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Ch. 14 Outline - Acids and Bases
A. Acids and bases
1. The Nature of Acids and Bases
a. Arrhenius definitions
- acid - produce hydrogen ions (H+) in solution
- base - produce hydroxide ions (OH-) in solution
b. Brønsted-Lowry definitions
- acid - proton (H+) donor
- base- proton acceptor - Note ammonia : HCl + NH3 ℑ NH4+ + Clc. Hydronium ion : H3O+ ( a hydrogen ion attached to a water molecule-H2O·H+)
- general reaction for an acid in water :
HA(aq)
Acid
+
H2O(l) ℑ
Base
H3O+
+
Conjugate acid
A-(aq)
Conjugate base
- conjugate base -what is formed when an acid loss the hydrogen ion
- conjugate acid- what is formed when a substance gains the hydrogen ion
- note : if A- has a stronger attraction to the proton the equilibrium position will be far to the left
d. Acid dissociation constant (Ka). Ka for the above reaction would be :
Ka = [H3O+][A-]/ [HA]
- if Ka > 1 the equilibrium position favors the products
- if Ka < 1 the equilibrium position favors the reactants
2. Acid Strength
a. a strong acid is one for which the equilibrium will lie far to the right - dissociates to a large extent
- a strong acid yields a weak conjugate base (much weaker than water)
- common strong acids : H2SO4, HCl, HBr, HI, HNO3, and HClO4
- strong acids have the equilibrium position so far to the right that the undissociated acid
concentration is so small it cannot be accurately measured, making it hard to calculate the Ka of
strong acids
-monoprotic (one ionizing hydrogen), diprotic (two ionizing hydrogens) triprotic (three ionizing
hydrogens)
b. a weak acid is one for which the equilibrium position will lie far to the left
- a weak acid yields a relatively strong conjugate base
Property
Strong Acid
Weak Acid
Ka value
large
small
Position of the dissociation Far to right
Far to left
(ionization) equilibrium
Equilibrium [H+] compared [H+] ≈ [HA]0
[H+] << [HA]0
to [HA]
Strength of conjugate base
A- is a much weaker conjugate A- is a much stronger
compared to water
base than water
conjugate base than water
c. Oxyacids - acids which have the proton attached to an oxygen molecule (e.g. H3PO4, HNO3, H2SO4)
d. Organic acids - acids with a carbon backbone-the hydrogen is commonly part of the carboxyl
group(-COOH)
- usually weak
-example : acetic acid HC2H3O2 (also CH3COOH)
e. Amphoteric substances (e.g. water) - can behave as either an acid or a base
-autoionization of water : 2 H2O ℑ H3O+ + OH-
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- in a neutral solution [H+] = [OH-]=1.0 x 10 mol/L
- in an acidic solution [H+] > [OH-]
- in a basic solution [H+] < [OH-]
- ion-product constant (dissociation constant) of water(Kw) at 25 ºC = [H+][OH-] = 1.0 x 10-14
- since Kw is an equilibrium constant it will vary with temperature
3. The pH scale
Rule for significant figures and logs : The number
a. pH = -log [H+]
of decimal places in the log is equal to then
b. The pH changes by 1 for every factor of
number of significant figures in the original
10 change in the [H+]
number.
- pH decreases as [H+] increases
c. pOH = -log[OH-]
d. pKw = 14.00 = pH + pOH
4. Calculating the pH of Strong Acid Solutions
a. focus on the major species - minor species like water can be ignored if their H+ contribution is very
small compared to the major species and vice versa
- in calculating the pH of .10 M HCl the contribution of water (10-7) can be ignored
- in calculating the pH of 1.0 x 10-12 HNO3, the [H+] contributed by the autoionization of water
will be much larger and the contribution of the 1.0 x 10-12 HNO3 can be ignored
5. Calculating the pH of Weak Acid Solutions
a. Steps to solving :
1. Write the major species in solution.
2. Determine which of the major species can contribute H+ ions.
- as in strong acid solutions any minor species can be ignored (use Ka values to determine)
3. Solve the problem as an equilibrium problem.
- 5% rule - if x / [HA]0 x 100% is less than5 then we can safely assume [HA]0- x ≈ [HA]0
b. Calculating the pH of a Mixture of Weak Acids
- follow same approach as above
c. Percent Dissociation
- percent dissociation = amount dissociated (mol/L) / initial concentration (mol/L) x 100%
- for weak acids the percent dissociation increases as the acid becomes more dilute
- calculation of Ka from percent dissociation
6. Bases
a. Group 1A and Group 2A bases are strong, although Group 2A bases are only slightly soluble and do
not produce high concentrations of OH- ions in solution (and therefore make good antacids)
- Ca(OH)2 is often called slaked lime and is used in industry to remove sulfur dioxide :
-steps : (1) SO2 + H2O ℑ H2SO3
(2) Ca(OH)2 + H2SO3 ℑ CaSO3(s) + 2 H2O
- slaked lime is also used to soften water by precipitating out the calcium ions
b. many substances will cause hydroxide ions to form in aqueous solution even though they do not
contain hydroxide ions
- ex. Ammonia, amines - contain unshared pairs of electrons that bond to the proton of water and
leave the OH- :
- general reaction : B(aq) + H2O ℑ BH+ + OH- (weak bases)
- Kb = [BH+][OH-] / [B]
- Kb - refers to the reaction of a base with water to form the conjugate acid and the hydroxide ion
-calculation of pH of weak bases- similar to the calculation of pH from weak acids
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7. Polyprotic Acids- contain more than one proton
a. polyprotic acids dissociate in steps and therefore have a Ka1, Ka2 etc.
- Ka1 > Ka2 > Ka3
- generally only the first dissociation (Ka1) is important for the calculation of pH (other Ka values
are small and can be ignored)
-only in dilute (less than 1.0 M) solutions of H2SO4 does the second dissociation contribute
significantly to [H+]
8. Acid-base Properties of Salts
a. salts that are formed from a strong acids and a strong base form neutral solutions(no effect on [H+])
b. salts formed from a strong base and a weak acid form a basic solution
- the anion has a strong affinity for the H+ ion
e.g. acetic acid and KOH form water and potassium acetate (KC2H3O2). The acetate ion reacts
with water to form OH- ions :
C2H3O2 + H2O ℑ HC2H3O2 + OHc. Base Strength in Aqueous Solution
- key concept -Which species competes the most strongly for the H+ ion, the OH- ion or the anion
from the acid?
d. salts formed from a strong acid and a weak base form acidic solutions
- the cation has a strong affinity for the hydroxide ion
- e.g. ammonium chloride:
NH4+ (aq) ℑ NH3(aq) + H+(aq)
e. salts of highly charged metallic ions form acidic solutions - high charge polarizes water making
hydrogen ions more likely to be released
- e.g. Al3+ forms the hydrated Al(H2O)63+ which is a weak acid :
Al(H2O)63+(aq) ℑ Al(OH)(H2O)52+(aq) + H+(aq)
f. for salts formed from weak acids and weak bases
- it depends on the K value
- if Ka = Kb, the solution will be neutral
- if Ka > Kb, the solution will be acidic
- if Ka < Kb, the solution will be basic
9. The Effect of Structure on Acid-Base Properties
a. The two main factors in determining whether a molecule containing a X-H bond will behave as a
Brønsted-Lowry acid are
- the strength of the X-H bond
- the polarity of the X-H bond
Molecule
Bond strength
Bond polarity
Acid Strength in
(kJ/mol)
water
H-F
565
Most polar
Weak
H-Cl
427
2nd most polar
Strong
H-Br
363
2nd least polar
Strong
H-I
295
Least polar
strong
- oxyacids- as the number of oxygen atoms increases the polarity of the O-H bond increases and
the bond strength weakens
Molecule
Ka value
HClO
3.5 x 10-8
HClO2
1.2 x 10-2
HClO3
~1
HClO4
Large (~107)
- the more highly charged a metallic ion is, the more acidic the hydrated ion becomes -see
discussion of aluminum ion above
10. Acid-Base Properties of Oxides
a. Consider H-O-X :
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-if O-X bond is strong due to high electronegativity of X, the O-X bond will remain in aqueous
solution forming an acidic solution (H+ ions by themselves)
- if O-X bond is weak due to low electronegativity of X, the H-O bond will remain intact and form
a basic solution (e.g. KOH, NaOH)
b. acidic oxides - form from elements with high electronegativities - nonmetallic (covalent) oxides
- e.g. SO2 : SO2(g) + H2O(l) Æ H2SO3(aq)
- e.g. CO2 : CO2(g) + H2O(l) Æ H2CO3(aq)
- e.g. NO2 : 2NO2(g) + H2O(l) Æ HNO3(aq) + HNO2(aq)
c. basic oxides - metallic oxides
- e.g. CaO(s) + H2O(l) Æ Ca(OH)2(aq)
- e.g. K2O(s) + H2O(l) Æ 2KOH(aq)
11. The Lewis Acid-Base Model
a. a Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor
- e.g. NH3
b. Summary of three acid-base models :
Model
Definition of Acid
Definition of Base
Arrhenius
H+ producer
OH- producer
Brønsted-Lowry
H+ donor
H+ acceptor
Lewis
Electron-pair acceptor
Electron-pair donor
12. Summary of strategies for solving acid-base problems
a. steps :
1. List the major species in solution.
2. Look for reactions that go to completion - strong acids or strong bases.
3. For reactions that go to completion :
- determine concentrations of products
- write down the major species in solution after the reaction
4. Look at each component ands determine if it is an acid or a base.
5. Pick the equilibrium that will control the pH - use values of dissociation constants to make this
determination
- write the equation for the reaction and the equilibrium expression
- calculate the initial concentrations
- define x
- calculate the equilibrium concentrations in terms of x
- substitute the concentrations into the equilibrium expression and solve for x
- check the validity of the approximation (5% rule)
- calculate the pH and other concentrations required