Download factors affecting strength of acids

ANALYTICAL CHEMISTRY
CHEM 3811
CHAPTER 8
DR. AUGUSTINE OFORI AGYEMAN
Assistant professor of chemistry
Department of natural sciences
Clayton state university
CHAPTER 8
ACIDS AND BASES
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
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
bitter taste, change red litmus paper to blue, slippery to touch
BRONSTED-LOWRY ACIDS
- Acids are proton (H+) donors
- Not restricted to aqueous solutions
HCl, HNO3, H2SO4
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
LEWIS ACIDS
- Acids are electron pair acceptors
- Not restricted to protons or aqueous solutions
BF3, B2H6, Al2Cl6, AlF3, PCl5
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
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
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+)
CONJUGATE ACID BASE PAIRS
HX(aq) + H2O(l)
X-(aq) + H3O+(aq)
- HX donates a proton to H2O to form XHX 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
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)
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)
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+
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)
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
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
AUTOPROTOLYSIS OF WATER
- Pure water molecules (small percentage) interact with one
another to form equal amounts of H3O+ and OH- ions
Kw
H3O+ + OH-
H2O + H2O
reduces to
Kw
H2O
H+ + OH-
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
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)
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
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-]
THE pH CONCEPT
pH
- Negative logarithm of the hydronium ion concentration
[H3O+] in an aqueous solution
pH = - log[H3O+]
[H3O+] = 10-pH
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
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
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
THE pH CONCEPT
pKw = -log(Kw) = -log(1.00 x 10-14) = 14
pOH = -log[OH-]
[H3O+][OH-] = Kw
Implies that
pH + pOH = pKw
pH + pOH = 14.00
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 (< 5%) of their protons
to H2O in aqueous solution
Amino acids, Organic acids: acetic acid, citric acid
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
STRENGTH OF ACIDS
H3O+(aq) + A-(aq)
HA(aq) + H2O(l)
- Equilibrium constant for the reaction of a weak acid with water
- Represented by Ka (acid dissociation constant)


[H 3O ][A ]
Ka 
[HA]
- 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)
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
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
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)
[BH  ][OH  ]
Kb 
[B]
- H2O is a pure liquid so not included
WEAK ACIDS AND BASES


[H 3O ][A ]
Ka 
[HA]
[BH  ][OH  ]
Kb 
[B]
Ka x Kb = [H3O+][OH-] = Kw = 1.00 x 10-14
- Reaction goes to completion when Ka value is very large
- Weak acids have small Ka values
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
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
METAL IONS
- Many metal ions bind with four or six water molecules
Mn+ + 6H2O
M(H2O)6n+
Ka
M(H2O)6n+
M(H2O)5(OH)(n-1)+ + H+
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
pH OF STRONG BASES
Find the pH of 3.9 x 10-2 M NaOH
[H3O+][OH-] = Kw = 1.0 x 10-14
[H3O+][3.9 x 10-2] = 1.0 x 10-14
[H3O+] = 2.6 x 10-13
pH = - log(2.6 x 10-13) = 12.59
pH OF STRONG BASES
Find the pH of 3.9 x 10-2 M NaOH
Alternatively
pOH = - log[OH-]
pOH = - log(3.9 x 10-2) = 1.41
pH + pOH = 14
pH = 14 - 1.41 = 12.59
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
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
WEAK ACID EQUILIBRIUM
For a weak acid HA
Ka
HA
A- + H+
cHA = total concentration = analytical concentration
= [HA] + [A-]
WEAK ACID EQUILIBRIUM
For a weak acid HA
Ka
HA
A- + H+
[A  ]
[A  ]
Fraction of Dissociati on  

[A ]  [HA] CHA
- Fraction of dissociation increases with increasing acid strength
- Fraction of dissociation increases with dilution
WEAK ACID EQUILIBRIUM
For a weak acid HA
Ka
HA
A- + H+
[H  ][A  ]
x2

 Ka
[HA]
Fx
- 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
WEAK ACID EQUILIBRIUM
[H  ][A  ]
x2

 Ka
[HA]
Fx
- The equation reduces to
[H  ][A  ] x 2

 Ka
[HA]
F
- If x ≤ 5% of F
That is F – x ≈ F if x ≤ 0.05F
WEAK BASE EQUILIBRIUM
For a weak base B
Kb
B + H2O
BH+ + OH-
[BH  ]
Fraction of Associatio n  
[B ]  [B]
WEAK BASE EQUILIBRIUM
For a weak base B
Kb
B + H2O
BH+ + OH-
[BH  ][OH  ]
x2

 Kb
[B]
Fx
- 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
WEAK BASE EQUILIBRIUM
[BH  ][OH  ]
x2

 Kb
[B]
Fx
- The equation reduces to
[BH  ][OH  ] x 2

 Kb
[B]
F
- If x ≤ 5% of F
That is F – x ≈ F if x ≤ 0.05F
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
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
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
FACTORS AFFECTING STRENGTH OF ACIDS
Example
Bond strength of hydrogen halides
HF > HCl > HBr > HI
Acidity of hydrogen halides
HF < HCl < HBr < HI
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 < NH3 < H2O < HF
FACTORS AFFECTING STRENGTH OF ACIDS
- Bond strength and electronegativity sometimes
predict opposite trends
- Bond strength dominates down a group
- Electronegativity dominates across a period
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