Download Acid Base Equilibria

ACID – BASE EQUILIBRIA
Introduction
Acids and bases are two important concepts in chemistry. Acids are substances with a pH of
anything below 7 while bases are substances with a pH of anything above 7, making them
opposites in a sense. However, some substances--such as distilled water--are neutral, which have
a pH of exactly 7
The Nature of Acids and Bases
Strong and weak acids are defined by their ease of losing (or donating) a proton.
The term "acid" refers to a group of compounds whose chemical properties are similar.
1. They dissolve in water to give a solution with pH less than 7
2. They react with bases to give a salt and water
3. They give distinctive colours when in contact with indicators
The term "base" refers to a group of compounds whose chemical properties are similar.
1. They dissolve in water to give a solution with pH more than 7.
2. They react with acids to give a salt and water
3. They give distinctive colours when in contact with indicators
There are several types of compounds which all come under the general heading of bases.
Base
Example
Formula
Solubility
Metal hydroxides
Sodium hydroxide
NaOH
very soluble
Metal oxides
Calcium oxide
CaO
reacts with water
Metal carbonates
Magnesium carbonate
MgCO3
insoluble
Metal hydrogen carbonates
Sodium hydrogen carbonate
NaHCO3
soluble
Metal sulphites (sulphate IV) Sodium sulphite
Na2SO3
soluble
Ammonia solution
NH3(aq)
very soluble
Ammonia (aq)
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Note: An alkali is simply a soluble metal hydroxide (base)
Indicators are compounds that change colour depending on whether they're in acidic or basic
conditions
Methyl orange Bromophenol blue
Methyl red Bromothymol blue
Phenolphthalein
Acid red
yellow
red
yellow
colourless
Base yellow
blue
yellow
blue
red
Each one changes colour at a different pH, and so there will be situations where one is useful and
others are not.
Reactions of acids with bases (neutralization reactions)
Reaction of acids with metal hydroxides or metal oxides: Makes a salt + water,
HCl + NaOH -> H2O + NaCl.
Reaction of acids with metals: Makes a salt + hydrogen
2HCl + Mg -> MgCl2 + H2.
Reaction of acids with carbonates: Makes a salt + water + carbon dioxide
2HCl + CaCO3 -> CO2 + H2O + CaCl2.
Experimental properties of acids and bases
When acids and bases neutralize, the reaction is noticeably exothermic (i.e. heat can be felt
coming from the reaction). Obviously, they will have an effect on the colour of indicators as
described above. The hydrogen produced in the reaction of acids with metal will produce a 'pop'
sound if a match is held to it, and the CO2 from the carbonate reaction will turn limewater a
milky white when bubbled through it.
Examples of some common acids: HCl, CH3COOH, H2SO4, HNO3
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Examples of bases: NaOH, NH3
Strong and Weak Acids and Bases
A strong acid, when placed in water, will almost fully ionize/dissociate straight away, producing
H3O+ ions from water.
A weak acid will, however, only partially dissociate into ions, leaving a high percentage of
unreacted molecules in the solution. Equilibrium is established, and so when some of the H3O+
ions produced by a weak acid react, Le Chatelier's principle means that more of the acid will
react to form H3O+ ions. This means that, given an equal number of moles of acid, they will be
neutralized by the same amount of strong base, but their solutions will have different pH values.
A weak base is the same as this, only it accepts protons and so produces OH- ions from water
rather than H3O+.
Conductivity
Any solution's ability to conduct electricity is conditioned by the concentration of ions it
contains. A strong acid has more ions than a weak one, and so its solution will be a better
electrical conductor than a weak acid. The same goes for strong/weak bases.




Strong acids: HCl, HNO3, H2SO4.
Weak acids: CH3COOH, H2CO3.
Strong bases: group 1 hydroxides (i.e. NaOH etc), or lower group 2 hydroxides Ba (OH)2.
Weak bases: NH3, CH3CH2NH2.
The strength of an acid or base can be measured with an indicator (universal) or a pH meter.
Also the rate of reaction...hydrogen production with metals or CO2 with CaCO3 will reveal the
strength of an acid.
Acid – Base Theories
There are three important theories regarding acids and bases.



Arrhenius
Bronsted Lowry
Lewis
Arrhenius refers to hydrogen ions in aqueous solution as being the fundamental cause of acidity.
In this theory all acids release free hydrogen ions into the solution by dissociation. These
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hydrogen ions are solvated by water molecules to form other species such as the hydronium ion,
H3O+. In general the hydrogen ions are more simply referred to as H+ (aq).
In Arrhenius theory a basic solution is caused by free hydroxide ions, OH-.
Bronsted Lowry looks at acidity from the point of view of hydrogen ions. This theory places no
restrictions, such as the need for aqueous solution. In Bronsted Lowry theory an acid is any
species that can release a hydrogen ion. Bases become any species that can accept a hydrogen
ion. An Arrhenius acid is, of course, also a Bronsted Lowry acid, but the reverse is not
necessarily true. Bronsted Lowry theory covers Arrhenius and extends it.
Lewis approaches acidity from the point of view of the electron pairs necessary to bond to the
hydrogen ions and extends the concept of acidity to include any species that is able to accept an
electron pair. A Lewis base is the species with the available electron pair.
Once again Lewis theory encompasses the previous two theories and extends them. Now there is
no need for hydrogen ion and any dative coordinate bonding becomes an acid - base reaction.
Arrhenius definition of acids and bases:
 Acid: a substance that, when dissolved in water, increases the concentration of H+ ions
An example is hydrogen chloride:
H2O
HCl(g)
H+(aq) + Cl-(aq)
 Base: a substance that, when dissolved in water, increases the concentration of OH- ions
An example is sodium hydroxide:
NaOH(s)
H2O
Na+(aq) + OH-(aq)
Bronsted – Lowry definition of acids and bases:
The proton; H+
Since the H+ ion is formed by the loss of the only electron in a H atom it is essentially a proton
without any surrounding valence electron. In water the proton interacts with water molecule to
form the hydronium ion:
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..
:O
H+
H
..
O
H
H
H
H2O
H+
H
H3O+
Note: H+ is usually used, but bear in mind that reality is that it exists as the hydrated hydronium
ion.
According to the Bronsted – Lowry theory of acids and bases:
 Acid: a substance (molecule or ion) that can transfer a proton to another substance
HCl(g)
H2O(l)
H3O+(aq) + Cl-(aq)
HCl acts as a Bronsted – Lowry acid – it donates a proton to H2O and H2O acts as a
Bronsted – Lowry base (it accepts a proton from HCl)
 Base: a substance that can accept a proton from another substance
Since emphasis in the Bronsted – Lowry concept is on proton transfer, the concept is also
applicable to reactions that do not occur in aqueous solution, for example the gas phase reaction:
HCl(g) + NH3(g)
NH4Cl(s) (NH4+ + Cl-)
A proton is transferred from the acid HCl to the base NH3.
Another example:
NH3(aq) + H2O(l)

NH4+(aq) + OH-(aq)
NH3 is an Arrhenius base – its addition to water leads in an increase in OH- ions. It is a
Bronsted – Lowry base because it accepts a proton from water.

H2O acts as a Bronsted – Lowry acid because it donates a proton to the NH3 molecule.
To be a Bronsted – Lowry acid, a molecule or ion must have a hydrogen atom to be able to
donate an H+ ion.
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To be a Bronsted – Lowry base, a molecule or ion must have a non-bonding pair of electrons that
it can use to bind the H+ ion.
Note: Some substances can act as an acid in one reaction and as a base in another, e.g. Water is a
Bronsted – Lowry base in its reaction with HCl and a Bronsted – Lowry acid in its reaction with
NH3. Such substances are said to AMPHOTERIC.
Conjugate acid – base pairs
Conjugate acids and bases is another important topic. Simply put, it's where two compounds
transform into each other by a gain or loss of a proton. The conjugate acid is seen as the chemical
substance that releases or donates a proton in a forward (arrow pointing right) chemical reaction,
while the conjugate base absorbs or gains the proton in a backward (arrow pointing left)
chemical reaction. Common conjugate pairs are water (H2O) and a hydronium ion (H3O+),
ammonium ion (NH4+) and ammonia (NH3), and hydrochloric acid (HCl) and a chloride ion (Cl). However, this is only viable within the Bronsted-Lowry acid-base theory.
In acid – base equilibria, both the forward and reverse reactions involve proton transfer.
Consider the reaction of an acid HX with water:
remove H+
HX(aq) + H2O(l)
Acid
X-(aq) + H3O+(aq)
Base
Conjugate
base
Conjugate
acid
add H+
An acid and a base such as HX and X- that differ only in the presence or absence of a proton are
called a conjugate acid – base pair. Every acid has a conjugate base and every base has a
conjugate acid.
Examples:
The acid HClO4 has a conjugate base ClO4-.
The conjugate acid for CN- is HCN
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Relative Strengths of Acids and Bases
The strength of an acid is measured by its ability to donate a proton and that of a base by its
ability to accept a proton.
Strong acid: one that completely transfers its proton to water, leaving no undissociated
molecules in solution
For example: HN03 dissociates completely in water to form H+ and N03-.
The reaction is
HNO3 (aq) → H+ (aq) + N03-(aq)
A 0.01 M solution of nitric acid contains 0.01 M of H+ and 0.01 M N03- ions and almost no
HN03 molecules. The pH of the solution would be 2.0.
There are only 6 strong acids: HCl, H2SO4, HNO3, HClO4, HBr and HI
The rest of the acids therefore are considered weak acids.
Weak acid: one that only partially dissociates in aqueous solution and therefore exists in the
solution as a mixture of acid molecules component ions:
For example, HF dissociates in water to give H+ and F-. It is a weak acid with a dissociation
equation that is
HF (aq) ↔ H+ (aq) + F-(aq)
Note the use of the double arrow with the weak acid. That is because equilibrium exists between
the dissociated ions and the undissociated molecule. In the case of a strong acid dissociating,
only one arrow ( → ) is required since the reaction goes virtually to completion.
An equilibrium expression can be written for this system:
Ka = [ H+][F-] / [HF]
 The stronger the acid the weaker is its conjugate base
 The stronger the base the weaker is its conjugate acid
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Strong Bases:
Dissociate 100% into the cation and OH- (hydroxide ion).
Example: NaOH (aq) → Na+(aq) + OH-(aq)
A 0.010 M NaOH solution will contain 0.010 M OH- ions (as well as 0.010 M Na+ ions) and
have a pH of 12.
Examples of strong bases are the hydroxides of Groups I and II.
Note: the hydroxides of Group II metals produce 2 mol of OH- ions for every mole of base that
dissociates. These hydroxides are not very soluble, but what amount that does dissolve
completely dissociates into ions.
Example: Ba (OH)2(aq) → Ba2+(aq) + 2OH-(aq)
A 0.000100 M Ba(OH)2 solution will be 0.000200 M in OH- ions (as well as 0.000100 M in
Ba2+ ions) and will have a pH of 10.3.
Weak Bases:
Most weak bases are anions of weak acids. Weak bases do not furnish OH- ions by dissociation.
They react with water to furnish the OH- ions.
Note that like weak acids, this reaction is shown to be at equilibrium, unlike the dissociation of a
strong base which is shown to go to completion.
When a weak base reacts with water the OH- comes from the water and the remaining H+
attaches itself to the weak base, giving a weak acid as one of the products. You may think of it as
a two-step reaction similar to the hydrolysis of water by cations to give acid solutions.
Examples:
NH3(aq) + H2O(aq) → NH4+(aq) + OH-(aq)
methylamine: CH3NH2(aq) + H20(l) → CH3NH3+(aq) + OH-(aq)
acetate ion: C2H3O2-(aq) + H2O(aq) → HC2H302(aq) + OH-(aq)
General reaction: weak base(aq) + H2O(aq) → weak acid(aq) + OH-(aq)
Since the reaction does not go to completion relatively few OH- ions are formed.
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Acid-Base Properties of Salt Solutions
Definition of a salt:


an ionic compound made of a cation and an anion, other than hydroxide.
the product besides water of a neutralization reaction
Determining acidity or basicity of a salt solution:
1. Split the salt into cation and anion
2. Add OH- to the cation
a. if you obtain a strong base, the cation is neutral
b. if you get a weak base, the cation is acidic
3. Add H+ to the anion
a. if you obtain a strong acid, the anion is neutral
b. if you obtain a weak acid, the anion is basic
4. Salt solutions are neutral if both ions are neutral
5. Salt solutions are acidic if one ion is neutral and the other is acidic
6. Salt solutions are basic is one of the ions is basic and the other is neutral.
The acidity or basicity of a salt made of one acidic ion and one basic ion cannot be determined
without further information
Acid-Base Titration
Acid-base titration is the determination of the concentration of an acid or base by exactly
neutralizing the acid/base with a different acid or base of a known concentration. This allows for
quantitative analysis of the concentration regarding the unknown acid or base solution, making
use of the neutralization reaction that commonly occurs between acids and bases. It can also be
used to find the chemical percent purity.
Common Equipment
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






Burette
White tile (used to see a colour change in the solution)
Pipette
pH Indicator (varies depending on reactants)
Erlenmeyer/Conical flask
Titrant (a standard solution of known concentration)
Analyte (a solution of unknown concentration)
Acid-base titrations usually react in three ways:
1. A strong acid will react with a weak base to form an acidic (pH<7) solution.
2. A strong acid will react with a strong base to form a neutral (pH=7) solution.
3. A weak acid will react with a strong base to form a basic (pH>7) solution.
The equivalence point (also known as the stoichiometric point) of an acid-base titration is when
a titrant is added and stoichiometrically equal to the amount of moles of the analyte present in
the sample; in simpler words, the smallest amount of titrant that is sufficient to fully neutralize or
react with the analyte. In some cases there are multiple equivalence points, though they are
simply multiples of the first equivalent point. Shown on a graph of a titration curve, the
equivalence point is the inflection point or the part of the graph where it is the steepest. The end
point (similar though not the same as the equivalence point) refers to the point at which the
indicator changes colour in a colorimetric titration.
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1. This titration curve shows the result of a titration between a strong acid and a weak base.
2. This titration curve shows the result of a titration between a strong acid and a strong base.
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3. This titration curve shows the result of a titration between a strong base and a weak acid.
Dissociation of Water
As mentioned earlier, water can act as either a Bronsted acid or base depending on
circumstances. In the presence of an acid water acts as a proton acceptor, and in the presence of a
base, acts as a proton donor. So one molecule of water can donate a proton to another water
molecule – process referred to as auto-ionization of water:
..
O:
H
H
H
..
O:
H
..
O
H
..
:O
..
H-
H
H
H3O+
H2O + H2O
+
OH-
The reaction is very fast in both directions, such that no individual molecule remains ionized for
long, which is why pure water is a poor conductor of electricity.
From the above reaction the equilibrium – constant expression is:
Kc
=
[H3O+][OH-]
[H2O]2
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Re-arranging gives:
Kc[H2O]2 = [H3O+][OH]As previously explained, we exclude the concentration of pure water, and the new constant is
denoted Kw – the ionic-product constant for water.
At 25oC Kw is 1.0 x 10-14:
Kw = [H3O+][OH]- = 1.0 x 10-14
Since H3O+ and H+ are used interchangeably to represent the hydrated proton, the above equation
is simplified to:
Kw = [H+][OH]- = 1.0 x 10-14
The equation applies to aqueous solutions and can be used to calculate either [H+] or [OH-] if one
of these concentrations are known.
If [H+] = [OH-], the solution is said to be neutral.
Example: What is the [H+] of a sample of lake water with [OH-] of 4.0 x 10-9 M? Is the lake
acidic, basic, or neutral?
Solution: [H+] = 1 x 10-14 / 4 x 10-9 = 2.5 x 10-6 M
Therefore the lake is slightly acidic.
The pH Scale
The pH scale is an important topic regarding acids and bases, as it not only tells whether or not
the substance is acidic or basic, but also measures its acidity/alkalinity/neutrality. The process is
done by the logarithm of hydrogen-ion (H+) concentration in gram atoms per litre.
pH:
the measure of concentration of protons (H
the proton donation reaction.
pH Definition:
ion) in water, or essentially the strength of
pH is the negative logarithm of the concentration of H . So an acid has
low pH, and therefore high concentrations of H , and can participate
more readily in reactions that require donation of a proton.
The [H+] in aqueous solutions is usually very small, and it is more convenient to express [H +] in
terms of pH, defined as the negative logarithm in base 10 of [H+]:
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pH = -log[H+]
Neutral pH: at neutral pH there are equal ``concentrations'' of H and OH in the solution.
At room temperature neutral pH is 7. ``Neutral'' really means that there is equal
tendency of solution for donation or acceptance of protons.
For a neutral solution [H+] = 1.0 x 10-7M thus,
pH = -log(1.0 x 10-7) = 7.00
The pH of a neutral solution is 7.
An acidic solution is one in which [H+] ˃ 1.0 x 10-7 M, so pH decreases as [H+] increases.
At 25oC the pH of an acidic solution is less than 7, and that of a basic solution is greater than 7.
pH depends strongly on temperature, as shown in the following graph:
Example 1: Calculate the [OH-] of a solution of baking soda with a pH of 8.5.
Solution: First calculate the [H+]
If pH is 8.5, then the antilog of -8.5 is 3.2 x 10-9. Thus the [H+] is 3.2 x 10-9 M
Next calculate the [OH-]
1.0 x 10-14 / 3.2 x 10-9 = 3.1 x 10-6 M
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Example 2: Calculate the pH of a solution of household ammonia whose [OH-] is 7.93 x 10-3 M.
Solution: This time you first calculate the [H+] from the [OH-]
7.93 x 10-3 M OH- = 1.26 x 10-12 M H+
Then find the pH
-log[1.26 x 10-12] = 11.9
Example
Calculate the concentration of H+(aq) and pH of
(a) a solution in which [OH]- is 0.010 M;
(b) a solution in which [OH]- is 2.09 x 10-9 M.
pH Problem Solving Diagram
Hydrolysis
Hydrolysis is a chemical reaction of a compound with water, usually resulting in the formation of
one or more new compounds. The most common hydrolysis occurs when a salt of a weak acid or
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weak base (or both) is dissolved in water. Water ionizes into negative hydroxyl ions (OH−) and
positive hydrogen ions (H+), which become hydrated to form positive hydronium ions (H3O+).
The salt also breaks up into positive and negative ions. For example, when sodium acetate is
dissolved in water it readily dissociates into sodium and acetate ions. Because sodium hydroxide
is a strong base, the sodium ions react only slightly with the hydroxyl ions already present in the
water to form sodium hydroxide molecules. Acetic acid is a weak acid, so the acetate ions react
readily with the hydrogen ions present in the water to form neutral acetic acid molecules. The net
result of these reactions is a relative excess of hydroxyl ions, causing an alkaline solution. A
chemical reaction has actually taken place between the water and the dissolved salt. There are
relatively few instances in which water reacts directly with organic compounds under ordinary
conditions. It does react with acid halides, acid anhydrides, and organo-metallic compounds, e.g.,
Grignard reagents. The addition of strong acids or bases or the use of steam will often bring
about hydrolysis where ordinary water has no effect. Some industrially important hydrolysis
reactions are the synthesis of alcohols from olefins (e.g., ethanol, CH3COOH, from ethene,
CH2CH2) in the presence of a strong acid catalyst, the conversion of starches to sugars in the
presence of a strong acid catalyst, and the conversion of animal fats or vegetable oils to glycerol
and fatty acids by reaction with steam. Hydrolysis is an important reaction in plants and animals
(see metabolism). The catalytic action of certain enzymes allows the hydrolysis of proteins, fats,
oils, and carbohydrates.
Some real life examples of hydrolysis
The most common hydrolysis occurs when a salt of a weak acid or weak base (or both) is
dissolved in water. Water auto-ionizes into negative hydroxyl ions and hydrogen ions. The salt
breaks down into positive and negative ions. For example, sodium acetate dissociates in water
into sodium and acetate ions. Sodium ions react very little with hydroxyl ions whereas acetate
ions combine with hydrogen ions to produce neutral acetic acid, and the net result is a relative
excess of hydroxyl ions, causing a basic solution.
However, under normal conditions, only a few reactions between water and organic compounds
occur. Generally, strong acids or bases must be added in order to achieve hydrolysis where water
has no effect. The acid or base is considered a catalyst. They are meant to speed up the reaction,
but are recovered at the end of it.
Acid–base-catalyzed hydrolyses are very common; one example is the hydrolysis of amides or
esters. Their hydrolysis occurs when the nucleophile (a nucleus-seeking agent, e.g., water or
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hydroxyl ion) attacks the carbon of the carbonyl group of the ester or amide. In an aqueous base,
hydroxyl ions are better nucleophiles than dipoles such as water. In acid, the carbonyl group
becomes protonated, and this leads to a much easier nucleophilic attack. The products for both
hydrolyses are compounds with carboxylic acid groups.
Perhaps the oldest example of ester hydrolysis is the process called saponification. It is the
hydrolysis of a triglyceride (fat) with an aqueous base such as sodium hydroxide (NaOH).
During the process, glycerol, also commercially named glycerin, is formed, and the fatty acids
react with the base, converting them to salts. These salts are called soaps, commonly used in
households.
Moreover, hydrolysis is an important process in plants and animals, the most significant example
being energy metabolism and storage. All living cells require a continual supply of energy for
two main purposes: for the biosynthesis of small and macromolecules, and for the active
transport of ions and molecules across cell membranes. The energy derived from the oxidation
of nutrients is not used directly but, by means of a complex and long sequence of reactions, it is
channeled into a special energy-storage molecule, adenosine triphosphate (ATP).
The ATP molecule contains pyrophosphate linkages (bonds formed when two phosphate units
are combined together) that release energy when needed. ATP can be hydrolyzed in two ways:
the removal of terminal phosphate to form adenosine diphosphate (ADP) and inorganic
phosphate, or the removal of a terminal diphosphate to yield adenosine monophosphate (AMP)
and pyrophosphate. The latter is usually cleaved further to yield two phosphates. This results in
biosynthesis reactions, which do not occur alone, that can be driven in the direction of synthesis
when the phosphate bonds are hydrolyzed.
In addition, in living systems, most biochemical reactions, including ATP hydrolysis, take place
during the catalysis of enzymes. The catalytic action of enzymes allows the hydrolysis of
proteins, fats, oils, and carbohydrates. As an example, one may consider proteases, enzymes that
aid digestion by hydrolyzing peptide bonds in proteins. They catalyze the hydrolysis of interior
peptide bonds in peptide chains, as opposed to exo-peptidases, another class of enzymes that
catalyze the hydrolysis of terminal peptide bonds, liberating one free amino acid at a time.
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However, proteases do not catalyze the hydrolysis of all kinds of proteins. Their action is stereoselective: Only proteins with a certain tertiary structure will be targeted. The reason is that some
kind of orienting force is needed to place the amide group in the proper position for catalysis.
The necessary contacts between an enzyme and its substrates (proteins) are created because the
enzyme folds in such a way as to form a crevice into which the substrate fits; the crevice also
contains the catalytic groups. Therefore, proteins that do not fit into the crevice will not be
hydrolyzed. This specificity preserves the integrity of other proteins such as hormones, and
therefore the biological system continues to function
Buffer Solutions
Buffers are aqueous solutions consisting of a mixture between a weak acid and its conjugate
base or a weak base and its conjugate acid. They possess the property that the pH of said solution
changes very little when a small amount of strong acid/base is added to it. They mostly are used
as a means of keeping pH at a nearly constant value in a wide variety of chemical applications.
There are two types of buffers: acidic and basic.
Acidic buffers comprise of a weak acid and its salt e.g. ethanoic acid and sodium ethanoate.
Basic buffers comprise of a weak base and its conjugate salt; e.g. aqueous ammonia and
ammonium chloride.
Buffers resist changes in pH value of solution on small dilutions.
A buffer is used to resist the change in pH when an acid or alkali is added to a solution.
The scope of buffers is to maintain the pH of a solution at a constant value.
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