Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
NOTES - Unit 8: Acids & Bases
PART 1: Acid/Base Theory & Properties
BrØnsted-Lowry: a theory of proton transfer
o A Bronsted-Lowry ACID is a _______________________________________________.
o A Bronsted-Lowry BASE is a _______________________________________________.
Conjugate pairs: Acids react to form bases and vice versa. The acid-base pairs related to each other in this way are
called conjugate acid-base pairs. They differ by just one proton.
HA + B  A- + BH+
Ex) List the conjugate acid-base pairs in the following reaction: CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq)
Ex) Write the conjugate base for each of the following.
a) H3O+
b) NH3
c) H2CO3
Ex) Write the conjugate acid for each of the following.
a) NO2b) OH-
c) CO32-
Amphoteric / __________________________ substances: substances which can act as Bronsted-Lowry acids and bases,
meaning they can accept or donate a proton (capable of both). These features enable them to have a “double-identity:”
1) To act as a Bronsted-Lowry acid, they must be able to dissociate and ________________________________.
2) To act as a B-L base, they must be able to accept H+, which means they must have a lone pair of electrons.
Water is a prime example – it can donate H+ and it has two lone pairs of electrons.
 Auto-ionization of water:
H2O + H2O  H3O+ + OH
Water reacting as a base with CH3COOH:
CH3COOH(aq) + H2O(l)  CH3COO- (aq) + H3O+ (aq)

Water reacting as an acid with NH3:
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
Ex) Write equations to show HCO3- reacting with water (a) as an acid and (b) as a base.
Lewis: a theory of electron pairs
o A Lewis ACID is an _______________________________________________.
o A Lewis BASE is an _______________________________________________.
Lewis acid-base reactions result in the formation of a covalent bond, which will always be a _______________________
bond (a.k.a. ___________________________________________ bond) because both the electrons come from the base.
1
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Example: (note – the “curly arrow” is a convention used to show donation of electons.)
Example: (note – boron has an incomplete octet, so it is able to accept an electron pair)
Example: (note – metals in the middle of the periodic table often form ions with vacant orbitals in their d subshell, so
they are able to act as Lewis acids and accept lone pairs of electrons when they bond with ligands to form complex ions.
Ligands, as donors of lone pairs, are therefore acting as Lewis bases)
Typical ligands found in complex ions include H2O, CN- and NH3. Note that they all have lone pairs of electrons, the
defining feature of their Lewis base properties.
Table 8.1: Acid-base theory comparison
Theory
Definition of acid
Bronsted-Lowry Proton donor
Lewis
Electron pair acceptor
Definition of base
Proton acceptor
Electron pair donor
Ex: For each of the following reactions, identify the Lewis acid and the Lewis base.
a) 4NH3(aq) + Zn2+(aq)  [Zn(NH3)4]2+(aq)
b) 2Cl-(aq) + BeCl2 (aq) +  [BeCl4]2- (aq)
c) Mg2+(aq) + 6H2O(l)  [Mg(H2O)6]2+(aq)
Ex: Which of the following could not act as a ligand in a complex ion of a transition metal?
a) Clb) NCl3
PCl3
d) CH4
2
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Properties of acids and bases
For acids and bases here, we will use the following definitions:
 Acid: a substance that donates H+ in solution
 Base: a substance that can neutralize an acid to produce water --- includes metal oxides, hydroxides, ammonia,
soluble carbonates (Na2CO3 and K2CO3) and hydrogencarbonates (NaHCO3 and KHCO3)
 Alkali: a soluble base. When dissolved in water, alkalis all release the hydroxide ion, OH-. For example:
o K2O(s) + H2O(l)  2K+(aq) + 2OH-(aq)
o NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
o CO32- (aq) + H2O(l)  HCO3-(aq) + OH-(aq)
o HCO3-(aq)  CO2(g) + OH-(aq)
 Neutralization: net ionic equation =
Acid-Base Indicators
Acid-Base indicators change color reversibly according to the concentration of H+ ions in solution.
Many indicators are derived from natural substances such as extracts from flower petals and berries. ______________,
a dye derived from lichens, can distinguish between acids and alkalis, but cannot indicate a particular pH. For this
purpose, _____________________________________ was created by mixing together several indicators; thus universal
indicator changes color many times across a range of pH levels.
Table 8.2: Some common acid-base indicators
Indicator
Color in acid
Litmus
Pink
methyl orange
Red
phenolphthalein
Colorless
Color in alkali
blue
yellow
pink
Figure 8.1: Universal Indicator (pH 014)
Acids react with metals, bases and carbonates to form salts
1. Neutralization reactions with bases: acid + base  salt + water
a. With hydroxide bases
b. With metal oxide bases
c. With ammonia (via ammonium hydroxide)
2. With reactive metals (those above copper in the reactivity series): acid + metal  salt + hydrogen
3. With carbonates (soluble or insoluble) / hydrogencarbonates: acid + carbonate  salt + water + carbon dioxide
3
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Strong, Concentrated and Corrosive
In everyday English, strong and concentrated are often used interchangeably. In chemistry, they have distinct meanings:
 strong: completely _______________________________ into ions
 concentrated: high number of __________________ of solute per liter (dm3) of solution
 corrosive: chemically ___________________________
Similarly, weak and dilute also have very different chemical meanings:
 weak: only slightly dissociated into ions
 dilute: a low number of moles of solute per liter (dm3) of solution
Strong and weak acids and bases
Consider the acid dissociation reaction: HA(aq)  H+(aq) + A-(aq)
Strong acid: equilibrium lies to the right (acid dissociates fully)  reversible rxn is negligible  exists entirely as ions
Ex:
Weak acid: equilibrium lies to the left (partial dissociation)  exists almost entirely in the undissociated form
Ex:
Similarly, the strength of a base refers to its degree of dissociation in water.
Strong base ex:
Weak base ex:
NOTE: Weak acids and bases are much more common than strong acids and bases.
Table 8.3: Strong and Weak Acids and Bases you should know
Strong Acids
Strong Bases
Weak Acids
Weak Bases
(only six; know 1 three for IB)
(Grp 1 hydroxides & barium hydroxide)
carboxylic and carbonic acids
ammonia and amines
H2SO4, sulfuric acid*
LiOH, lithium hydroxide
CH3COOH, ethanoic acid
C2H5NH2, ethylamine
and other organic acids
and other amines
H2CO3, carbonic acid
NH3, ammonia
Note CO2(aq) = H2CO3(aq)
Note NH3(aq) = NH4OH(aq)
st
HNO3, nitric acid
HCl, hydrochloric acid
HI, hydroiodic acid
HBr, hydrobromic acid
HClO4, perchloric acid
NaOH, sodium hydroxide
KOH, potassium hydroxide
Ba(OH)2, barium hydroxide
H3PO4, phosphoric acid
*NOTE: Sulfuric acid, H2SO4, is a diprotic acid which is strong in the dissociation of the first H+ and weak in the dissociation of the second H+.
For purposes of IB, only monoprotic dissociations are considered.
Experimental methods for distinguishing between strong and weak acids and bases
1. Electrical conductivity: strong acids and bases will have a higher conductivity (higher concentration of mobile ions)
2. Rate of reaction: faster rate of rxn with strong acids (higher concentration of ions)
3. pH: measure of H+ concentration in sol’n. A 1.0 M sol’n of strong acid will have lower pH than 1.0 M sol’n of weak
acid; 1.0 M sol’n of strong base will have higher pH than 1.0 M sol’n of weak base
4
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
PART 2: pH, pOH & pKw
The pH Scale
 pH is a value chemists use to give a measure of the acidity or alkalinity of a solution.
 Used because [H+] is usually very small
 pH stands for pouvoire of hydrogen.
o Pouvoire is French for “power.”
o The normal range of the pH scale is 0-14.
o However, it is possible (if the hydronium or hydroxide concentrations get above 1 Molar) for the pH to go
beyond those values.
 pH= -log[H+]
 [H+] = 10-pH
 As pH decreases, [H+] increases exponentially (a change of one pH unit represents a 10-fold change in [H+]
Example: If the pH of a solution is changed from 3 to 5, deduce how the hydrogen ion concentration changes.
Calculations involving acids and bases
 Sig figs for Logarithms (see page 631): The rule is that the number of decimal places in the log is equal to the number
of significant figures in the original number. Another way of saying this is only numbers after decimal in pH are
significant.
Example: [H+] = 1.0 x 10-9 M (2 significant figures)  pH = -log(1.0 x 10-9) = 9.00 (2 decimal places)
 Ion product constant of water, Kw
o Recall that water autoionizes: H2O(l)  H+(aq) + OH-(aq)
(endothermic)
o Therefore Kc =

o
The concentration of water can be considered to be constant because so little of it ionizes, and it can
therefore be combined with Kc to produce a modified equilibrium constant known as kw. In fact, liquids and
solids never appear in equilibrium expressions for this reason.
o
Therefore, Kw =
o At 25C, Kw = 1.00 x 10-14
o In pure water, because [H+]=[OH-], it follows that [H+]=
o So at 25C, [H+] = 1.0 x 10-7, which gives pH = 7.00
Kw is temperature dependent
o Since the dissociation of water reaction in endothermic (bonds breaking), an increase in temperature will
shift the equilibrium to the ________________, thus ___________________________ the value of K w.
o As Kw increases, so do the concentrations of H+(aq) and OH-(aq)  pH decreases
o However, since hydronium and hydroxide concentrations remain equal, water does not become acidic or
basic as temperature changes, but the measure of its pH does change.
5
Unit 8: Acids & Bases
o
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Example: Fill in the rest of the table below
Table 8.4: Kw is temperature dependent

Temp (C)
Kw
[H+] in pure water
pH of pure water
0
10
20
25
30
40
50
1.5 x 10-15
3.0 x 10-15
6.8 x 10-15
1.0 x 10-14
1.5 x 10-14
3.0 x 10-14
5.5 x 10-14
0.39 x 10-7
7.47
0.82 x 10-7
7.08
1.22 x 10-7
1.73 x 10-7
6.92
6.77
H+ and OH- are inversely related
o Because the product [H+] x [OH-] is constant at a given temperature, it follows that as one goes up, the other
must go down (since Kw = [H+][OH-])
Table 8.5 Solutions are defined as acidic, basic, or neutral based on the relative concentrations of H+ and OH-
Type of sol’n
Relative concentrations
pH at 25C
Acid
Neutral
Alkaline
o
Example: A sample of blood at 25C has [H+]=4.60 x 10-8 mol dm-3. Calculate the concentration of OH- and
state whether the blood is acidic, neutral or basic.


How would you expect its pH to be altered at body temperature (37C)?
pH and pOH scales are inter-related
o pOH=




KW = [H+][OH-]
-log KW = -log([H+][OH-])
-log KW = (-log[H+]) + (-log[OH-])
pKW = pH + pOH
-14
at 25C, KW = 1.0 x10 , thus 14.00 = pH + pOH at 25C
Given any one of the following we can find the other three: [H+],[OH-],pH and pOH
Example: Lemon juice has a pH of 2.90 at 25C. Calculate its [H+],[OH-], and pOH.
From the relationship:
6
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
PART 3: Weak Acids & Bases
Strong acids and bases: pH and pOH can be deduced from their concentrations: since we assume strong acids and
bases dissociate completely, pH and pOH can be calculated directly from the initial concentration of solution.
Example: Calculate the pH of a 0.10 M solution of NaOH at 298 K.
Example: Calculate the pH of a 0.15 M solution of HNO3 at 298 K.
Dissociation constants express the strength of weak acids and bases
Since equilibrium for weak acids and bases lies far to the left (they do not dissociate fully), concentrations of ions in
solution cannot be determined by the initial concentrations without knowing the extent of dissociation.
Consider the equilibrium expression for the dissociation of any weak acid in water:
o
o
o
Ka is known as the acid dissociation constant.
It has a fixed value for a particular acid at a specified temperature.
Since the value of Ka depends on the position of equilibrium of acid dissociation, it gives us a direct measure of the
strength of an acid.
o The higher the value of Ka at a particular temperature, the greater the dissociation and so the stronger the acid.
o Note: because Ka is an equilibrium constant, its value does not change with the concentration of the acid or in the
presence of other ions.
Consider the equilibrium expression for the dissociation of any weak base in water:
Kb is known as the base dissociation constant. It has the same characteristics as those described above for Ka.
Calculations involving Ka and Kb
The values of Ka and Kb enable us to compare the strengths of weak acids and bases and to calculate ion concentrations
present at equilibrium (and therefore the pH and pOH values). Keep the following in mind:
o The given concentration of an acid or base is its initial concentration (before dissociation occurs).
o The pH (or pOH) of a solution refers to the concentration of H+ ions (or OH-ions) at equilibrium.
o The concentration values substituted into the expressions for Ka and Kb must be the equilibrium values for all
reactants and products.
o When the extent of dissociation is very small (very low value for Ka or Kb) it is appropriate you use the
approximations [acid]initial  [acid]equilibrium and [base]initial  [base]equilibrium.
7
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Calculation of Ka and Kb from pH and initial concentration
Example: Calculate Ka at 25C for a 0.0100 mol dm-3 solution of ethanoic acid, CH3COOH. It has a pH value of 3.40 at this
temperature.
Example: Calculate Kb at 25C for a 0.100 mol dm-3 solution of methylamine, CH3NH2. Its pH value is 11.80 at this
temperature.
Calculation of [H+] and pH, [OH-] and pOH from Ka and Kb
A real, but ugly example: Calculate the pH of a 0.10M solution of HNO2 (Ka = 4.0 x 10-4)
Step 1: Write out the dissolving equation & the equilibrium law expression
Step 2: Set up ICE and let x = [H+]
Step 3: Substitute equilibrium values into the
equilibrium law expression.
Step 4: Solve the problem using the assumed values.
Step 4¼: Check assumption for validity (5% rule)
8
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Step 4½: Solve the problem again using quadratic since the assumed values made an __ __ __ out of U and ME.
Step 5: Solve for pH
NOTE: Due to the nature of the timed exams you will be taking (IB and/or AP), assumptions of weak acids and bases
dissociating less than 5% will always be considered valid. ASSUME away and do not use the quadratic formula. You
must state your assumption, but you need not check on the validity of said assumption.
Pretty little AP/IB example: Determine the pH of a 0.75 mol dm-3 solution of ethanoic acid (Ka = 1.8 x 10-5).
Another pretty little AP/IB example: Determine the pH of a 0.20 mol dm-3 solution of ammonia (Kb = 1.8 x 10-5).
9
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
pKa and pKb:
1) pKa and pKb numbers are usually positive and have no units
2) The relationship between Ka and pKa and between Kb and pKb is inverse. Stronger acids or bases with higher
values for Ka or Kb have lower values for pKa or pKb.
3) A change of one unit in pKa or pKb represents a 10-fold change in the value of Ka or Kb.
4) pKa and pKb must be quoted at a specified temperature (they are derived from temp-dependent Ka and Kb).
Relationship between Ka and Kb, pKa and pKb for a conjugate pair:
Consider the Ka and Kb expressions for a conjugate acid-base pair HA and A-.
Therefore, Ka  Kb = [H+] [OH-] = Kw
By taking the negative logarithms of both sides: pKa + pKb = Kw
At 25C Kw = 1.0 x 10-14, so pKw = 14.00 
therefore pKa + pKb = 14.00 (at 25C)
Thus, the higher the Ka for the acid, the lower the value of Kb for its
conjugate base. In other words, stronger acids have weaker conjugate
bases and vice versa.
Practice Problems (Parts 1-3):
1)
2)
3)
4)
5)
6)
7)
Write the dissociation rxn and corresponding Ka equilibrium expression for the following acids in water: a) HCN and b) HF
Identify the acid, base, conjugate base and conjugate acid:
a) C5H5NH+ + H2O  C5H5N + H3O+
b) Al(H2O)63+ + H2O  H3O+ + Al(H2O)5(OH)2+
List the 6 strong acids.
List the four strong bases you should know.
List the weak acids you should know.
List the weak bases you should know.
Fill in the chart below:
pH
pOH
[H+]
[OH-]
Acidic, alkaline, or neutral?
5.02
3.9E-9 M
0.27 M
8.5
8) Calculate the Kb of ethylamine, C2H5NH2 given that a 0.10 mol dm-3 solution has a pH of 11.86.
9) What are the [H+] and [OH-] in a 0.10 mol dm-3 solution of an acid that has a Ka = 1.0 x 10-7?
10) The pKa of ethanoic acid, CH3COOH, at 25 C is 4.76. What is the pKb of its conjugate base, CH3COO-?
10
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
PART 4: Salt Hydrolysis & Buffer Solutions
BUFFER SOLUTIONS: solutions that resists a change in pH on the addition of small amounts of acid or alkali.
 Two types:
o acidic buffers – maintain a pH value _______________________
o basic buffers – maintain a pH value _______________________
 Mixture of two solutions, combined in such a way that they each contain the two species of a conjugate acidbase pair.
1) Acidic Buffers
Composition of the buffer solution:
Made by mixing an aqueous solution of a weak acid with a solution of its salt of a strong alkali. For example:
The following equilibria exist in a solution of this mixture:
(weak acid – eq’m lies to the left)
(soluble salt, fully dissociated in solution)
So the mixture contains relatively high concentrations of both CH3COOH and CH3COO-. Note that this is an acid
and its conjugate base. These can be considered “reservoirs,” ready to react with added OH- and H+ respectively
in neutralization reactions.
Response to added acid and base:
Addition of acid (H+): H+ combines with the base (CH3COO-) to form CH3COOH, removing most of the added H+.
Addition of base (OH-): OH- combines with the acid (CH3COOH) to form CH3COO- and H2O, thereby removing
most of the added OH-.
Consequently, as the added H+ and OH- are used in these rxns, they do not persist in solution and so pH does
not change significantly.
2) Basic Buffers
Composition of the buffer solution:
Made by mixing an aqueous solution of a weak base with a solution of its salt of a strong acid. For example:
The following equilibria exist in solution:
(weak base – eq’m lies to the left)
(soluble salt, fully dissociated in solution)
11
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
So the mixture contains relatively high concentrations of both NH3 and NH4+. Note that this is a base and its
conjugate acid. These act as “reservoirs,” ready to react with added H+ and OH- respectively in neutralization
reactions.
Response to added acid and base:
Addition of acid (H+): H+ combines with the base (NH3) to form NH3+, removing most of the added H+.
Addition of base (OH-): OH- combines with acid (NH4+) to form NH3 and H2O, removing most of the added OH-.
So, as with the acidic buffer, the removal of the added H+ and OH- by reactions with components of the buffer
solution keeps the pH of the buffered solution relatively constant. (For further proof that buffers resist pH
change: see Sample Exercise 15.3, p. 685)
Determining the pH of a buffer solution
Consider an acidic buffer made of the generic weak acid HA and its salt MA.
We can make two approximations based on some assumptions about these reactions:
1. The dissociation of the weak acid is so small that it can be considered negligible  thus [HA]initial  [HA]eq’m
2. The salt is considered to be fully dissociated into its ions  thus [MA]initial  [A-]eq’m
Henderson-Hasselbalch Equation(s):
pH =
pOH =
Example: Calculate the pH of a buffer solution at 298 K, prepared by mixing 25 cm3 of 0.10 mol dm-3 ethanoic acid,
CH3COOH, with 25 cm3 of 0.10 mol dm-3 sodium ethanoate, NaCH3COO. Ka of CH3COOH = 1.8 x 10-5 at 298 K.
Note: when a buffer solution contains equal amounts in moles of acid and salt (or base and salt), the last term in the
Henderson-Hasselbalch expression becomes 0, so pH = pKa (or pOH = pKb).
Making buffer solutions
pH of a buffer depends on:
1. The pKa (or pKb) of its acid or base
2. The ratio of the initial concentration of acid and salt (or base and salt) used in its preparation
12
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Acidic buffer: React a weak acid (whose pKa value is as close as possible to the desired pH of the buffer) with enough
strong alkali such that one half of it is converted into salt.
For example:
CH3COOH(aq) + NaOH(aq)  NaCH3COO(aq) + H2O(l)
Initial amt. (mol)
Change (mol)
Eq’m (mol)
Basic buffer: React a weak base (whose pKb value is as close as possible to the desired pOH of the buffer) with enough
strong acid such that one half of it is converted into salt.
Example: How would you prepare a buffer solution of pH 3.75? (See IB Data Booklet)
Factors that can influence buffers
 _______________: does not change pH (Ka & Kb are constant & ratio of acid or base to salt does not change), but
does decrease buffering capacity (amt. of acid or base the sol’n can absorb without a significant change in pH)
 _______________: As temp. affects the values of Ka and Kb, the pH of the buffer is affected accordingly.
SALT HYDROLYSIS
Neutralization rxn: acid + base  salt + water (pH of resulting solution depends on relative strengths of acid and base)
Anion hydrolysis: The anion (A-) is a conj. base of the parent acid. When the acid is weak, the conj. base is strong enough
to hydrolyze water (and the release of OH- causes the pH of the sol’n to ↑):
Cation hydrolysis: The cation (M+) is a conj. acid of the parent base. When the base is weak and this conj. is a non-metal
(i.e. NH4+), the conj. acid is strong enough to hydrolyze water (and the release of H+ causes the pH of the sol’n to ↓).
When the cation is a metal, the outcome depends on charge density. Metal ions that are either small with a +2 or +3
charge, or are transition metals (i.e. Fe3+) hydrolyze water (attracting OH-; releasing H+ into sol’n). Group 1 cations and
Group 2 cations below Be on the P.T. do not have sufficient charge density to hydrolyze water.
Neutralization Rxn
strong acid and
strong base
weak acid and
strong base
strong acid and
weak base
weak acid and
weak base
Ex: parent acid & base
HCl + NaOH
Salt formed
Hydrolysis of ions
neither ion
hydrolyzes
CH3COOH + NaOH
anion hydrolysis
HCl + NH3
HCl + Al(OH)3
cation hydrolysis
CH3COOH + NH3
anion and cation
hydrolyze
Type of salt sol’n
pH of salt sol’n
Depends on relative
strengths of conjugates
cannot
generalize
13
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
PART 5: Titration Curves & Indicators
ACID-BASE TITRATIONS



Units


Commonly used to determine the amount of acid or base in a solution.
Complete by adding a solution of known concentration until the substance being tested is consumed.
 This is called the _____________________ point (a.k.a.__________________________ point)
Titration curve - graph of pH vs. volume of added acid/base
millimole (mmol) = 1/1000 mol = 10-3 mol
Molarity = mol/L = mmol/mL  This makes calculations easier because we will rarely add liters of solution.
Strong acid & Strong Base: Do the stoichiometry. There is no equilibrium. They both dissociate completely.
Weak acid & Strong base: There is an equilibrium. Do stoichiometry. Then do equilibrium (ICE or H-H).
Titration Curves
1) Strong acid and strong base
2) Weak acid and strong base
3) Weak base and strong acid
4) Weak acid and weak base
Titration Curves: Weak Acid and Strong Base
(Comparing Weak Acids)
1. The solution of a weak acid has a higher initial pH than a
solution of a strong acid of the same concentration
2. The weaker the acid the more rapidly the pH rises in the early
part of the titration, but more slowly near the equivalence point.
This is because of the buffering action of the formed salt.
3. The pH at the equivalence point is not 7.00. The weaker the
acid, the higher the pH at the EQ point.
14
Indicators must be chosen based on the pH of the
equivalence point
Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Calculations involving Titrations of Weak Acids/Bases
Divide the titration into the following points:
1.
2.
 Use a two step procedure when calculating pH prior to the equivalence point whenever a strong acid/base is added
to a weak acid/base.
o 1. Stoichiometric Calculation: Allow the strong acid/base to react to completion with the weak acid/base,
producing a solution that contains the weak acid and its common ion. (its conjugate base/acid)
o 2. Equilibrium Calculation: Use the value of Ka/Kb and the equilibrium expression to calculate the
equilibrium concentrations of the weak acid/base, and the common ion, and H+ (ICE notation or HH Eq’n)
3.
4.
Example: Titration of a weak acid with a strong base
Calculate the pH when the following quantities of 0.050M KOH solution have been added to 50.0 ml of a 0.025M
solution of benzoic acid (HC7H5O2 Ka = 6.5x10-5).
A) 20.0ml B) 25.0ml C) 30.0ml
Step 0: Find the EQ point
Step 1: Starting point
This is just a “weak acid in water” problem…..
Step 2: Before the EQ point (20.0 ml KOH added)
First perform stoichiometric calculations…
Calculate moles of OH- present…
Calculate moles HC7H5O2 present…
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Unit 8: Acids & Bases
IB Topics 8 & 18
AP Chapters (Zumdahl): 4.8; 14(all); 15(all)
Next, using these values, perform equilibrium calculations…
Step 3: At the Eq point (25.0 ml KOH added)
Step 4: Beyond the Eq point (30.0 ml KOH added)
Beyond the equivalence point, there is no longer any benzoic acid to neutralize the additional KOH added. The pH of the
solution will be determined by the amount of excess KOH (the effect of the salt is negligible compared to the KOH)
We have used three different methods of calculating the pH during a titration….
1. Before the equivalence point has been reached. (Weak acid/base + common ion, buffered region)
2. At the equivalence point. (hydrolysis of a salt)
3. Beyond the equivalence point. (Strong Acid/Base in water)
INDICATORS
Signal change in pH; change color when the pH = their pKa; can be used to signal the equivalence point in titrations
An indicator is a weak acid (or base) in which the dissociated form is a different color than the undissociated form.
Color changes when [I-] = [HIn]. Thus, indicator changes color when Kin = [H+] …or when pKin = pH
Indicator
pKin
pH range
Use
methyl orange
3.7
3.1-4.4
Titrations with strong acids
phenolphthalein
9.6
8.3-10.0
Titrations with strong bases
16