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LECSS Chemistry 12 Unit 4 Topic A Notes

Revised October 23, 2013 © Don Bloomfield

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Unit 4: Acids and Bases

Topic A: Definitions of Acids and Bases and the Relative Strength of Acids and Bases

In this topic we will examine:

Various definitions of acids and bases

Brønsted-Lowry definitions

The meaning of pH (an introduction)

Conjugate acid-base pairs

Relative strength of Brønsted-Lowry acids and bases

The “levelling effect”

Hydrolysis of salts (part 1)

A.1: Definitions of Acids and Bases

Aqueous acids and bases may be defined in a number of ways. Some definitions are “more inclusive”

than others.

A.1.1: The Operational Definitions of Acids and Bases

The operational definition of acids and bases defines these substances according to observed chemical

and physical properties.

According to this definition, acids:

Taste sour. Example: citric acid which is in lemon juice

Conduct electric currents. That is, they are good electrolytes.

Cause certain dyes to change colour in a specific way. Example: they cause blue

litmus to turn red.

Produce H2(g) when reacted with certain metals. Example: the reaction between

aqueous HCl and solid Zn.

React with bases in such a way that all the above properties are lost, except for the

ability to conduct an electric current.

In contrast to this, bases:

Taste bitter.

Feel slippery. Example: soap.

Conduct electric currents. That is, they are good electrolytes.

Cause certain dyes to change colour in a specific way. Example: they cause red

litmus to turn blue.

React with acids in such a way that all the above properties are lost, except for the

ability to conduct an electric current.

While these definitions give a useful method of identifying acids and bases, they do not provide us

with an explanation of their behaviour.



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A.1.2: The Arrhenius Definitions of Acids and Bases

Svante Arrhenius defined acids and bases (and salts) on the molecular level.

According to the Arrhenius definitions:

Acids release H+ ions in aqueous solutions.

Bases release OH- ions in aqueous solutions.

An example of an Arrhenius acid is hydrochloric acid, HCl:

HCl(aq) H+

(aq) + Cl-(aq)

An example of an Arrhenius base is sodium hydroxide, NaOH:

NaOH(aq) Na+

(aq) + OH-(aq)

The Arrhenius definitions of acids and bases are able to explain why a reaction between them causes

them to lose their physical properties (except for the ability to conduct an electric current). Consider

the reaction between hydrochloric acid and sodium hydroxide:

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

The result of the reaction is an aqueous solution of sodium chloride. Because such a solution contains

a dissolved ionic solid (or salt), it is able to conduct and electric current. In general, by the Arrhenius

definitions:

ACID + BASE SALT + WATER

Because neither salts nor water fall under the Arrhenius definition of acids or bases, they are classified

as neutral substances. As a result, the reaction between an acid and a base is often referred to as a

neutralization reaction. Another example is the reaction between hydrofluoric acid and barium

hydroxide:

2HF(aq) + Ba(OH)2(aq) BaF2(aq) + 2H2O(l)

acid + base salt + water

In very general terms, the Arrhenius definitions define any compound having a chemical formula

starting with one or more “hydrogens” as an acid (e.g. HF, H2SO4, etc.).

Similarly, any compound having a chemical formula ending with one more “hydroxides” is a base

(e.g. LiOH, Ba(OH)2, etc.). The Arrhenius definitions also tell us that any ionic compound made up

of ions other than H+ or OH

- ions is neither an acid nor a base, but a salt.

Eliminating the spectator ions in either of the neutralization reactions above gives the following net

ionic equation:

H+

(aq) + OH-(aq) H2O(l)



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While Arrhenius’ definitions help to explain some of the chemical behaviour of acids and bases, it

does not explain why solutions of a large number of compounds (such as NH3 and NaF) are able to

partially neutralize acids such as HCl. Similarly, the Arrhenius definitions are unable to explain why

some compounds (such as NH4NO3) are able to partially neutralize bases such as NaOH.

Hebden Reference: Section IV.1

Practice Exercises:

Complete the following exercises from the Hebden text:

Pages 110 to 112 Exercises 1 to 3.

A.1.2.1: Some Common Acids and Bases

In Section IV.2, Hebden outlines some common acids and bases, their uses, and some alternate names

for them. These include:

sulphuric acid (battery acid)

hydrochloric acid (muriatic acid)

nitric acid

acetic acid (ethanoic acid, vinegar)

sodium hydroxide (caustic soda, lye)

potassium hydroxide (caustic potash)

aqueous ammonia (a common household cleaner)

You will need to be familiar with the acids and bases listed above and some of their uses.

In addition to those above, you also need to be familiar with:

carbonic acid (carbonated water)

citric acid (found in oranges, lemons, etc.)

ascorbic acid (vitamin C)

acetylsalicylic acid (A.S.A; found in pain relievers)

Perhaps unknown to Hebden is that it is actually possible to obtain sulphuric acid in concentrations

greater than 100%!

SO3(g) will dissolve in 98% H2SO4 solutions. When added to water, the dissolved SO3(g) combines

with water:

SO3(g) + H2O(l) H2SO4(aq)

Concentrated H2SO4 that has SO3(g) dissolved into it is called oleum.

If time permits, I will tell you why some frozen french fries are a good source of vitamin C.



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In this section, Hebden discusses the fact that NaOH (and KOH) rapidly absorb water from the air.

He uses the term deliquescent to describe such substances. A better term for this property of some

substances is hygroscopic. You need to know this term.

Notice that acetic acid (CH3COOH(aq)) does not “fit in” with the Arrhenius definition of an acid; its

chemical formula does not start with an “H”. While commonly considered to be an acid, acetic acid

solutions are rather poor electrolytes. This too, goes against the definition of an acid as we have

described thus far.

On the other hand, aqueous ammonia (NH3(aq)) does not “fit in” with the Arrhenius definition of a

base. Ammonia solutions may be quite basic, yet the chemical formula for ammonia does not end

with an “OH”. As well, it is a poor electrolyte, meaning that the concentration of dissolved ions is

rather low.

In the next section, we will examine a different way of defining acids and bases. This method of

defining acids and bases is much more inclusive and explains the acid-properties of substances such

as acetic acid and the base properties of substances such as ammonia.

Hebden Reference: Section IV.2

Practice Exercises:


Page 114 Exercises 5 to 9.

A.1.3: The Brønsted-Lowry Definitions of Acids and Bases

Johannes Brønsted and Thomas Lowry simultaneously presented a more general method of defining

acids and bases.

A major flaw of Arrhenius’ definitions is that H+ ions do not independently exist in aqueous solutions.

Once released into a solution, they bond with other molecules that are present. Very often these are

water molecules.

For example, in a solution of HCl:

HCl(aq) + H2O(l) H3O+

(aq) + Cl-(aq) (i)

H3O+ is called the HYDRONIUM ion, or a hydrated proton.

Production of H3O+ is the result of transferring a proton (H

+) from the HCl to H2O molecules.

H+

(aq) + H2O(l) H3O+

(aq)

In equation (i), HCl is said to be “donating a proton” to another chemical substance.



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According to the Brønsted-Lowry Definition an acid is any substance that donates a proton in an

aqueous solution.

According to the Brønsted-Lowry Definition, a base is any substance that accepts a proton in aqueous

solution.

In equation (i), H2O is acting as a Brønsted-Lowry (B-L) base.

The B-L definitions extend to a number of substances that the Arrhenius definitions do not define as

either an acid or a base.

For example, aqueous ammonia is a B-L base because NH3 accepts protons from water molecules:

NH3(aq) + H2O(l) ⇌ NH4+

(aq) + OH-(aq) (ii)

Notice that in equation (i), H2O acts as a B-L base, while in equation (ii) above, it is acting as a (B-L)

acid.

Substances that can act as either a B-L acid or a B-L base (depending on the conditions of the

solution) are said to be AMPHIPROTIC. As we will see, many substances are amphiprotic.

With the exception of H2O, all other amphiprotic substances that we will encounter are anions

containing a hydrogen atom. For example,

Some amphiprotic substances: HPO42-

, HS-, HCO3

-, etc.

To attract a positive proton, a negative charge is required. Water, being extremely polar, does not

need a net negative electric charge. Of course, a substance cannot donate protons if it does not

contain hydrogen.

Comparing equations (i) and (ii), you may have noticed that different kinds of “arrow” were used.

This is due to the fact that many aqueous solutions of acids and bases form equilibrium systems.

Some do not. We will address this later.

Looking at the reverse reaction in equation (ii), we see that in the reverse reaction a proton is being

donated to the hydroxide ion by the ammonium ion,

NH4+

(aq) + OH-(aq) ⇌ NH3aq) + H2O(l)

acid + base base + acid

In EVERY Brønsted-Lowry acid-base reaction, an acid and a base react to form another acid and

another base. We can write two general equations for the reaction between a B-L acid and a B-L

base.

Either:

ACID + BASE ACID + BASE

or,

ACID + BASE ⇌ ACID + BASE



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The type of arrow used depends on whether the reaction goes to completion or establishes a state of

ionic equilibrium. This is determined by the nature of the acids and bases involved. We will study

this in-depth later.

A.1.3.1: A Different Definition of “Neutral”

Under the B-L definitions of acids and bases, “neutral” is defined differently than under the Arrhenius

definitions.

Definition of Neutral:

Any aqueous solution in which [H3O+] = [OH

-] is defined to be NEUTRAL.

In pure water at any temperature, [H3O+] = [OH

-]. Thus, pure water is ALWAYS neutral.

As we will see, the role temperature plays in determining the [H3O+] and [OH

-] in pure water is very

important.

With a different definition of neutral comes a new definition of acidic and basic.

Definitions of Acidic and Basic:

Any solution in which [H3O+] > [OH

-] is an ACIDIC solution.

Any solution in which [H3O+] < [OH

-] is a BASIC solution.

A.1.3.2: The pH Scale – An Introduction

Whether a solution is acidic, neutral, or basic (a.k.a. alkaline) is often indicated by a measure of its

pH.

For now, all that you need to know is that pH is a measure of the [H3O+] in a solution. Its exact

mathematical definition does not concern us just yet, but it will become a major part of this course in

due time.

Typically, measures of pH range from pH = 0 (very acidic) to pH = 14 (very basic). Basic solutions

are sometimes said to be “alkaline”.



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At 25C, and ONLY at 25C, a pH of 7.0 is considered neutral ([H3O+] = [OH

-]). For the purposes of

this course, you may always assume a temperature of 25C unless told otherwise.

Hebden Reference: Sections IV.3 and IV.4.

Practice Exercises:



Provincial Exam Examples: Definitions of Acids and Bases

1. Which of the following is a general characteristic of Arrhenius acids?

A. They produce H+ in solution. C. They accept an H

+ from water.

B. They turn bromthymol blue to a blue colour. D. They react with H3O+ ions to produce H2.

2. Which of the following is generally true of acids, but not for bases?

A. pH 7 C. release H+ in solution

B. conduct current when in solution D. cause indicators to change color.

3. Which of the following is a property of all acidic solutions at 25C?

A. They have a pH less than 7.0.

B. They have a pH greater than 7.0.

C. They cause phenolphthalein to turn pink.

D. They release hydrogen when placed on copper metal.

4. A substance that produces hydroxide ions in solution is a definition of which of the following?

A. an Arrhenius acid C. a Brønsted-Lowry acid

B. an Arrhenius base D. a Brønsted-Lowry base.

5. Which of the following represents the complete neutralization of H3PO4 by NaOH?

A. H3PO4 + NaOH NaH2PO4 + H2O

B. H3PO4 + 3NaOH Na3PO4 + 3H2O

C. H3PO4 + 2NaOH Na2HPO4 + 2H2O

D. H3PO4 + NaOH NaH + HPO4 + H2O



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6. Which of the following represents the neutralization reaction between Ca(OH)2(s) and HCl(aq)?

A. H2O(l) H+

(aq) + OH-(aq)

B. Ca2+

(aq) + 2Cl-(aq) CaCl2(aq)

C. Ca(OH)2(s) + 2HCl(aq) CaCl2(aq) + 2H2O(l)

D. Ca2+

(aq) + 2OH-(aq) + 2H

+(aq) + 2Cl

-(aq) CaCl2(aq) + 2H2O(l)

7. Which of the following household products could have a pH = 12.0?

A. soda pop C. lemon juice

B. tap water D. oven cleaner

8. Identify the common acid in the stomach.

A. nitric acid C. perchloric acid

B. sulphuric acid D. hydrochloric acid

9. Which of the following is a typical pH value for dishwashing solutions?

A. 2.0 C. 10.0

B. 4.0 D. 14.0

10. A Brønsted-Lowry acid is defined as a substance that

A. releases H+

(aq). C. accepts a proton.

B. releases OH-(aq). D. donates a proton.

11. What is a general characteristic of all Brønsted-Lowry bases?

A. They all accept H+.

B. They all accept OH-.

C. They will all turn litmus a pink color.

D. They will all react with acids to produce H2 gas.

12. Which of the following represents the reaction of H2PO4- acting as an acid?

A. H2PO4- + H2O ⇌ H3PO4 + OH

-

B. H2PO4- + H2O ⇌ H3O

+ + H3PO4

C. H2PO4- + H2O ⇌ H3O

+ + HPO4

2-

D. H2PO4- + H2O ⇌ H4PO4

+ + 2OH

-



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13. Consider the following equilibrium:

HS- + H3BO3 ⇌ H2BO3

- + H2S

The two species acting as Brønsted-Lowry bases in the equilibrium are

A. HS- and H2S C. H3BO3 and H2S

B. HS- and H2BO3

- D. H3BO3 and H2BO3

-

14. In which of the following is HSO3- acting as a Brønsted-Lowry acid?

A. HSO3- + H2O ⇌ H2SO3 + OH

-

B. NH3 + HSO3- ⇌ NH4

+ + SO3

2-

C. HSO3- + HPO4

2- ⇌ H2SO3 + PO43-

D. H2C2O4 + HSO3- ⇌ HC2O4

- + H2SO3

15. Write a chemical reaction showing an amphiprotic anion reacting as a base in water.

A.2: Conjugate Acid-Base Pairs and Relative Acid Strength

As mentioned previously, every Brønsted-Lowry acid-base reaction has an acid and a base on each

side of the reaction equation.

(iii) HNO3(aq) + H2O(l) H3O+

(aq) + NO3-(aq)

acid base acid base

(iv) HF(aq) + H2O(l) ⇌ H3O+

(aq) + F-(aq)

acid base acid base

(v) CO32-

(aq) + H2O(l) ⇌ HCO3

-(aq) + OH

-(aq)

base acid acid base

Each combination of 1 acid and 1 base with chemical formulae that differ by one proton constitute a

conjugate acid-base pair.

The conjugate pairs in the above reactions are:

(iii) _____________________and_________________________

(iv) _____________________and_________________________

(v) _____________________and_________________________



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It is VERY IMPORTANT that you remember that a conjugate pair differs by one, and only one,

proton.

e.g., the conjugate base of H2CO3 is HCO3- and NOT _______________.

Amphiprotic substances are tricky as they all have BOTH a conjugate acid AND a conjugate base.

e.g. H2O: conjugate acid: ____________ H2PO4- : conjugate acid: ____________

conjugate base: ____________ conjugate base: ___________

Notice that in equation (iii) a one-way arrow ( ) was used, while in (iv) and (v), a reversible arrow

(⇌ ) was used.

Use of the two different arrows indicates the degree to which each substance ionizes in an aqueous

solution.

HNO3 completely ionizes in an aqueous solution. Therefore, a solution of HNO3 contains only H3O+,

NO3-, and H2O molecules (i.e., no HNO3 molecules exist in their original form in solution).

Acids which completely ionize in aqueous solutions are called STRONG ACIDS.

There are 6 strong acids that you MUST know:

Perchloric acid HClO4

Hydroiodic acid HI

Hydrobromic acid HBr

Hydrochloric acid HCl

Nitric acid HNO3

Sulphuric acid H2SO4

Strong acids will react to completion with any Brønsted-Lowry base present in solution with it.

e.g., HCl(aq) + F-(aq) HF(aq) + Cl

-(aq)

acid base acid base

Acids that do not completely ionize in aqueous solutions are called WEAK ACIDS. In aqueous

solutions, weak acids form an equilibrium system.

e.g., HCN(aq) + H2O(l) ⇌ H3O+

(aq) + CN-(aq)

acid base acid base

Equilibrium systems such as the one above exhibit the same characteristics as all other equilibrium

systems we have encountered.



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It is important that one does not associate acid strength with concentration. They are distinctly

different things. We can have:

concentrated solutions of strong acids

dilute solutions of strong acids

concentrated solutions of weak acids

dilute solutions of weak acids

The Chemistry 12 Data Booklet contains a table of Relative Strengths of Brønsted-Lowry Acids and

Bases on page 6.

Notice that in each entry, the acid is shown to dissociate into a single H+ ion and the conjugate base of

the acid. e.g.,

HCOOH ⇌ H+ + COOH

-

This shows that the acid can donate a proton to a Brønsted-Lowry base when in solution. If the only

base present is H2O, then the reaction is:

HCOOH(aq) + H2O(l) ⇌ H3O+

(aq) + COOH-(aq)

The seventh entry from the top is:

H3O+ ⇌ H

+ + H2O

This may also be written as:

H3O+ + H2O ⇌ H3O

+ + H2O

This tells us that H3O+

(aq) and H+

(aq) are equivalent. The two-way arrow above means that the reaction

may be read left to right or right to left.

The H3O+ ions on the left side of the above equations represents the hydronium ions produced by the

ionization of any one of the six strong acids., e.g., HNO3(aq) + H2O(l) H3O+

(aq) + NO3-(aq).

Because all six strong acids completely ionize to form solutions of H3O+, all six strong acids have

equal strength when in an aqueous solution. The levelling effect is the name given to the effect water

has on the strong acids which causes them to all have the same strength. As we will see, the levelling

effect also applies to solutions of strong bases.

Another result of the complete ionization of the strong acids is that H3O+ is the strongest acid that can

exist in an aqueous solution.

As discussed at other times in this course, the ability of a solution to conduct an electric current

depends upon the concentration of dissolved ions in the solution. As strong acids (and strong bases)

are 100% ionized, they are good electrolytes. Solutions of weak acids (and weak bases) are only

partially ionized. This makes them weak electrolytes. As we move down the left side of the table, the

acids get weaker and weaker.



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On the right side of the table are the conjugate bases of the acids, with the weakest of the bases on the

top, the strongest on the bottom.

The six bases at the top are the conjugate bases of the strong acids. These are so weak that they will

not act as Brønsted-Lowry bases. The conjugate base of sulphuric acid (HSO4-) is, in fact, a rather

strong weak acid (see the eleventh entry from the top).

From the table we can conclude the following:

The stronger an acid is, the weaker is its conjugate base.

The three strongest bases listed are OH-, O

2- and NH2

-. Being anions, basic solutions containing these

anions must also contain cations (Why?). As we saw in Unit 3, an ion must be introduced into a

solution in the form of a soluble compound.

The OH-, O

2-, and NH2

- ions are only considered as strong bases when combined with one of seven

different metal ions: Li+, Na

+, K

+, Rb

+, Cs

+, Sr

2+ and Ba

2+.

The 21 Strong Bases

LiOH Li2O LiNH2

NaOH Na2O NaNH2

KOH K2O KNH2

RbOH Rb2O RbNH2

CsOH Cs2O CsNH2

Sr(OH)2 SrO Sr(NH2)2

Ba(OH)2 BaO Ba(NH2)2

You need to know these!! On page 122, Hebden states that Mg(OH)2, Ca(OH)2, Fe(OH)3 and

Zn(OH)2 are strong bases. He is absolutely incorrect.

In an aqueous solution, each of the amides and oxides listed above will first dissociate and then the

anion released will react with water. For example,

(i) NaNH2(s) Na+

(aq) + NH2-(aq) then NH2

-(aq) + H2O(l) NH3(aq) + OH

-(aq)

(ii) Na2O(s) 2Na+

(aq) + O2-

(aq) then O2-

(aq) + H2O(l) 2OH-(aq)

The hydroxides dissociate to release OH-,

NaOH(s) Na+

(aq) + OH-(aq)

Because the reactions of the NH2- and O

2- ions with water are not reversible, the strongest base that

exists in water is OH-. The reaction of the hydroxide ions from any strong base will accept a proton



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from any available acid. This reaction is given in the third entry from the bottom of the table when

read left to right:

H2O H+ + OH

-

If the only acid available is water, the reaction is:

OH- + H2O ⇌ H2O + OH

-

The two-way arrow in the above equation means that the reaction may be read in either direction.

In the presence of any acid, the OH- from a strong base will react to completion:

NaOH(aq) + HF(aq) H2O(l) + NaF(aq)

NOTE: It is common for students to think that NH3 is a strong base because of its location on the very

bottom of the left side of the table. This is absolutely incorrect. The presence of NH3 at this location

shows that it is the conjugate acid of NH2-. Mind you, it is such a weak acid that it will not act as an

acid at all, in the same way that the conjugate bases of the strong acids are too weak to ever act as

bases.

The ranking of the bases on the table lead to the following conclusion:

The stronger a base is, the weaker is its conjugate acid.

In a solution of a weak Brønsted-Lowry base, the base only partially reacts with water to form OH-

ions. For example:

CN-(aq) + H2O(l) ⇌ HCN(aq) + OH

-(aq)

The above equation represents an equilibrium system that behaves just as all other equilibrium

systems we have examined.

We can write four general equations to represent the reactions between Brønsted-Lowry acids and

bases with water.

Strong Acid: HA(aq) + H2O(l) H3O+

(aq) + A-(aq)

Weak Acid: HA(aq) + H2O(l) ⇌ H3O+

(aq) + A-(aq)

Strong Base: B-(aq) + H2O(l) HB(aq) + OH

-(aq)

Weak Base: B-(aq) + H2O(l) ⇌ HB(aq) + OH

-(aq)



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We have already discussed the fact that the strength of a Brønsted-Lowry acid is a measure of the

degree to which it will ionize in an aqueous solution. Since the result of the ionization of an acid is

the formation of H3O+, we can conclude that:

For any given concentration, the weaker an acid is, the lower the [H3O+].

Similarly:

For any given concentration, the weaker a base is, the lower the [OH-].

A.3: The Equilibrium Position of Brønsted-Lowry Acid/Base Reactions

We have discussed the idea that the products of every Brønsted-Lowry acid-base reaction are also a

Brønsted-Lowry acid and base.

During reaction of a strong acid and any base (weak or strong), the reaction will always go to

completion.

HA(aq) + B-(aq) HB(aq) + A

-(aq)

strong base acid base

acid

Similarly, the reaction of a strong base with any acid (weak or strong), the reaction will go to

completion.

HA(aq) + B-(aq) HB(aq) + A

-(aq)

acid strong acid base

base

In either case, the reaction reaches a final position of static equilibrium with only products present.

In the reaction between a weak acid and a weak base, the reaction will reach a state of dynamic

equilibrium, with reactants and products present in appreciable concentrations. Whether the products

are present in the greater concentration or the reactants are present in the greater concentration at

equilibrium depends on the relative strength of the acids and bases present in the equilibrium system.

Consider the general reaction between a weak acid and a weak base:

HA(aq) + B-(aq) ⇌ HB(aq) + A

-(aq)

acid base acid base

Both of the acids and both of the bases in the above equation are weak substances.



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If HA is the stronger of the two acids, then _______________ will be the weaker base.

If HB is the stronger of the two acids, then _______________ will be the weaker base.

The result of every weak acid/weak base reaction is that the stronger acid and stronger base always

appear on the same side of the reaction equation, and the weaker acid and weaker base always appear

on the same side of the equation.

In every weak acid/weak base reaction, the stronger of the weak substances will react to a greater

extent than will the weaker of the weak substances. We know this because the stronger a substance is,

the more closely it will react to completion.

In the equation above, if HA is the stronger of the two acids, B- will be the stronger of the two bases.

Since they react to a greater degree than do HB and A-, HB and A

- will be present in greater

concentrations than HA and B- at equilibrium. That is, the products will be “favoured at equilibrium”.

Conversely, if the stronger of the two weak acids and the stronger of the two weak bases are on the

product side of the reaction equation, they will react to the greater degree and the reactants will be

present in greater concentrations at equilibrium (i.e., the reactants will “be favoured”).

Determining which of the acids is stronger in the equilibrium resulting from a weak acid/weak base

reaction is simply a matter of locating each on the table. Usually there is no need to locate the weak

bases, as we know that the stronger of the weak acids has the weaker conjugate base.

Example:

Consider the following equilibrium:

HNO2(aq) + CN-(aq) ⇌ HCN(aq) + NO2

-(aq)

To determine whether the products or reactants are favoured at equilibrium, we simply need to locate

HNO2 and HCN on the table of relative acid strengths. We find the following:

HNO2(aq) + CN-(aq) ⇌ HCN(aq) + NO2

-(aq)

stronger weaker

This result also tells us that CN- is a stronger base than NO2

-. Since the reactants are the stronger of

the substances, the products are favoured at equilibrium.

In every acid-base reaction, the side of the reaction with the weaker acid and weaker base is

favoured at equilibrium.



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Practice Examples: Complete the following acid-base equilibrium reaction equations and state

whether the reactants or products are favoured at equilibrium.

(a) HIO3(aq) + SO32-

(aq)

(b) HCN(aq) + H2BO3-(aq)

(c) HF(aq) + HS-(aq)

Hebden Reference: Sections IV.5 and IV.6.

Practice Exercises:



Page 133 Exercises 38, 39, 43, and 44.

Provincial Exam Examples: Acid-Bases Pairs and Relative Acid Strength

1. Identify a conjugate pair from the following equilibrium:

PO43-

+ HCO3- ⇌ HPO4

2- + CO3

2-

A. CO32-

and PO43-

C. PO43-

and HPO42-

B. PO43-

and HCO3- D. HCO3

- and HPO4

2-

2. What is the conjugate base of H2PO4- ?

A. OH- C. HPO4

2-

B. PO43-

D. H3PO4

3. The conjugate base of HBO32-

is

A. BO32-

C. HBO3-

B. BO33-

D. H2BO3-

…more on next page



17

4. Consider the following reaction:

HCN + CH3NH2 ⇌ CN- + CH3NH3

+

Which of the following describes a conjugate acid-base pair in the equilibrium above?

Acid Base

A. CN-

HCN

B. CH3NH3+

CN-

C. HCN CH3NH3+

D. CH3NH3+ CH3NH2

5. a) Write the formula equation to represent the complete neutralization reaction between household

vinegar (acetic acid) and drain cleaner (sodium hydroxide).

b) Write the formula for the conjugate base of the reactant acid.

6. An acid-base reaction occurs between HSO3- and IO3

- .

a) Write the equation for the equilibrium that results.

b) Identify one conjugate acid-base pair in the reaction.

c) State whether reactants or products are favoured, and explain how you arrived at your answer.

__________________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________

7. The two reactants in an acid-base reaction are HNO2(aq) and HCO3-(aq).

a) Write the equation for the above reaction.



18

b) Define the term conjugate acid-base pair.

__________________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________

8. Which of the following is the weakest base?

A. F- C. CN

-

B. HS- D. IO3

-

9. List the bases C2O42-

, NH3, and PO43-

in order from strongest to weakest.

A. PO43-

NH3 C2O42-

B. C2O42-

NH3 PO43-

C. NH3 PO43-

C2O42-

D. PO43-

C2O42-

NH3

10. Water has the greatest tendency to act as an acid with which of the following?

A. Cl- C. H2PO4

-

B. NO2- D. CH3COO

-

11. When comparing equal volumes of 0.10 M HNO3 with 0.10 M HNO2, what would be observed?

A. The pH values would be the same.

B. The electrical conductivities would be different.

C. The effects on blue litmus paper would be different.

D. The volumes of 0.10 M NaOH needed for neutralization would be different.

12. Which of the following 1.0 M solutions would have the highest electrical conductivity?

A. HI C. HCN

B. HF D. HNO2

13. Which of the following solutions will have the lowest electrical conductivity?

A. 1.0 M HI C. 1.0 M NaOH

B. 1.0 M H2S D. 1.0 M NaNO3



19

14. Consider the following equilibrium:

HF(aq) + HPO42-

(aq) ⇌ F-(aq) + H2PO4

-(aq)

For the above equilibrium, identify the weaker acid and determine whether reactants or products are

favoured.

Weaker Acid Side Favoured

A. HF Products

B. HF Reactants

C. H2PO4- Products

D. H2PO4-

Reactants

15. Which of the following is the strongest acid that can exist in an aqueous solution?

A. O2-

C. H3O+

B. NH2- D. HClO4

16. Consider the following equilibria:

I. CH3COOH + OCN- ⇌ HOCN + CH3COO

-

II. CH3COOH + ClO- ⇌ HClO + CH3COO

-

a) In equation I above, the reactants are favoured. Identify the stronger acid.

b) In equation II above, the products are favoured. Identify the stronger acid.

c) Consider the following reaction:

HOCN + ClO- ⇌ OCN

- + HClO

Does this reaction favour the reactants or products? Explain.

__________________________________________________________________________________

__________________________________________________________________________________

__________________________________________________________________________________



20

A.4: Hydrolysis of Salts (Part 1)

According to Arrhenius:

Acid + Base Salt + Water

This is called a neutralization reaction.

However, in many cases where we react an acid and base in equivalent amounts, the resulting salt

solution may be slightly acidic (pH 7.0 at 25C) or slightly basic (pH 7.0 at 25C).

For example, the neutralization reaction between nitric acid and ammonium hydroxide is:

HNO3(aq) + NH4OH(aq) NH4NO3(aq) + H2O(l)

The resulting solution of NH4NO3 is slightly acidic.

The neutralization of acetic acid by sodium hydroxide is:

CH3COOH(aq) + NaOH(aq) NaCH3COO(aq) + H2O(l)

The resulting NaCH3COO solution is slightly basic.

The above results can be very easily explained if one examines the table of relative acid strengths.

In the case of the NH4NO3 solution, we can see that NH4+ is a weak acid:

NH4+

(aq) + H2O(l) ⇌ H3O+

(aq) + NH3(aq)

In the case of the NaCH3COO solution, CH3COO- is a weak base:

CH3COO-(aq) + H2O(l) ⇌ CH3COOH(aq) + OH

-(aq)

When an anion or cation reacts with water as an acid or a base, it is said to hydrolyze. Similarly, if a

salt contains an ion that reacts with water as an acid or a base, that salt is said to hydrolyze. The

equation showing an ion acting as an acid or a base is called a HYDROLYSIS REACTION

EQUATION.

As we have seen, a large number of Brønsted-Lowry bases are anions. In addition, some anions and a

few cations hydrolyze as Brønsted-Lowry acids.

Any salt containing an ion that hydrolyzes will form a solution that may be acidic or basic.

So which ions hydrolyze and which ones do not?



21

Cations That Hydrolyze

Neither the Group I nor the Group II metal ions hydrolyze. In fact, for our purposes, we only need to

concern ourselves with four cations that hydrolyze, all as acids. These are NH4+, Al

3+, Fe

3+ and Cr

3+.

We have already examined the hydrolysis of NH4+. The other three cations hydrolyze in a manner

similar to one another.

Because of the large positive charge on these ions, water molecules (being very polar) are strongly

attracted to them, forming a complex hydrated ion.

Cr3+

(aq) + 6H2O(l) Cr(H2O)63+

(aq)

Al3+

(aq) + 6H2O(l) Al(H2O)63+

(aq)

Fe3+

(aq) + 6H2O(l) Fe(H2O)63+

(aq)

Each of these complex ions will donate a proton to a water molecule and produce a H3O+ ion.

Cr(H2O)63+

(aq) + H2O(l) ⇌ H3O+

(aq) + Cr(H2O)5(OH)2+

(aq)

Al(H2O)63+

(aq) + H2O(l) ⇌ H3O+

(aq) + Al(H2O)5(OH)2+

(aq)

Fe(H2O)63+

(aq) + H2O(l) ⇌ H3O+

(aq) + Fe(H2O)5(OH)2+

(aq)

Anions That Hydrolyze

None of the anions that are conjugate bases of the six strong acids hydrolyze EXCEPT HSO4-. HSO4

-

hydrolyzes as a weak acid in water as shown below:

HSO4-(aq) + H2O(l) ⇌ H3O

+(aq) + SO4

2-(aq)

The conjugate base of HSO4- is SO4

2-. It will hydrolyze as a pathetically weak base (remember, the

stronger an acid is, the weaker is its conjugate base):

SO42-

(aq) + H2O(l) ⇌ HSO4-(aq) + OH

-(aq)

As we know, the anions OH-, NH2

- and O

2- are all strong bases in solution.

ALL OTHER anions not containing a hydrogen atom will tend to hydrolyze as weak bases. For

example:

SO32-

(aq) + H2O(l) ⇌ HSO3-(aq) + OH

-(aq)

A difficulty arises when dealing with anions containing hydrogen atom. As we saw earlier, such ions

are amphiprotic, and could hydrolyze as and acid or as a base!



22

To determine whether an amphiprotic ion hydrolyzes primarily as an acid or as a base requires the use

of mathematical methods that we will learn later in this course.

Examples:

Predict whether the following salts form solutions that are acidic, basic or neutral. Write the

corresponding hydrolysis reaction equation(s) when applicable.

(a) NaBr

(b) KNO2

(c) Fe(NO3)3

(d) NH4CN

Provincial Exam Examples: The Hydrolysis of Salts

1. Consider the following reaction:

NO2-(aq) + H2O(l) ⇌ HNO2(aq) + OH

-(aq)

This reaction represents which of the following?

A. the titration of NO2- C. the ionization of HNO2

B. the hydrolysis of NaNO2 D. the dissociation of NaNO2

2. Which of the following describes the net ionic equation reaction for the hydrolysis of NH4Cl(s)?

A. NH4+

(aq) + Cl-(aq) ⇌ NH4Cl(s)

B. NH4Cl(s) ⇌ NH4+

(aq) + Cl-(aq)

C. Cl-(aq) + H2O(l) ⇌ HCl(aq) + OH

-(aq)

D. NH4+

(aq) + H2O(l) ⇌ H3O+

(aq) + NH3(aq)



23

3. Which of the following is the net ionic equation describing the hydrolysis of KCN?

A. K+

(aq) + H2O(l) ⇌ KOH(aq) + H+

(aq)

B. KCN(aq) + H2O(l) ⇌ K+

(aq) + CN-(aq)

C. CN-(aq) + H2O(l) ⇌ HCN(aq) + OH

-(aq)

D. CN-(aq) + H2O(l) ⇌ 2H

+(aq) + CNO

-(aq)

It is now time to complete and submit the Unit 4 Topic A Hand-in Assignment.

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