SCH4U1
UNIT 4 – Acid-Base Equilibrium (Chapter 8)
8.1 The Nature of Acid-Base Equilibria
Bronsted-Lowry Theory
*In most acid-base interactions, a ___________ (H+ ion) is transferred from one reactant to another.
*Bronsted-Lowry acid = proton ____________
*Bronsted-Lowry base = proton ____________
Ex.1
H2O(l)
_____
+
HCl(g)
_____
=====>
H3O+(aq)
conjugate______
+
Cl-(aq) *Conjugate acid-base pair #1 =
conjugate- *Conjugate acid-base pair #2 =
______
*HCl(g) is a ___________ acid - it ionizes >>__________% in water to form H3O+ ions. A ________ arrow
in the equation is used to indicate this.
**Remember: The reaction of a strong acid with water is essentially ______% complete. (i.e. the
___________ reactions of these acids are strongly favoured.)
**The stronger an acid, the _________ its conjugate base. *The weaker an acid, the _________ its
conjugate base.
**Weak acids have _____ percent reactions (usually less than 3%).
Strong Acids
Weak Acids
Weak vs. Strong Acids  http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/acid13.swf
*Other strong acids include:
Strong Acids
Name
List and name some other weak oxyacids:
monoprotic
monoprotic
monoprotic
monoprotic
monoprotic
diprotic
*Illustrate the “diprotic nature” of sulfuric acid:
1
*Draw Lewis structure for nitric, perchloric and sulfuric acids:
Ex. 2
NH3(g) + H2O(l) <===> NH4+(aq) + OH-(aq)
base
______
conjugateconjugate______
______
*Conjugate acid-base pair #1 =_______/_______
*Conjugate acid-base pair #2 =_______/_______
*The __________ arrows in the equation indicates that NH3(g) is a weak _______ – it ionizes ___________
and forms a dynamic equ’m with the products in the reaction.
*Notice from the above examples that H2O is ______________; that is, it may ___________ (act as a
base) or ___________ (act as an acid) a proton depending on the chemical reaction.
Practice: Represent the amphoteric nature of the hydrogen carbonate ion (HCO3-) using two separate
equations.
Strong Bases  an ionic substance that (according to the Arrhenius definition) dissociates _________ in
water to release _____________ ions. All hydroxides of Group ____ and Group ____ elements are strong
bases (i.e. dissociate _______________).
*Examples:
*The Autoionization of Water
 the reaction between two water molecules producing a ____________ ion and a __________ ion
i.e.
H2O(l) + H2O(l) <===> H3O+(aq) + OH-(aq)
*Simplification:
(*Remember: H+ = H3O+)
*H2O + proton = ______________ ion ; H2O – proton = ____________ ion
 This reaction is extremely _____. On average, at 25C, only about _______ water molecules in a billion
are ionized at any given moment. At higher temperatures this number ______________.
2
The Ion-Product Constant of Water, Kw
The product of the molar concentrations of _______________ ion and _______________ ion in water is
a ___________________ called the ion-product constant of water, Kw.
*At SATP (i.e. 25˚ C / 101.3 kPa):
Kw = [
Kw = (
(aq)][
H2O(l)
H+(aq) + OH-(aq)
(aq)]
mol/L)(
mol/L)
Kw =
*How can the Kw formula be rearranged to calculate [H+(aq)] and [OH-(aq)]?
[H+(aq)] =
[OH-(aq)] =
The value of Kw is __________________________dependent.
Temperature (˚C)
Kw
0
1.5 x 10-15
10
3.0 x 10-15
20
6.8 x 10-15
*25
*1.0 x 10-14
30
1.5 x 10-14
40
3.0 x 10-14
50
5.5 x 10-14
60
9.5 x 10-14
-14
*Use 1.0 x 10 in calculations, unless otherwise stated.
In _______________solutions: [H+(aq)] = [OH-(aq)]
In _______________solutions : [H+(aq)] > [OH-(aq)]
In _______________solutions : [H+(aq)] < [OH-(aq)]
Important Formulas used to Determine: pH, pOH and pKw
*The pH concept was devised to make it easier to express and compare the concentration of H + ions when
they have very ___________ (i.e. < 1 mol H+/L) values involving awkward exponential terms.
pH =
pOH =
pKw = -logKw
and
and
pKw = -log(
pKw =
[H+(aq)]=
[OH-(aq)] =
pH + pOH = pKw
*at SATP
)
pH + pOH =
Q. What is the pH of neutral water at: a) 0˚C? b) 60˚C
*Try the following: p. 532#1-3 ; p. 537 #4-7 ; p. 540 # 8-10 ; p. 546 #12-15 ; p.549 #17-19
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Activity  The pH/pOH scale:
(@ 25˚C:)
[H+(aq)]
pH
Sample pH calculation:
[OH-(aq)]
pOH
Acidic,
Basic, or
Neutral?
[H+(aq)] x
[OH-(aq)]
pH +
pOH
Sample pOH calculation:
 What is the relationship between the concentrations of hydroxide and hydronium in
neutral, acidic and basic solutions, respectively?
 When given the pH, how can the concentration of hydronium be calculated?
 The pH scale is a logarithmic scale. What does this mean?
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8.2 Weak Acids and Bases
*Weak Acids  do not ionize completely in water (Most common acids are weak.)
Ex. 1 - A weak _________________ acid: acetic acid [CH3COOH   HC2H3O2 - *The
“acidic” hydrogen(s) is/are represented at the beginning of the chemical formula.]
Now W ritten
*The ionization of acetic acid:
*Abbreviation:
*The Acid Ionization Constant, Ka:
Ka =
Ex. 2 - A weak __________________ acid: carbonic acid (H2CO3)
Ionization #1:
Ionization #2:
Ka1 =
Ka2 =
=
=
Ka1 > Ka2
Ex. 3 - A weak ___________________ acid: phosphoric acid (H3PO4)
Ionization #1:
Ionization #2:
Ionization #3:
Ka1 =
Ka2 =
Ka3 =
=
=
=
Ka1 > Ka2 > Ka3
*Weak Bases  react incompletely with water; have a weak attraction for protons
Example: aniline, C5H5NH2
The ionization equation:
*Can’t abbreviate.
*The Base Ionization Constant, Kb:
Kb =
5
*Alternate method by which to determine Kb:
*Show how this formula is derived.
Kb x K a = Kw
*Kw =
Therefore:
Kb =
Kw
* Ka
*Ka of the conjugate acid.
E.g., Use this formula to determine the Kb of ammonia.
Ionization equation:
Try #1-13 from the section 8.2.
8.3 Acid-Base Properties of Salt Solutions
Dissolving Animation  http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf
Salts:
 strong ________________ (i.e. conduct an ______________l current when ____________ or_______)
 composed of ___________ and ____________ in a rigid _______________ lattice
*Hydrolysis (“water splitting”)  a reaction of an ion with water to produce an ____________ or
____________ solution (i.e. hydronium or hydroxide ions)
*Some Guidelines to Remember
In general:
(i) Strong acids have very weak (i.e. do not react with water) conjugate _________.
(ii) Weak acids have weak (do react with water) conjugate __________.
(iii)Weak bases have weak (do react with water) conjugate __________.
(iv) Strong bases have very weak (do not react with water) conjugate _________.
1. Salts that form Neutral Solutions:
*In general, salts that consist of the cations of ___________ bases and the anions of _________ acids have
_____ effect on the pH of an aqueous solution.
e.g.
Salt
NaCl
KCl
NaI
NaNO3
Cation from Strong Base
Anion from Strong Acid
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*In general:
 Group ____and _____cations do not react with water.
 The conjugate bases of very strong _________(i.e. Cl-, I-, Br-, NO3-, ClO4- & SO42-) do not react with
water. (Exception: The conjugate base of sulfuric acid, HSO 4-, does react with water, but only as an acid
(i.e. it is not considered amphoteric.)
2. Salts that form Basic Solutions:
Example: Na2CO3(s)
Dissociation:
Na2CO3(s)
2O
H


*Na+(aq) does not react with water. (i.e., does not affect the pH of the solution.)
*CO32-(aq) , however, is the conjugate base of a weak acid and will react with water.
CO32-(aq) + H2O(l) 
Base Hydrolysis:
Summary:
CO32-(aq) is
Kb =
a weak base and will therefore _______________ the pH.
3. Salts that form Acidic Solutions:
Example: NH4Cl(s)
Dissociation:
NH4Cl(s)
2O
H


*Cl-(aq) does not react with water. (i.e., does not affect the pH of the solution.)
*NH4+(aq) , however, is the conjugate acid of a weak base (NH3) and will react with water.
NH4+(aq) + H2O(l) 
Acid Hydrolysis:
Ka =
Summary: NH4+(aq) is a very weak acid and will ____________ the pH.
4. Metal Salts that form Acidic Solutions:
*_________ charged__________metal ions such as Al3+(aq), Fe3+(aq), and Sn2+(aq) form _________solutions.
Example: Al(NO3)3(s)
Dissociation:
Al(NO3)3(s)
2O
H


Hydration (formation of hydrated complex ion):
Acid hydrolysis:
*NO3- does affect pH.
Al3+(aq) + 6H2O(l) 
Al(H2O)63+(aq) + H2O(l) 
Summary: Al3+(aq) __________ the pH.
5. Salts that act as Acids and Bases:
Example: NH4CN(s)
Dissociation:
NH4CN(s)
2O
H


7
Hydrolysis (2):
#1(acid hydrolysis):
NH4+(aq) + H2O(l) 
Ka =
CN-(aq) + H2O(l) 
Kb =
and
#2(base hydrolysis):
*Net effect?  Acidic or Basic?
Kb _____ Ka, therefore, solution will be __________. (i.e. pH >7)
6. Hydrolysis of Amphoteric Ions:
Example: NaHCO3(s)
Dissociation:
NaHCO3(s)
2O
H


Hydrolysis(2):
#1(acid hydrolysis):
HCO3-(aq) + H2O(l) 
Ka =
HCO3-(aq) + H2O(l) 
Kb =
and
#2(base hydrolysis):
*Net Effect  Acidic or Basic?
Kb _______ Ka, therefore, solution will be ___________ (i.e. pH > 7)
Try: pp. 588-589 #1-6, 8
8.4 Acid-Base Titration
Titration
A titration is a controlled addition and measurement of the amount of a
solution of known concentration (mol dm-3) required to react completely
with a measured amount of a solution of unknown concentration (mol dm-3).
In a titration we are trying to
reach the _______________
point. This is the point at
which the two solutions are
present in exactly the same
amounts.(moles)
In a titration we use
___________solutions to show
us when the _____-point has
been reached. The end point is
the point at which the indicator
changes ______. It is
important that we pick an
indicator that will change color
in a range that includes the pH of the equivalence point.
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Titration Calculations: Moles acid = Moles base; if the mole ratio is a 1:1 acid : base
ratio, otherwise, we must take into account the mole ratio
Summary of Terms:
*titration  the precise addition of solution in a ________ into a _________ volume of a sample solution
*titrant  the solution in a buret during a titration; concentration is _____________
 The titrant is added to the _________ until it is ____________. The volume of titrant needed to
neutralize the sample is recorded and used to calculate the ______________ of the sample.
*sample  the solution being analyzed in a titration; _________ is known; _______________ is unknown
*endpoint  the point in a titration at which a sharp change in a measurable and characteristic property
occurs [e.g. a ___________ change in an acid-base ______________ / e.g. phenolphthalein indicator is
___________in acid solution and _____ in basic solution – when it turns pink the __________ has been
reached] *Indicators are usually __________ acids themselves.
*equivalence point  the point in the titration where the amount of _________ added reacts with the
___________ amount of reactant in the ___________
*When performing a titration it is important to choose an indicator that has an endpoint that is equal (or
very close to) the equivalence point.
Titration Type #1  Strong Acid with Strong Base
*Equivalence point: pH = _____ (i.e. the titration is complete when the pH of the resulting solution
equals ______. Therefore, the chosen indicator should turn colour (i.e. reach its endpoint) at a pH of ____.
Ex. 1: A sample of hydrochloric acid needs to be analyzed; that is, the concentration of the sample needs to
be determined. A 10.0 mL sample of the acid is titrated with 0.10 M sodium hydroxide. 13.5 mL of titrant
(i.e. NaOH) are added before the endpoint is reached. Determine the concentration of the acid.
*Always start with a balanced chemical equation: *How do the products of the reaction affect the pH?
Ex. 2: A 25.0 mL sample of calcium hydroxide is titrated with 0.22 M hydrobromic acid. It takes 18.3 mL
of titrant to neutralize the sample. Determine the concentration of the sample.
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Titration Type #2  Weak Acid with Strong Base
*Equivalence point: pH ___ 7.0
(Why?...How do the products of the neutralization reaction affect the pH?)
Ex.: It is found that 42.5 mL of 1.02 M NaOH have been added to 50.0 mL of vinegar when
phenolphthalein (endpoint occurs at pH of ~ _____) in the solution turns pink (i.e. equivalence point is
reached). What is the concentration of the vinegar?
Titration Type #3  Strong Acid with Weak Base
*Equivalence point: pH___ 7.0
(Why?...How do the products of the neutralization reaction affect the pH?)
Ex. Write the equation for a neutralization reaction involving ammonia (a weak base) and hydrochloric
acid.
Summary:
“Type” of Titration
pH at equivalence point
Entity determining pH at
equivalence point
<7
7
>7
Titration Calculations (More Practice):
1. In a titration, 27.4 ml of 0.0154M Ba(OH)2 is added to a 20.0 ml sample of HCl solution of unknown
concentration. What is the molarity of the acid solution?
2. A 15.5 ml sample of 0.215M KOH solution required 21.2 ml of acetic acid solution to reach the
equivalence point in an experiment. Calculate the molarity of the acetic acid solution.
3. If 20.0 ml of 0.0100M aqueous HCl is required to neutralize 30.0 ml of an aqueous solution of NaOH,
determine the molarity of the base.
4. Suppose that 20.0 ml of 0.10M Ca(OH)2 is required to neutralize12.0 ml of aqueous HCL solution.
What is the molarity of the HCl solution?
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Titration Curves:
Strong Acid / __________
______________/_____________
_______________/_______________
*Choosing an appropriate indicator:
*At the equivalence point:
pH =
8.5 Buffers: The Control of pH
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf
*If a small volume of a strong acid or base is added to water, the pH of the water will change significantly.
e.g. adding  2 drops of 1.0 M HCl(aq) to a litre of water will change the pH from ____ to ____.
*When a solution is __________, its solutes protect it against large changes in pH even when _______
volumes of ___________ acids or bases are added. (i.e. they buffer the system against a pH)
*Every life form is extremely sensitive to slight changes in pH. The pH of blood is _______. If it were to
change to 7.00 or 8.00, for example, you would _________. Blood is a ___________ solution.
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Components of a Buffer:
Usually a buffer consists of two solutes:
1. a ________ Bronsted _______ (HA(aq)); and,
2. the weak acid’s conjugate ___________. (supplied by a soluble _________ of the acid such as NaA(aq))
e.g.
*The capacity of a buffer is determined by the _______________ of its acid-base
conjugate pair.
Equilibrium of a Buffer:
Generic equation: HA(aq)  H+(aq) + A-(aq)
Specific example - acetic acid / sodium acetate buffer:
*If H+ is added, there will be a shift to the left.
The net change is:
Generic equation: H+(aq) + A-(aq)  HA(aq)
Ex.
*If OH-(aq) is added, there will be a shift to the right
The net change is:
Generic equation: HA(aq) + OH-(aq)  A-(aq) + H2O(l)
Ex.
The Carbonate Ion in Blood:
*Your body uses H2CO3(aq) and its conjugate base, HCO3-(aq) (i.e. the bicarbonate ion)
CO2(aq) + H2O(l)  H2CO3(aq)  HCO3-(aq) + H+(aq)
Neutralizing OH-(aq):
H2CO3(aq) + OH-(aq) 
Neutralizing H+(aq)
HCO3-(aq) + H+(aq) 
**The Buffer equation:
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