net equation is:
H+(aq) + NH3(aq)  NH4+(aq)
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Acid-Base Reactions in
Aqueous Solutions
4. Weak acids - weak bases
–
•
forming soluble salts
This is one example of many possibilities.
acetic acid + ammonia
•
The molecular equation is:
CH3COOH(aq) + NH3(aq)  NH4CH3COO(aq)
•
The total ionic equation is:
CH3COOH(aq) + NH3(aq)  NH4+(aq) + CH3COO-(aq)
•
The net ionic equation is:
CH3COOH(aq) + NH3(aq)  NH4+(aq) + CH3COO-(aq)
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Acidic Salts and Basic Salts
• Acidic salts are formed by the reaction of polyprotic
acids with less than the stoichiometric amount of base.
• E.g. if sulfuric acid and sodium hydroxide are reacted
in a 1:1 ratio.
H2SO4(aq) + NaOH(aq)  NaHSO4(aq) + H2O()
The acidic salt sodium hydrogen sulfate is formed.
• If sulfuric acid and sodium hydroxide are reacted in a
1:2 ratio.
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O()
The normal salt sodium sulfate is formed.
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Acidic Salts and Basic Salts
• Similarly, basic salts are formed by the reaction of
polyhydroxy bases with less than the stoichiometric
amount of acid.
• E.g. If barium hydroxide and hydrochloric acid are
reacted in a 1:1 ratio.
Ba(OH)2(aq) + HCl(aq)  Ba(OH)Cl(aq) + H2O()
The basic salt is formed.
• If the reaction is in a 1:2 ratio.
Ba(OH)2(aq) + 2HCl(aq)  BaCl2(aq) + 2H2O()
The normal salt is formed.
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Most familiar e.g. of an acidic salt
is sodium hydrogen carbonate or
baking soda (NaHCO3)
Basic aluminum salts e.g.
Al(OH)2Cl, aluminum dihydroxide
chloride and Al(OH)Cl2 aluminum
hydroxide dichloride are
components of some
antiperspirants
p. 361
Acidic Salts and Basic Salts
• Both acidic and basic salts can neutralize acids and
bases.
– However the resulting solutions are either acidic or basic
because they form conjugate acids or bases.
• Another example of BrØnsted-Lowry theory.
• This is an important concept in understanding buffers.
• An acidic salt neutralization example is:
NaHSO4(aq) + NaOH(aq)  Na2SO4 (aq) + H2O()
• A basic salt neutralization example is:
Ba(OH)Cl(aq) + HCl(aq)  BaCl2(aq) + H2O()
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The Lewis Theory
• This is the most general of the present day
acid-base theories.
• Emphasis on what the electrons are doing
as opposed to what the protons are doing.
– Acids are defined as electron pair acceptors.
– Bases are defined as electron pair donors.
• Neutralization reactions are accompanied by
coordinate covalent bond formation.
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The Lewis Theory
• One Lewis acid-base example is the ionization of
ammonia.
NH3  H 2O 
 NH 4  OH-
base acid
• Look at this reaction in more detail paying
attention to the electrons.
+
H
N
H
H
H
+
O
N
H
H
+
-
O
H
H
Base- e- pair donor
H
H
Acid- e- pair acceptor
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The Lewis Theory
• A second example is the ionization of HBr.
HBr + H2O H3O+ + Bracid
Acid- e- pair acceptor
Br
H
+
base
+
_
H
Br
O
H
+
O
H
H
H
Base- e- pair donor
The H that came from
the acid is bonded to
water via a dative or
coordinate bond
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The Lewis Theory
• The reaction of sodium fluoride and boron
trifluoride provides an example of a reaction
that is only a Lewis acid-base reaction.
– It does not involve H+ at all, thus it cannot be an
Arrhenius nor a Brønsted-Lowry acid-base
reaction.
NaF + BF3  Na+ + BF4• You must draw the detailed picture of this
reaction to determine which is the acid and
which is the base.
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The Lewis Theory
_
F
+
F
F
F
B
B
F
F
F
F
BF3 is a strong Lewis acid
It accepts an e- pair from
the fluoride ion
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The Lewis Theory
• BF3 is a strong Lewis acid. Another
example of it reacting with NH3 is shown in
this movie.
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Acid-Base Theories
•
•
Look at the reaction of ammonia and
hydrobromic acid.
NH3 + HBr NH4++ BrIs this reaction an example of:
1.
2.
3.
4.
•
Arrhenius acid-base reaction,
Brønsted-Lowry acid base reaction,
Lewis acid-base reaction,
or a combination of these?
You do it!
It is a Lewis and Brønsted-Lowry acid base
reaction but not Arrhenius.
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