ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 8

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ANALYTICAL CHEMISTRY CHEM 3811CHAPTER 8
DR. AUGUSTINE OFORI AGYEMAN Assistant professor
of chemistry Department of natural
sciences Clayton state university
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CHAPTER 8 ACIDS AND BASES
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ARRHENIUS ACIDS
- Acids are substances that ionize in aqueous
solutions to produce hydrogen ions (proton,
H) HCl, HNO3, H2SO4 - Arrhenius acids are
covalent compounds in the pure state Properties
sour taste, change blue litmus paper to red,
corrosive
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ARRHENIUS BASES
- Bases are substances that ionize in aqueous
solutions to produce hydroxide ions
(OH-) NaOH, KOH, Ca(OH)2 - Arrhenius bases are
ionic compounds in the pure state Properties bit
ter taste, change red litmus paper to blue,
slippery to touch
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BRONSTED-LOWRY ACIDS
- Acids are proton (H) donors - Not restricted
to aqueous solutions HCl, HNO3, H2SO4
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BRONSTED-LOWRY BASES
- Bases are proton acceptors - Not restricted
to aqueous solutions NH3, dimethyl sulfoxide
(DMSO) - Proton donation cannot occur unless an
acceptor is present
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LEWIS ACIDS
- Acids are electron pair acceptors - Not
restricted to protons or aqueous solutions BF3,
B2H6, Al2Cl6, AlF3, PCl5
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LEWIS BASES
- Bases are electron pair donors - Not
restricted to protons or aqueous solutions NH3,
ethers, ketones, carbon monoxide, sulfoxides -
The product of a Lewis acid-base reaction is
known as an adduct - The base donates an
electron pair to form coordinate covalent bond
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ACIDS
Monoprotic Acid - Donates one proton per molecule
(HNO3, HCl) Diprotic Acid - Donates two protons
per molecule (H2SO4, H2CO3) Triprotic Acid -
Donates three proton per molecule (H3PO4,
H3AsO4) Polyprotic Acid - Donates two or more
protons per molecule
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CONJUGATE ACID BASE PAIRS
- Most Bronsted-Lowry acid-base reactions do not
undergo 100 conversion - Acid-base
equilibrium is established - Every acid has a
conjugate base associated with it (by removing
H) - Every base has a conjugate acid associated
with it (by adding H)
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CONJUGATE ACID BASE PAIRS
HX(aq) H2O(l)
X-(aq) H3O(aq)
- HX donates a proton to H2O to form X- HX is
the acid and X- is its conjugate base - H2O
accepts a proton from HX H2O acts as a base and
H3O is its conjugate acid
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CONJUGATE ACID BASE PAIRS
NH3(aq) H2O(l)
NH4(aq) OH-(aq)
HNO3(aq) H2O(l)
H3O(aq) NO3-(aq)
HF(aq) H2O(l)
H3O(aq) F-(aq)
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AMPHOTERIC SUBSTANCES
- A substance that can lose or accept a proton -
A substance that can function as either
Bronsted-Lowry acid or Bronsted-Lowry base - H2O
is the most common (refer to previous slide for
examples)
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REACTIONS OF ACIDS AND BASES
Arrhenius acid Arrhenius base ? salt
water HCl NaOH ? NaCl H2O B-L acid
B-L base ? conjugate base conjugate
acid H3PO4 H2O ? H2PO4- H3O
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SALTS
- Salts are ionic compounds - The positive ion
is a metal or polyatomic ion - The negative ion
is a nonmetal or polyatomic ion exception is
the hydroxide ion (OH-) - Salts dissociate
completely into ions in solution - A reaction
between an acid and a hydroxide base produces
salt (cation from the base and anion from the
acid)
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SALTS
- Solutions of salts may be acidic, basic, or
neutral - Acidity depends on relative values of
Ka of the cation and Kb of the anion - The
conjugate base of a strong acid (anion from a
strong acid) has no net effect on the pH of a
solution (spectator ion) Cl- from HCl, NO3- from
HNO3 - Cation from a strong base has no net
effect on the pH of a solution (spectator
ion) Na from NaOH, K from KOH
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SALTS
- NaCl solution contains Na and Cl- ions -
Both ions are spectator ions and do not affect
the pH of the solution - pH is determined by
autoionization of water
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AUTOPROTOLYSIS OF WATER
- Pure water molecules (small percentage)
interact with one another to form equal
amounts of H3O and OH- ions
Kw
H2O H2O
H3O OH-
reduces to
Kw
H OH-
H2O
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AUTOPROTOLYSIS OF WATER
- The number of H3O and OH- ions present in a
sample of pure water at any given time is
small - At equilibrium (24oC) H3O OH-
1.00 x 10-7 M - H3O hydronium ion
concentration - OH- hydroxide ion
concentration
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AUTOPROTOLYSIS OF WATER
- The ion product constant of water H3O x
OH- (1.00 x 10-7) x (1.00 x 10-7) 1.00 x
10-14 - Valid in all solutions (pure water and
water with solutes)
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AUTOPROTOLYSIS OF WATER
Addition of Acidic Solute - increases H3O -
OH- decreases by the same factor to make
product 1.00 x 10-14 Addition of Basic Solute
- increases OH- - H3O decreases by the
same factor to make product 1.00 x 10-14
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THE pH CONCEPT
Acidic Solution - An aqueous solution in which
H3O is higher than OH- Basic Solution - An
aqueous solution in which OH- is higher than
H3O Neutral Solution - An aqueous solution
in which H3O is equal to OH-
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THE pH CONCEPT
pH - Negative logarithm of the hydronium ion
concentration H3O in an aqueous
solution pH - logH3O H3O 10-pH
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THE pH CONCEPT
- For H3O coefficient of 1.0 - Expressed in
exponential notation - The pH is the negative of
the exponent value
H3O 1.0 x 10-5 M, then pH 5.00 H3O
1.0 x 10-3 M, then pH 3.00 H3O 1.0 x
10-11 M, then pH 11.00
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THE pH CONCEPT
- For neutral solutions pH is equal to 7.00 -
For acidic solutions pH is less than 7.00 - For
basic solutions pH is greater than 7.00 -
Increasing H3O lowers the pH
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THE pH CONCEPT
- A change of 1 unit in pH corresponds to a
tenfold change in H3O pH 3.00 implies
H3O 1.0 x 10-3 M 0.0010 M pH 2.00
implies H3O 1.0 x 10-2 M 0.010 M which
is tenfold - The pH meter and the litmus paper
are used to determine pH values of solutions
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THE pH CONCEPT
pKw -log(Kw) -log(1.00 x 10-14) 14 pOH
-logOH- H3OOH- Kw Implies that pH
pOH pKw pH pOH 14.00
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STRENGTH OF ACIDS
Strong Acids - Transfer 100 (or very nearly
100) of their protons to H2O in aqueous
solution - Completely or nearly completely ionize
in aqueous solution - Strong electrolytes HCl,
HBr, HClO4, HNO3, H2SO4 Weak Acids - Transfer
only a small percentage (lt 5) of their protons
to H2O in aqueous solution Amino acids,
Organic acids acetic acid, citric acid

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STRENGTH OF ACIDS
- Equilibrium position lies to the far right for
strong acids
HA(aq) H2O(l)
H3O(aq) A-(aq)
- Predominant species are H3O and A-
- Equilibrium position lies to the far left for
weak acids
HA(aq) H2O(l)
H3O(aq) A-(aq)
- Predominant species is HA
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STRENGTH OF ACIDS
HA(aq) H2O(l)
H3O(aq) A-(aq)
- Equilibrium constant for the reaction of a weak
acid with water - Represented by Ka (acid
dissociation constant)
- H2O is a pure liquid so not included - Acid
strength increases with increasing Ka value - For
polyprotic acids, Ka for each dissociation step
is smaller than the previous step (weaker acid)
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STRENGTH OF BASES
Strong Bases - Completely or nearly completely
ionize in aqueous solution - Strong
electrolytes Hydroxides of Groups IA and IIA
are strong bases LiOH, CsOH, Ba(OH)2, Ca(OH)2
most common in lab NaOH and KOH Weak bases -
produce small amounts of OH- ions in aqueous
solution Organic bases, methylamine, cocaine,
morphine most common NH3
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STRENGTH OF BASES
- Weak bases produce small amounts of OH- ions in
aqueous solution (NH3)
NH3(g) H2O(l)
NH4(aq) OH-(aq)
- Equilibrium position lies to the far left -
Small amounts of NH4 and OH- ions are
produced - The name aqueous ammonia is preferred
over ammonium hydroxide
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STRENGTH OF BASES
B(aq) H2O(l)
BH(aq) OH-(aq)
- Equilibrium constant for the reaction of a weak
base with water - Represented by Kb (base
hydrolysis constant)
- H2O is a pure liquid so not included
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WEAK ACIDS AND BASES
Ka x Kb H3OOH- Kw 1.00 x 10-14 -
Reaction goes to completion when Ka value is very
large - Weak acids have small Ka values
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WEAK ACIDS AND BASES
pKa - logKa pKb - logKb pKa pKb
pKw - The stronger an acid the smaller its
pKa - The stronger the acid the weaker its
conjugate base - The stronger the base the
weaker its conjugate acid
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METAL IONS
- Metal ions with 1 charge have negligible
acidity - Metal ions with 2 charge or higher
are acidic - The higher the charge the more
acidic the metal Example Fe2 Ka 4 x 10-10,
Fe3 Ka 6.5 x 10-3 - Solutions of metal salts
are usually acidic - Metal ions are Lewis acids
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METAL IONS
- Many metal ions bind with four or six water
molecules
M(H2O)6n
Mn 6H2O
Ka
M(H2O)6n
M(H2O)5(OH)(n-1) H
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pH OF STRONG ACIDS
- Differences in acidities of strong acids cannot
be measured since they all ionize completely -
This phenomenon is known as leveling effect
Find the pH of 3.9 x 10-2 M HCl HCl is a
strong acid and dissociates completely HCl(aq)
? H(aq) Cl-(aq) pH - log(3.9 x 10-2)
1.41
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pH OF STRONG BASES
Find the pH of 3.9 x 10-2 M NaOH H3OOH-
Kw 1.0 x 10-14 H3O3.9 x 10-2 1.0 x
10-14 H3O 2.6 x 10-13 pH - log(2.6 x
10-13) 12.59
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pH OF STRONG BASES
Find the pH of 3.9 x 10-2 M NaOH Alternatively p
OH - logOH- pOH - log(3.9 x 10-2)
1.41 pH pOH 14 pH 14 - 1.41 12.59
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pH OF STRONG ACIDS AND BASES
- For dilute solutions the contribution of H2O
should not be neglected - Acids and bases
suppress water ionization What concentrations of
H and OH- are produced by H2O dissociation in
1.0 x 10-3 M HCl? pH 3 OH- Kw/H3O 1.0
x 10-11 OH- is produced from the dissociation of
H2O Implies H2O dissociation OH- H3O
1.0 x 10-11
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pH OF STRONG ACIDS AND BASES
- For dilute solutions the contribution of H2O
should not be neglected - Acids and bases
suppress water ionization What concentrations of
H and OH- are produced by H2O dissociation in
1.0 x 10-4 M KOH? H3O Kw/OH- 1.0 x
10-10 H3O (or H) is produced from the
dissociation of H2O Implies H2O dissociation
OH- H3O 1.0 x 10-10
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WEAK ACID EQUILIBRIUM
For a weak acid HA
Ka
HA
A- H
cHA total concentration analytical
concentration HA A-
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WEAK ACID EQUILIBRIUM
For a weak acid HA
Ka
HA
A- H
- Fraction of dissociation increases with
increasing acid strength - Fraction of
dissociation increases with dilution
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WEAK ACID EQUILIBRIUM
For a weak acid HA
Ka
HA
A- H
- Assume H A- - F is the initial (formal)
concentration of HA - Initial concentration of H
and A- is 0 each - Final concentration of H and
A- is x each - The iCe table may be used for such
problems
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WEAK ACID EQUILIBRIUM
- The equation reduces to
- If x 5 of F That is F x F if x 0.05F
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WEAK BASE EQUILIBRIUM
For a weak base B
Kb
B H2O
BH OH-
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WEAK BASE EQUILIBRIUM
For a weak base B
Kb
B H2O
BH OH-
- Assume BH OH- - F is the initial
(formal) concentration of B - Initial
concentration of BH and OH- is 0 each - Final
concentration of BH and OH- is x each - The iCe
table may be used for such problems
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WEAK BASE EQUILIBRIUM
- The equation reduces to
- If x 5 of F That is F x F if x 0.05F
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MIXTURES OF ACIDS
- When determining the pH of a mixture of
acids only the pH of the strongest acid is
considered - Contributions by the weaker acids
towards pH are neglected - A weak acid produces
fewer protons in the presence of a strong
acid Similarly - A weak base produces fewer
hydroxide ions in the presence of a strong base
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FACTORS AFFECTING STRENGTH OF ACIDS
- Key factors are the strength of the H A bond
and the stability of the A- ion Binary Acid
(HA) - An acidic compound composed of hydrogen
and one other element (mostly a nonmetal) HCl,
HI, HBr, H2S, H2O
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FACTORS AFFECTING STRENGTH OF ACIDS
Bond Strength of Binary Acids - Generally
decreases down the groups of the periodic
table - Due to increasing size of the other
element - Acidity increases down the groups of
the periodic table - Due to decreasing bond
strength
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FACTORS AFFECTING STRENGTH OF ACIDS
Example Bond strength of hydrogen halides HF gt
HCl gt HBr gt HI Acidity of hydrogen halides HF
lt HCl lt HBr lt HI
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FACTORS AFFECTING STRENGTH OF ACIDS
Stability of the A- Anion - Depends on the
ability of the A atom to accept additional
negative charge - Electronegativity is the
factor - A more electronegative atom results in
a stronger acid - Acidity of nonmetal hydrides
increases across periods of the periodic
table CH4 lt NH3 lt H2O lt HF
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FACTORS AFFECTING STRENGTH OF ACIDS
- Bond strength and electronegativity sometimes
predict opposite trends - Bond strength
dominates down a group - Electronegativity
dominates across a period
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FACTORS AFFECTING STRENGTH OF ACIDS
Oxyacids - Acids containing hydrogen, oxygen, and
a third element The third element may be a -
Nonmetal HNO3, H2SO4, H3PO4 - A transition metal
with high oxidation state H2CrO4 - Carbon in
organic acids CH3COOH - Acidity increases with
electronegativity of the third element -
Hypohalous acids (H O X), X Cl, Br, I