8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 1/31

 ARRHENIUS DEFINITION (1890)This considered that acids were substances that released H+ ions in solution and that bases were

substances that produced OH- ions in solution. It explained adequately how acids and bases neutralized

each other

i.e. H+(aq) + OH-

(aq)  H2O(l) 

acid base neutral

However, as time passed the Arrhenius definition began to seem less satisfactory. For one thing, there

seemed to be two kinds of base. Metal hydroxides such as NaOH produce OH- ions in water by ionic

dissociation. Bases, such as ammonia produce OH- ions in aqueous solution by undergoing a reaction

with water.

NaOH(s)  Na+

(aq) + OH-(aq) 

NH3(g) + H2O(l) NH+

4(aq) + OH-(aq)

But more important, the Arrhenius definition of acids and bases was narrow.

It defined acids and bases in aqueous solution only.

THE LOWRY-BRONSTED DEFINITION (1923) 

Lowry-Bronsted produced more general definitions of acids and bases.

These were in terms of H+

ions (called a protons, since a H+

ion has neither electrons nor neutrons).

According to Bronsted - an acid is a substance that can donate a proton

a base is a substance that can accept a proton 

This definition can be represented by the general chemical reaction

A B + H+ 

which does not attempt to show electrical charge balance.

In this equation -

A is the acid.,

B is the base and

H+ (a hydrogen atom without an electron) is a proton.

Together A and B are called a CONJUGATE ACID AND BASE PAIR.

We can call B the conjugate base of A and call A the conjugate acid of B. This Lowry-Bronsted definition

is more general than the Arrhenius definition. It does not refer to any specific solvent. We can simply say

that any substance that can lose a proton is an acid. The following are examples of Bronsted acids -

Molecules Anions CationsHCl HSO4

-NH4

+ 

H2SO4 HCO3-

CH3NH3+ 

HCN HPO42- 

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 2/31

Similarly the Bronsted bases are substances which can accept protons and include the following examples

:-

Molecules Anions Cations

NH3 OH- Mg(OH)+ 

CH3NH2 HCO3- Al(OH)2+ HPO4

2- 

e.g. Mg(OH)+ + H+ = Mg2+ + H2O

While the relationship A B + H+

is a very good general definition of an acid and a base, there is a

problem in applying this definition to an acid/base system in solution. The problem arises because the

free protons (H+) cannot exist in solution to any great extent. In many cases the protons interact with the

solvent

e.g. H2O(l) + H+(aq)  H3O+(aq) (a hydroxonium ion)

The solvent is then acting as a Bronsted base in accepting a proton.

The same reasoning can be applied to a base in solution. The base must obtain a proton from a proton

source. Very often the proton source is the solvent.

H2O

NH3(g) NH+4(aq) 

The solvent is then acting as a Bronsted acid by providing a proton.

We can see that acid base reactions in solution are not simply processes in which an acid loses a proton or a base gains a proton. We must regard all acid/base reactions in solution as PROTON TRANSFER

REACTIONS. Any such reaction includes TWO ACID - BASE CONJUGATE PAIRS. The original

acid base pair, plus another conjugate pair, to accept the proton from the acid or donate the proton to the

base. As we have said, this second acid-base pair is often derived from the solvent.

NAME OF ACID ACID BASE NAME OF BASE

Hydrochloric acid HCl Cl- Chloride

Hydrogen sulphate HSO4- SO4

2- Sulphate

Ammonium NH4+

NH3 Ammonia

Ammonia NH3 NH2- Amide

Water H2O OH-

Hydroxide

Each substance in the acid column is converted to its partner in the base column by the removal of a

proton. Each substance in the base column is converted to its partner in the acid column by the addition of a proton.

ANY TWO CONJUGATE ACID/BASE PAIRS CAN BE C0MBINED IN AN ACID/BASE REACTION

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 3/31

e.g. HCl + OH- H2O + Cl-

HCl and Cl- are one conjugate acid-base pair and H2O and OH- are another pair. The proton donated by

the acid HCl is accepted by the base OH- in the forward reaction and a proton accepted by the Cl - is

donated by the H2O in the reverse reaction. 

The Bronsted definition can be summarized in a few sentences.An acid is -

a proton donor

a base a proton acceptor.

We emphasize the relationship by designating conjugate acid base pairs.

e.g. HCl/Cl- 

Acid base reactions in solution require two conjugate acid/base pairs because free protons (H+) do not

exist in solution

THE LEWIS DEFINITION (1938)

The Lowry-Bronsted theory has its short-comings. Since it defines an acid as a proton donor, it excludesfrom the category of acids, substances that have no protons to donate. This limitation was overcome by a

more general definition of acids and bases, BASED ON ELECTRONIC STRUCTURE, which was

proposed by G N Lewis.

By the Lowry-Bronsted definition, a substance must accept a proton to be classified as a base. In other

words the base must form a chemical bond with the proton (H+

ion). Since the proton has no electrons the

base must have an electron pair available to form a bond.

In the Lewis definition:-

A BASE IS A SUBSTANCE THAT HAS A NON-BONDING VALENCE ELECTRON PAIR

THAT CAN BE USED TO FORM A CHEMICAL BOND.

MORE SIMPLY - A LEWIS BASE IS AN ELECTRON - PAIR DONOR.

THE LEWIS DEFINITION DOES NOT GREATLY EXPAND OUR IDEAS ABOUT BASES, BUT IT

DOES SIGNIFICANTLY BROADEN THE CATEGORY OF SUBSTANCES THAT CAN BE

CLASSIFIED AS ACIDS. WHEN A BASE ACCEPTS A PROTON IT DONATES ELECTRONS TO

FORM THE BOND WITH THE PROTON. THEREFORE THE PROTON ACCEPTS THE

ELECTRONS.

A LEWIS ACID IS A SUBSTANCE THAT IS AN ELECTRON-PAIR ACCEPTOR

A proton (H+) is the simplest Lewis acid but many other substances fit the definition, including a number

of cations,  e.g Zn2+

 

Zn2+

+ 2OH-  Zn(OH)2

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 4/31

 

Here the Zn2+ is the Lewis acid and OH- the Lewis base.

In organic chemistry it is especially common to find cations that behave as Lewis acids., accepting

electrons. The reaction :-

C2H4 + Br2 C2H4Br2 

can be considered to proceed in 3 steps.

1. Br2 forms Br+

and Br 

ions. The Br+

is the Lewis acid.

2. This joins to C2H4 (a Lewis base).

3. In the third step Br-, a Lewis base, donates an electron pair and joins to the cation formed in the first

step. The cation is a Lewis acid

H H H H

Br C C + + Br Br C C Br

H H H H

A great many reactions can be understood as the combination of a Lewis acid and a Lewis base. In fact acommon method of classifying chemical reagents uses the Lewis acid and Lewis base definition.

A LEWIS ACID, WHICH IS A SUBSTANCE SEEKING ELECTRONS, IS CALLED AN

ELECTROPHILE. 

A LEWIS BASE, WHICH IS AN ELECTRON DONOR (OR A SUBSTANCE THAT SEEKS

SUBSTANCES THAT ARE ELECTRON DEFICIENT), CAN BE CALLED A NUCLEOPHILE. 

In the above reaction between C2H4 and Br2, Br+ is an electrophile and Br- is a nucleophile.

Many Lewis acids are substances that are chemically neutral. These substances can accept electrons and

form additional bonds, either because they have incomplete octets or octet expansion is possible.

Many compounds of the elements of Group III are Lewis acids, the best known being the halides such as

BF3, BCl3, AlCl3 and AlBr3 

Question: Use the Lewis theory to explain what is happening when AlCl3 molecules dimerise

Cl Cl Cl

Al Al

Cl Cl Cl

Co-ordinate (dative covalent) bond in Al2Cl6 

OTHER EXAMPLES INCLUDE -

BF3 + NH3 F3B NH3

Lewis Acid Lewis Base

AlCl3 + Cl- AlCl4

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 5/31

Lewis Acid Lewis Base

Conclusion.The Lowry-Bronsted definition is quite adequate for the chemistry of acids and bases in aqueous solution.

The Lewis definition is useful because it allows us to classify so many substances as acids or bases, and

because it provides an understanding of the details of many chemical reactions.

Acid-Base Reaction TheoriesAcids and bases are everywhere. Some foods contain acid, like the citric acid in lemons and the lactic acid

in dairy. Cleaning products like bleach and ammonia are bases. Chemicals that are acidic or basic are an

important part of chemistry.

Helpful Hint!

You may need to refresh your memory on

naming acids. 

Several different theories explain what composes an acid and a base. The first scientific definition of an

acid was proposed by the French chemist Antoine Lavoisier in the eighteenth century. He proposed that

acids contained oxygen, although he did not know the dual composition of acids such as hydrochloric

acid (HCl). Over the years, much more accurate definitions of acids and bases have been created.

Arrhenius TheoryThe Swedish chemist Svante Arrhenius published his theory of acids and bases in 1887. It can be simply

explained by these two points:

Arrhenius Acids and Bases1.  An acid is a substance which dissociates in water to produce one or more hydrogen ions (H+).

2.  A base is a substance which dissociates in water to produce one or more hydroxide ions (OH-

).

Svante Arrhenius

Based on this definition, you can see that Arrhenius acids must be soluble in water. Arrhenius acid-base

reactions can be summarized with three generic equations:

An acid will dissociate in water producinghydrogen ions.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 6/31

 

A base (usually containing a metal) will dissociate

in water to product hydroxide ions.

Acids and bases will neutralize each other when

mixed. They produce water and an ionic salt,

neither of which are acidic or basic.

The Arrhenius theory is simple and useful. It explains many properties and reactions of acids and bases.For instance, mixing hydrochloric acid (HCl) with sodium hydroxide (NaOH) results in a neutral solution

containing table salt (NaCl).

However, the Arrhenius theory is not without flaws. There are many well known bases, such as ammonia(NH3) that do not contain the hydroxide ion. Furthermore, acid-base reactions are observed in solutions

that do not contain water. To resolve these problems, there is a more advanced acid-base theory.

Brønsted-Lowry TheoryThe Brønsted-Lowry theory was proposed in 1923. It is more general than the Arrhenius theory — all

Arrhenius acids/bases are also Brønsted-Lowry acids/bases (but not necessarily vice versa).Brønsted-Lowry Acids and Bases

1.  An acid is a substance from which a proton (H+ ion) can be removed. Essentially, an acid

donates protons to bases.2.  A base is a substance to which a proton (H+) can be added. Essentially, a base accepts 

protons from acids.

Acids that can donate only one proton are monoprotic, and acids that can donate more than one proton

are polyprotic.

These reactions demonstrate the behavior of Brønsted-Lowry acids and bases:

An acid (in this case, hydrochloric acid) will donate

a proton to a base (in this case, water is the base).

The acid loses its proton and the base gains it.

Water is not necessary. In this case, hydrochloric

acid is still the acid, but ammonia acts as the base.

The same reaction is happening, but now in reverse.

What was once an acid is now a base (HCl → Cl-)and what was once a base is now an acid (NH3 →NH4

+). This concept is called conjugates, and it will

be explained in more detail later.

Two examples of acids (HCl and H3O+) mixing with

bases (NaOH and OH-) to form neutral substances

(NaCl and H2O).

A base (sodium hydroxide) will accept a proton

from an acid (ammonia). A neutral substance isproduced (water), which is not necessarily a part of 

every reaction. Compare this reaction to the second

one. Ammonia was a base, and now it is an acid.This concept, called amphoterism, is explained later.

The Brønsted-Lowry theory is by far the most useful and commonly-used definition. For the remainder of General Chemistry, you can assume that any acids/bases use the Brønsted-Lowry definition, unless stated

otherwise.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 7/31

This Brønsted-Lowry acid donates a proton (in green) to water (the base).

Lewis TheoryThe Lewis definition is the most general theory, having no requirements for solubility or protons.

Lewis Acids and Bases

1.  An acid is a substance that accepts a lone pair of electrons.

2.  A base is a substance that donates a lone pair electrons.Lewis acids and bases react to create an adduct, a compound in which the acid and base have bonded by

sharing the electron pair. Lewis acid/base reactions are different from redox reactions because there is no

change in oxidation state.

This reaction shows a Lewis base (NH3) donating an electron pair to a Lewis acid (H+) to form an adduct

(NH4+).

[edit] Amphoterism and WaterSubstances capable of acting as either an acid or a base are amphoteric. Water is the most important

amphoteric substance. It can ionize into hydroxide (OH-, a base) or hydronium (H3O

+, an acid). By doing

so, water is

1.  Increasing the H+ or OH- concentration (Arrhenius),

2.  Donating or accepting a proton (Brønsted-Lowry), and

3.  Accepting or donating an electron pair (Lewis).

;Important A bare proton (H+

ion) cannot exist in water. It will form a hydrogen bond to the

nearest water molecule, creating the hydronium ion (H3O+). Although many equations and

definitions may refer to the "concentration of H+ ions", that is a misleading abbreviation.

Technically, there are no H+ ions, only hydronium (H3O+) ions. Fortunately, the number of 

hydronium ions formed is exactly equal to the number of hydrogen ions, so the two can be used

interchangeably.

H+ ions actually exist as hydronium, H3O+.

Water will dissociate very slightly (which further explains is amphoteric properties).

The presence of hydrogen ions indicates an acid, whereas the presence

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 8/31

of hydroxide ions indicates a base. Being neutral, water dissociates into

both equally.

This equation is more accurate — hydrogen ions do not exist in water

because they bond to form hydronium.

[edit] Ammonia

Another common example of an amphoteric substance is ammonia. Ammonia is a normally a base, but insome reactions it can act like an acid.

Ammonia acts as a base. It accepts a proton to form

ammonium.

Ammonia also acts as an acid. Here, it donates a proton toform amide.

Ammonia's amphoteric properties are not often seen because ammonia typically acts like a base. Water,

on the other hand, is completely neutral, so its acid and base behaviors are both observed commonly.

[edit] Conjugate Acids and BasesIn all the theories, the products of an acid-base reaction are related to the initial reactants of the reaction.For example, in the Brønsted-Lowry theory, this relationship is the difference of a proton between a

reactant and product. Two substances which exhibit this relationship form a conjugate acid-base pair.

Brønsted-Lowry Conjugate Pairs  An acid that has donated its proton becomes a conjugate base.

  A base that has accepted a proton becomes a conjugate acid .

Hydroiodic acid reacts with water (which

serves as a base). The conjugate base is the

iodide ion and the conjugate acid is the

hydronium ion. The acids are written in red,

and the bases are written in blue. Oneconjugate pair is written bold and the other

conjugate pair is in cursive.

Ammonia (basic) reacts with water (the acid).

The conjugate acid is ammonium and theconjugate base is hydroxide. Again, acids are

written in red, and the bases are written in blue.

The conjugate pairs are distinguished with

matching fonts.

[edit] Strong and Weak Acids/BasesA strong acid is an acid which dissociates completely in water. That is, all the acid molecules break up

into ions and solvate (attach) to water molecules. Therefore, the concentration of hydronium ions in a

strong acid solution is equal to the concentration of the acid.

The majority of acids exist as weak acids, an acid which dissociates only partially. On average, only

about 1% of a weak acid solution dissociates in water in a 0.1 mol/L solution. Therefore, the

concentration of hydronium ions in a weak acid solution is always less than the concentration of the

dissolved acid.Strong bases and weak bases do not require additional explanation; the concept is the same.

The conjugate of a strong acid/base is very weak. The conjugate of a weak acid/base is not

necessarily strong.

This explains why, in all of the above example reactions, the reverse chemical reaction does not occur.The stronger acid/base will prevail, and the weaker one will not contribute to the overall acidity/basicity.

For example, hydrochloric acid is strong, and upon dissociation chloride ions are formed. Chloride ions

are a weak base, but the solution is not basic because the acidity of HCl is overwhelmingly stronger than

basicity of Cl-.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 9/31

Helpful Hint!

Although the other halogens make strong acids,

hydrofluoric acid (HF) is a weak acid. Despite

being weak, it is incredibly corrosive — hydrofluoric acid dissolves glass and metal!

Most acids and bases are weak. You should be familiar with the most common strong acids and assumethat any other acids are weak.

Formula Strong Acid

HClO4 Perchloric acid

HNO3 Nitric acid

H2SO4 Sulfuric acid

HCl, HBr, HI Hydrohalic acids

Within a series of oxyacids, the ions with the greatest number of oxygen molecules are the strongest. For

example, nitric acid (HNO3) is strong, but nitrous acid (HNO2) is weak. Perchloric acid (HClO4) is

stronger than chloric acid (HClO3), which is stronger than the weak chlorous acid (HClO2). Hypochlorous

acid (HClO) is the weakest of the four.

Common strong bases are the hydroxides of Group 1 and most Group 2 metals. For example, potassium

hydroxide and calcium hydroxide are some of the strongest bases. You can assume that any other bases

(including ammonia and ammonium hydroxide) are weak.

Formula Strong Base

LiOH Lithium hydroxide

NaOH Sodium hydroxide

KOH Potassium hydroxide

RbOH Rubidium hydroxide

CsOH Cesium hydroxide

Ca(OH)2 Calcium hydroxideSr(OH)2 Strontium hydroxide

Ba(OH)2 Barium hydroxide[1] 

Acids and bases that are strong are not necessarily concentrated, and weak acids/bases are not

necessarily dilute. Concentration has nothing to do with the ability of a substance to dissociate.Furthermore, polyprotic acids are not necessarily stronger than monoprotic acids.

[edit] Properties of Acids and BasesNow that you are aware of the acid-base theories, you can learn about the physical and chemical

properties of acids and bases. Acids and bases have very different properties, allowing them to be

distinguished by observation.

[edit] Indicators

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 10/31

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 11/31

 

Acids react with metal

carbonates to produce water,

CO2 gas bubbles, and a salt.

Acids react with metal

oxides to produce water and

a salt.

[edit] BasesBases are typically less reactive and violent than acids. They do still undergo many chemical reactions,

especially with organic compounds. A common reactions is saponificiation: the reaction of a base with fator oil to create soap.

Saponification converts an "ester" into an "alcohol" and salt. This is an organic reaction outside the scope

of General Chemistry.

The Arrhenius Theory of acids and bases 

The theory   Acids are substances which produce hydrogen ions in solution.

  Bases are substances which produce hydroxide ions in solution.

Neutralisation happens because hydrogen ions and hydroxide ions react to produce water.

Limitations of the theory Hydrochloric acid is neutralised by both sodium hydroxide solution and ammonia solution. In both cases,

you get a colourless solution which you can crystallise to get a white salt - either sodium chloride or

ammonium chloride.

These are clearly very similar reactions. The full equations are:

In the sodium hydroxide case, hydrogen ions from the acid are reacting with hydroxide ions from thesodium hydroxide - in line with the Arrhenius theory.

However, in the ammonia case, there don't appear to be any hydroxide ions!

You can get around this by saying that the ammonia reacts with the water it is dissolved in to produce

ammonium ions and hydroxide ions:

This is a reversible reaction, and in a typical dilute ammonia solution, about 99% of the ammonia remains

as ammonia molecules. Nevertheless, there are hydroxide ions there, and we can squeeze this into the

Arrhenius theory.

However, this same reaction also happens between ammonia gas and hydrogen chloride gas.

In this case, there aren't any hydrogen ions or hydroxide ions in solution - because there isn't any solution.

The Arrhenius theory wouldn't count this as an acid-base reaction, despite the fact that it is producing thesame product as when the two substances were in solution. That's silly!

The Bronsted-Lowry Theory of acids and bases 

The theory 

  An acid is a proton (hydrogen ion) donor.

  A base is a proton (hydrogen ion) acceptor.

The relationship between the Bronsted-Lowry theory and the Arrhenius theory 

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 12/31

The Bronsted-Lowry theory doesn't go against the Arrhenius theory in any way - it just adds to it.

Hydroxide ions are still bases because they accept hydrogen ions from acids and form water.

An acid produces hydrogen ions in solution because it reacts with the water molecules by giving a proton

to them.

When hydrogen chloride gas dissolves in water to produce hydrochloric acid, the hydrogen chloride

molecule gives a proton (a hydrogen ion) to a water molecule. A co-ordinate (dative covalent) bond is

formed between one of the lone pairs on the oxygen and the hydrogen from the HCl. Hydroxonium ions,H3O

+, are produced.

Note: If you aren't sure about co-ordinate bonding you should follow this link. Co-ordinate bonds will be

mentioned several times over the course of the rest of this page.Use the BACK button on your browser to return quickly to this page.

When an acid in solution reacts with a base, what is actually functioning as the acid is the hydroxonium

ion. For example, a proton is transferred from a hydroxonium ion to a hydroxide ion to make water.

Showing the electrons, but leaving out the inner ones:

It is important to realise that whenever you talk about hydrogen ions in solution, H+

(aq), what you areactually talking about are hydroxonium ions.

The hydrogen chloride / ammonia problem This is no longer a problem using the Bronsted-Lowry theory. Whether you are talking about the reaction

in solution or in the gas state, ammonia is a base because it accepts a proton (a hydrogen ion). The

hydrogen becomes attached to the lone pair on the nitrogen of the ammonia via a co-ordinate bond.

If it is in solution, the ammonia accepts a proton from a hydroxonium ion:

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 13/31

If the reaction is happening in the gas state, the ammonia accepts a proton directly from the hydrogen

chloride:

Either way, the ammonia acts as a base by accepting a hydrogen ion from an acid.

Conjugate pairs 

When hydrogen chloride dissolves in water, almost 100% of it reacts with the water to producehydroxonium ions and chloride ions. Hydrogen chloride is a strong acid, and we tend to write this as a

one-way reaction:

Note: I am deliberately missing state symbols off this and the next equation in order to concentrate on

the bits that matter.

You will find more about strong and weak acids on another page in this section.

In fact, the reaction between HCl and water is reversible, but only to a very minor extent. In order to

generalise, consider an acid HA, and think of the reaction as being reversible.

Thinking about the forward reaction:  

  The HA is an acid because it is donating a proton (hydrogen ion) to the water.

  The water is a base because it is accepting a proton from the HA.

But there is also a back reaction between the hydroxonium ion and the A- ion:

  The H3O+

is an acid because it is donating a proton (hydrogen ion) to the A-ion.

  The A-ion is a base because it is accepting a proton from the H3O

+.

The reversible reaction contains two acids and two bases. We think of them in pairs, called conjugate

 pairs.

When the acid, HA, loses a proton it forms a base, A-. When the base, A-, accepts a proton back again, it

obviously refoms the acid, HA. These two are a conjugate pair.

 Members of a conjugate pair differ from each other by the presence or absence of the transferable

hydrogen ion. If you are thinking about HA as the acid, then A- is its conjugate base.

If you are thinking about A- as the base, then HA is its conjugate acid.

The water and the hydroxonium ion are also a conjugate pair. Thinking of the water as a base, the

hydroxonium ion is its conjugate acid because it has the extra hydrogen ion which it can give away again.

Thinking about the hydroxonium ion as an acid, then water is its conjugate base. The water can accept a

hydrogen ion back again to reform the hydroxonium ion. A second example of conjugate pairs This is the reaction between ammonia and water that we looked at earlier:

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 14/31

Think first about the forward reaction. Ammonia is a base because it is accepting hydrogen ions from the

water. The ammonium ion is its conjugate acid - it can release that hydrogen ion again to reform the

ammonia.

The water is acting as an acid, and its conjugate base is the hydroxide ion. The hydroxide ion can accept a

hydrogen ion to reform the water.

Looking at it from the other side, the ammonium ion is an acid, and ammonia is its conjugate base. The

hydroxide ion is a base and water is its conjugate acid.Amphoteric substances You may possibly have noticed (although probably not!) that in one of the last two examples, water was

acting as a base, whereas in the other one it was acting as an acid.

A substance which can act as either an acid or a base is described as being amphoteric.

Note: You might also come across the term amphiprotic in this context. The two words are related and

easily confused.

An amphiprotic substance is one which can both donate hydrogen ions (protons) and also accept them.

Water is a good example of such a compound. The water acts as both an acid (donating hydrogen ions)and as a base (by accepting them). The "protic" part of the word refers to the hydrogen ions (protons)

either being donated or accepted. Other examples of amphiprotic compounds are amino acids, and ions

like HSO4-(which can lose a hydrogen ion to form sulphate ions or accept one to form sulphuric acid).

But as well as being amphiprotic, these compounds are also amphoteric. Amphoteric means that they

have reactions as both acids and bases. So what is the difference between the two terms?

All amphiprotic substances are also amphoteric - but the reverse isn't true. There are amphoteric

substances which don't either donate or accept hydrogen ions when they act as acids or bases. There is a

whole new definition of acid-base behaviour that you are just about to meet (the Lewis theory) which

doesn't necessarily involve hydrogen ions at all.

A Lewis acid is an electron pair acceptor; a Lewis base is an electron pair donor (see below).

Some metal oxides (like aluminium oxide) are amphoteric - they react both as acids and bases. For

example, they react as bases because the oxide ions accept hydrogen ions to make water. That's not aproblem as far as the definition of amphiprotic is concerned - but the reaction as an acid is. The

aluminium oxide doesn't contain any hydrogen ions to donate! But aluminium oxide reacts with bases like

sodium hydroxide solution to form complex aluminate ions.

You can think of lone pairs on hydroxide ions as forming dative covalent (coordinate) bonds with emptyorbitals in the aluminium ions. The aluminium ions are accepting lone pairs (acting as a Lewis acid). So

aluminium oxide can act as both an acid and a base - and so is amphoteric. But it isn't amphiprotic

because both of the acid reaction and the base reaction don't involve hydrogen ions.

I have gone through 40-odd years of teaching (in the lab, and via books and the internet) without once

using the term amphiprotic! I simply don't see the point of it. The term amphoteric takes in all the cases of substances functioning as both acids and bases without exception. The term amphiprotic can only be used

where both of these functions involve transference of hydrogen ions - in other words, it can only be used

if you are limited to talking about the Bronsted-Lowry theory. Personally, I would stick to the older, moreuseful, term "amphoteric" unless your syllabus demands that you use the word "amphiprotic".

The Lewis Theory of acids and bases This theory extends well beyond the things you normally think of as acids and bases.

The theory   An acid is an electron pair acceptor.

  A base is an electron pair donor.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 15/31

 

The relationship between the Lewis theory and the Bronsted-Lowry theory 

 Lewis bases It is easiest to see the relationship by looking at exactly what Bronsted-Lowry bases do when they accept

hydrogen ions. Three Bronsted-Lowry bases we've looked at are hydroxide ions, ammonia and water, and

they are typical of all the rest.

The Bronsted-Lowry theory says that they are acting as bases because they are combining with hydrogen

ions. The reason they are combining with hydrogen ions is that they have lone pairs of electrons - which

is what the Lewis theory says. The two are entirely consistent.So how does this extend the concept of a base? At the moment it doesn't - it just looks at it from a

different angle.

But what about other similar reactions of ammonia or water, for example? On the Lewis theory, any 

reaction in which the ammonia or water used their lone pairs of electrons to form a co-ordinate bond

would be counted as them acting as a base.

Here is a reaction which you will find talked about on the page dealing with co-ordinate bonding.Ammonia reacts with BF3 by using its lone pair to form a co-ordinate bond with the empty orbital on the

boron.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 16/31

 As far as the ammonia is concerned, it is behaving exactly the same as when it reacts with a hydrogen ion

- it is using its lone pair to form a co-ordinate bond. If you are going to describe it as a base in one case, it

makes sense to describe it as one in the other case as well.

Note: If you haven't already read the page about co-ordinate bonding you should do so now. You will

find an important example of water acting as a Lewis base as well as this example - although the term

 Lewis base isn't used on that page.

Use the BACK button on your browser to return quickly to this page.

 Lewis acids Lewis acids are electron pair acceptors. In the above example, the BF3 is acting as the Lewis acid by

accepting the nitrogen's lone pair. On the Bronsted-Lowry theory, the BF3 has nothing remotely acidic

about it.

This is an extension of the term acid well beyond any common use.

What about more obviously acid-base reactions - like, for example, the reaction between ammonia and

hydrogen chloride gas?

What exactly is accepting the lone pair of electrons on the nitrogen. Textbooks often write this as if the

ammonia is donating its lone pair to a hydrogen ion - a simple proton with no electrons around it.That is misleading! You don't usually get free hydrogen ions in chemical systems. They are so reactivethat they are always attached to something else. There aren't any uncombined hydrogen ions in HCl.

There isn't an empty orbital anywhere on the HCl which can accept a pair of electrons. Why, then, is the

HCl a Lewis acid?

Chlorine is more electronegative than hydrogen, and that means that the hydrogen chloride will be a polar

molecule. The electrons in the hydrogen-chlorine bond will be attracted towards the chlorine end, leavingthe hydrogen slightly positive and the chlorine slightly negative.

Note: If you aren't sure about electronegativity and bond polarity it might be useful to follow this link.

Use the BACK button on your browser to return quickly to this page.

The lone pair on the nitrogen of an ammonia molecule is attracted to the slightly positive hydrogen atom

in the HCl. As it approaches it, the electrons in the hydrogen-chlorine bond are repelled still further

towards the chlorine.

Eventually, a co-ordinate bond is formed between the nitrogen and the hydrogen, and the chlorine breaks

away as a chloride ion.

This is best shown using the "curly arrow" notation commonly used in organic reaction mechanisms.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 17/31

 

The whole HCl molecule is acting as a Lewis acid. It is accepting a pair of electrons from the ammonia,

and in the process it breaks up. Lewis acids don't necessarily have to have an existing empty orbital.

You can calculate the pH of a buffer solution or the concentration of the acid and base using the

Henderson Hasselbalch equation. Here's a look at the Henderson Hasselbalch equation and a workedexample that explains how to apply the equation.

The Lewis Definitions of Acids and Bases

In 1923 G. N. Lewis suggested another way of looking at the reaction between H+

and OH-ions. In the

Brnsted model, the OH-ion is the active species in this reaction it accepts an H

+ion to form a covalent

bond. In the Lewis model, the H+ ion is the active species it accepts a pair of electrons from the OH- ion to form a covalent bond.

In the Lewis theory of acid-base reactions, bases donate pairs of electrons and acids accept pairs of 

electrons. A Lewis acid is therefore any substance, such as the H+ ion, that can accept a pair of 

nonbonding electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis base is any

substance, such as the OH-ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore

an electron-pair donor.

One advantage of the Lewis theory is the way it complements the model of oxidation-reduction reactions.

Oxidation-reduction reactions involve a transfer of electrons from one atom to another, with a net change

in the oxidation number of one or more atoms.

The Lewis theory suggests that acids react with bases to share a pair of electrons, with no change in the

oxidation numbers of any atoms. Many chemical reactions can be sorted into one or the other of theseclasses. Either electrons are transferred from one atom to another, or the atoms come together to share a

pair of electrons.

The principal advantage of the Lewis theory is the way it expands the number of acids and therefore the

number of acid-base reactions. In the Lewis theory, an acid is any ion or molecule that can accept a pair

of nonbonding valence electrons. In the preceding section, we concluded that Al3+

ions form bonds to six

water molecules to give a complex ion.

Al3+

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

(aq)

This is an example of a Lewis acid-base reaction. The Lewis structure of water suggests that this moleculehas nonbonding pairs of valence electrons and can therefore act as a Lewis base. The electron

configuration of the Al3+

ion suggests that this ion has empty 3s, 3 p, and 3d orbitals that can be used to

hold pairs of nonbonding electrons donated by neighboring water molecules.

Al3+

= [Ne] 3s0

3 p0

3d 0 

Thus, the Al(H2O)63+ ion is formed when an Al3+ ion acting as a Lewis acid picks up six pairs of electrons

from neighboring water molecules acting as Lewis bases to give an acid-base complex, or complex ion.

The Lewis acid-base theroy explains why BF3 reacts with ammonia. BF3 is a trigonal-planar molecule

because electrons can be found in only three places in the valence shell of the boron atom. As a result, the

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 18/31

boron atom is sp2 hybridized, which leaves an empty 2 p z orbital on the boron atom. BF3 can therefore act

as an electron-pair acceptor, or Lewis acid. It can use the empty 2 p z orbital to pick up a pair of 

nonbonding electrons from a Lewis base to form a covalent bond. BF3 therefore reacts with Lewis bases

such as NH3 to form acid-base complexes in which all of the atoms have a filled shell of valence

electrons, as shown in the figure below.

The Lewis acid-base theory can also be used to explain why nonmetal oxides such as CO2 dissolve in

water to form acids, such as carbonic acid H2CO3.

CO2(g) + H2O(l) H2CO3(aq)In the course of this reaction, the water molecule acts as an electron-pair donor, or Lewis base. The

electron-pair acceptor is the carbon atom in CO2. When the carbon atom picks up a pair of electrons from

the water molecule, it no longer needs to form double bonds with both of the other oxygen atoms as

shown in the figure below

One of the oxygen atoms in the intermediate formed when water is added to CO2 carries a positive

charge; another carries a negative charge. After an H+ ion has been transferred from one of these oxygenatoms to the other, all of the oxygen atoms in the compound are electrically neutral. The net result of the

reaction between CO2 and water is therefore carbonic acid, H2CO3.

Dissociation constantFrom Wikipedia, the free encyclopedia"Kd" redirects here. For other uses, see KD (disambiguation). 

In chemistry, biochemistry, and pharmacology, a dissociation constant is a specific type of equilibrium

constant that measures the propensity of a larger object to separate (dissociate) reversibly into smaller

components, as when a complex falls apart into its component molecules, or when a salt splits up into its

component ions. The dissociation constant is usually denoted K d and is the inverse of the association

constant. In the special case of salts, the dissociation constant can also be called an ionization constant. 

For a general reaction

in which a complex A xB y breaks down into x A subunits and y B subunits, the dissociation constant is

defined

where [A], [B], and [AxBy] are the concentrations of A, B, and the complex AxBy, respectively.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 19/31

One reason for the popularity of the dissociation constant in biochemistry and pharmacology is that in the

frequently encountered case where x=y=1, Kd has a simple physical interpretation: when [A]=Kd,

[B]=[AB] or equivalently [AB]/([B]+[AB])=1/2. That is, Kd, which has the dimensions of concentration,

equals the concentration of free A at which half of the total molecules of B are associated with A. This

simple interpretation does not apply for higher values of x or y.

Henderson-Hasselbalch EquationpH = pKa + log ([A

-]/[HA]) 

[A-] = molar concentration of a conjugate base

[HA] = molar concentration of a undissociated weak acid (M)

The equation can be rewritten to solve for pOH:

pOH = pKb + log ([HB+]/[ B ]) 

[HB+] = molar concentration of the conjugate base (M)

[ B ] = molar concentration of a weak base (M)

Example Problem Applying the Henderson-Hasselbalch EquationCalculate the pH of a buffer solution made from 0.20 M HC2H3O2 and 0.50 M C2H3O2

- that has an acid

dissociation constant for HC2H3O2 of 1.8 x 10-5

.

Solve this problem by plugging the values into the Henderson-Hasselbalch equation for a weak acid and

its conjugate base.pH = pKa + log ([A-]/[HA])

pH = pKa + log ([C2H3O2-] / [HC2H3O2])

pH = -log (1.8 x 10-5

) + log (0.50 M / 0.20 M)

pH = -log (1.8 x 10-5

) + log (2.5)pH = 4.7 + 0.40

pH = 5.1

Henderson – Hasselbalch equationFrom Wikipedia, the free encyclopedia

In chemistry, the Henderson – Hasselbalch equation describes the derivation of pH as a measure of 

acidity (using pKa, the acid dissociation constant) in biological and chemical systems. The equation is

also useful for estimating the pH of a buffer solution and finding the equilibrium pH in acid-base

reactions (it is widely used to calculate the isoelectric point of proteins).Two equivalent forms of the equation are

and

Here, pKa is − log(K a) where K a is the acid dissociation constant, that is:

for the non-specific Brønsted acid-

base reaction:

In these equations, A−

denotes the ionic form of the relevant acid. Bracketed quantities such as [base] and

[acid] denote the molar concentration of the quantity enclosed.

In analogy to the above equations, the following equation is valid:

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 20/31

Where BH+ denotes the conjugate acid of the corresponding base B.

DerivationThe Henderson – Hasselbalch equation is derived from the acid dissociation constant equation by the

following steps:

The ratio [ A −

] / [ HA] is unitless, and as such, other ratios with other units may be used. For example, the

mole ratio of the components, or the fractional concentrations where

will yield the same answer. Sometimes these other units are more convenient to

use.

[edit] HistoryLawrence Joseph Henderson wrote an equation, in 1908, describing the use of carbonic acid as a buffer

solution. Karl Albert Hasselbalch later re-expressed that formula in logarithmic terms, resulting in the

Henderson – Hasselbalch equation [1]. Hasselbalch was using the formula to study metabolic acidosis. 

[edit] Limitations

There are some significant approximations implicit in the Henderson – Hasselbalch equation. The mostsignificant is the assumption that the concentration of the acid and its conjugate base at equilibrium will

remain the same as the formal concentration. This neglects the dissociation of the acid and the hydrolysis

of the base. The dissociation of water itself is neglected as well. These approximations will fail when

dealing with relatively strong acids or bases (pKa more than a couple units away from 7), dilute or very

concentrated solutions (less than 1 mM or greater than 1M), or heavily skewed acid/base ratios (more

than 100 to 1). Also, the equation does not take into effect the dilution factor of the acid and conjugate

base in water. If the proportion of acid to base is 1, then the pH of the solution will be different if the

amount of water changes from 1mL to 1L.

[edit] Estimating blood pHA modified version of the Henderson – Hasselbalch equation can be used to relate the pH of blood to

constituents of the bicarbonate buffering system:[1] 

, where:

  pKa H2CO3 is the acid dissociation constant of carbonic acid. It is equal to 6.1.

  [HCO3-] is the concentration of bicarbonate in the blood

  [H2CO3] is the concentration of carbonic acid in the blood

This is useful in arterial blood gas, but these usually state pCO2, that is, the partial pressure of carbon

dioxide, rather than H2CO3. However, these are related by the equation:[1]

 

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 21/31

 , where:

  [H2CO3] is the concentration of carbonic acid in the blood

  k  H CO2 is a constant including the solubility of carbon dioxide in blood. k  H CO2 is approximately

0.03 mmol /mmHg

   pCO2 is the partial pressure of carbon dioxide in the blood

Taken together, the following equation can be used to relate the pH of blood to the concentration of bicarbonate and the partial pressure of carbon dioxide:[1] 

, where:

  pH is the acidity in the blood

  [HCO3-] is the concentration of bicarbonate in the blood

   pCO2 is the partial pressure of carbon dioxide in the blood

defining & Using pH & pOH

25

o

C Acid neutral BasepH  0 1 2 3 4 5 6 7 8 9 10 11 12 13 14  

 E

 x

 a

m

 p

l 

e

 s 

1M

HCl

0.1M

HCl

Gastric

 juice,

antvenom

coca

cola,

lemon juice,

vinegar

wine

coffee,

toma-toes

tap

water,

saliva,cow's

milk 

pure

water,

NaCl(aq),KNO3(aq)

sea

water,

soap,

bakingsoda

deter-

gents,

tooth-paste

deter-

gents,

washingsoda

house-

holdcleaner

0.1M

NaOH,

causticoven

cleaner

1M

NaOH

OH  14 13 12 11 10 9 8 7 6 5 4 3 2 1 0  p

mostacidic

<  ---- ----- ---- --------- leastacidic

neutral leastbasic

----- -------- ----- ----- >  mostbasic

pH pOH

pH is a measure of the hydrogen ion concentration,

[H+]

pOH is a measure of the hydroxide ion

concentration, [OH-]

pH is calculated using the following formula:

pH = -log10[H+]

pOH is calculated using the following formula:

pOH = -log10[OH-]

Example 1:Find the pH of a 0.2mol L

-1 

(0.2M) solution of HCl

  Write the balanced equation for the

dissociation of the acid

HCl -----> H+(aq) + Cl

-(aq)

  Use the equation to find the [H+]:

0.2 mol L- HCl produces 0.2 mol L-1 H+ 

since HCl is a strong acid that fully

dissociates

Example 1:Find the pOH of a 0.1mol L

- 

(0.1M) solution of NaOH

  Write the balanced equation for the

dissociation of the alkali

NaOH -----> OH-(aq) + Na

+(aq)

  Use the equation to find the [OH-]:

0.1 mol L-1 NaOH produces 0.1 mol L-1 OH- 

since NaOH is a strong alkali that fully

dissociates

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 22/31

  Calculate pH: pH = -log10[H+]

pH = -log10[0.2] = 0.7  Calculate pOH: pOH = -log10[OH-]

pOH = -log10[0.1] = 1

Example 2:Find the pH of a 0.2 mol L

-1 

(0.2M) solution of H2SO4 

  Write the balanced equation for the

dissociation of the acid

H2SO4 -----> 2H+(aq) + SO42-(aq)

  Use the equation to find the [H+]:

0.2 mol L-1

H2SO4 produces 2 x 0.2 = 0.4

mol L-1

H+

since H2SO4 is a strong acidthat fully dissociates

  Calculate pH: pH = -log10[H+]

pH = -log10[0.4] = 0.4

Example 2:Find the pOH of a 0.1mol L

-1 

(0.1M) solution of Ba(OH)2 

  Write the balanced equation for the

dissociation of the alkali:

Ba(OH)2 -----> 2OH-(aq) + Ba2+(aq)

  Use the equation to find the [OH-]:

0.1mol L-1

Ba(OH)2 produces 2 x 0.1 = 0.2

mol L-1

OH-since Ba(OH)2 is a strong

alkali that fully dissociates

  Calculate pOH: pOH = -log10[OH-]

pOH = -log10[0.2] = 0.7

Hydrogen ion concentration, [H+], can be

calculated using the following formula:

[H+] = 10-pH 

Hydroxide ion concentration, [OH-], can be

calculated using the following formula:

[OH-] = 10-pOH 

Example:Find the [H

+] of a nitric acid solution with a pH of 

3.0

pH= 3.0

[H+] = 10-pH 

[H+] = 10

-3.0= 0.001mol L

-1 

You can check this answer by using the calculated

value [H+] in the equation for pH to make sure you

arrive at the original pH

pH = -log10[H+]

pH = -log10[0.001] = 3We get the same value for pH using the calculated

value for [H+], so the calculated value for [H

+] is

correct.

Example:Find the [OH

-] of a sodium hydroxide solution with

a pOH of 1

pOH = 1

[OH-] = 10-pOH 

[OH-] = 10

-1= 0.1 mol L

-1 

You can check this answer by using the calculated

value [OH-] in the equation for pOH to make sure

you arrive at the original pOH

pOH = -log10[OH-]

pOH = -log10[0.1] = 1We get the same value for pOH using the calculated

value for [OH-], so the calculated value for [OH

-] is

correct.

pH + pOH = 14 (25oC)

Example A(1):Find the pH of a solution of sodium hydroxide

that has a pOH of 2

pH = 14 - pOH

pH = 14 - 2 = 12

Example B(1):Find the pOH of a solution of hydrochloric acid

that has a pH of 3.4

pOH = 14 - pH

pOH = 14 - 3.4 = 10.6

Example A(2):Find the [H

+] in a solution of sodium hydroxide

that has a pOH of 1

  Calculate the pH

pH = 14 - pOH

pH = 14 - 1 = 13

  Calculate [H+]

[H+] = 10-pH 

Example B(2):Find the [OH

-] of a sulfuric acid solution

with a pH of 3

  Calculate the pOH

pOH = 14 - pH

pOH = 14 - 3 = 11

  Calculate [OH-]

[OH-] = 10-pOH 

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 23/31

[H+] = 10-13 = 10-13mol L-1 [OH-] = 10-11 = 10-11 mol L-1 

Example A(3):Find the pH of 0.2mol L-1 sodium hydroxide

  Write the equation for the dissociation of 

NaOH:NaOH -----> Na+(aq) + OH-(aq)

  Use the equation to find [OH-]:

0.2mol L-1 NaOH produces 0.2mol L-1 

OH-since NaOH is a strong base that fully

dissociates

  Calculate the pOH:pOH = -log10[OH-]

pOH = -log10[0.2] = 0.7

  Calculate pH:

pH = 14 - pOH

pH = 14 - 0.7 = 13.3

Example B(3):Find the pOH of 0.2mol L-1 sulfuric acid

  Write the equation for the dissociation of 

H2SO4:H2SO4 -----> 2H+(aq) + SO4

2-(aq)

  Use the equation to find [H+]:

0.2 mol L-1 H2SO4 produces 2 x 0.2 = 0.4

mol L-1

H+

since H2SO4 is a strong acid that

fully dissociates

  Calculate the pH:pH = -log10[H

+]:

pH = -log10[0.4] = 0.4

  Calculate pOH:

pOH = 14 - pH

pOH = 14 - 0.4 = 13.6

pH can be viewed as an abbreviation for power of Hydrogen - or more completely, power of theconcentration of the Hydrogen ion.

The mathematical definition of pH is a bit less intuitive but in general more useful. It says that the pH is

equal to to the negative logarithmic value of the Hydrogen ion (H+) concentration, or

 pH = -log [H +] 

pH can alternatively be defined mathematically as the negative logarithmic value of the Hydroxonium

ion (H3O+) concentration. Using the Bronsted-Lowry approach

 pH = -log [H 3O+] 

pH values are calculated in powers of 10. The hydrogen ion concentration of a solution with pH 1.0 is 10

times larger than the hydrogen concentration in a solution with pH 2.0. The larger the hydrogen ion

concentration, the smaller the pH.This allows the definition of the following series of quantities.

pOH = -log[OH-]

the negative log of the hydroxide ion molarity

pKw = -log Kw the negative log of the water ion product , Kw

pKa = -log Ka the negative log of the acid dissociation constant, Ka

pKb = -log Kb the negative log of the base dissociation constant, Kb

The relationship pH + pOH = 14In a water solution the ion product for water is:

[H+] [OH

-] = Kw = 1 X 10

-14 

Take the -log of both sides of the equation

- log [H+] +(- log [OH

-]) = - log [1 X 10

-14] 

pH + pOH = 14 

Calculations of pHFor strong acids like HCl the molar concentrations are essentially the hydronium ion concentration. These

strong acids can produce solutions where the pH can be equal to or less than 1, the pH value would have avalue from 0-14.

Example: Determination of pH from [H3O+]

What is the pH of a solution whose [H3O+] = 1 X 10

-4M

pH = -log[H3O+]

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 24/31

pH = - log[1 X 10-4]

pH = - [ log 1 + log 10-4 ]

Note: When you multiply numbers you always ADD their log forms

log 1 is always zero

log 10x

= x so log 10-4

= -4

pH = - [ log 1 + log 10-4 ] = - [ 0 + (-4) ] = - [-4 ] = +4 

Example: Determination of pH from [H3O+] when coefficient is other than "1" 

What is the pH of a solution whose [H3O+] = 2.5 X 10

-5M

pH = -log[H3O+]

pH = - log[2.5 X 10-5

]

pH = - [ log 2.5 + log 10 -5 ]

Note: When you multiply numbers you always ADD their log forms

log 10x = x so log 10-5 = -5

log 2.5 can be determined using a calculator having the log

function key:Enter the number in this case 2.5

depress the log key

Read the display which should be .3979 for this problem

pH = - [.3979 - 5] = 4.6021 or +4.602 

Alternately if you can enter a number in scientific notation into your calculator key in 2.5 X 10-5 

depress the log key

Read the display which should be -4.602 for this problem

Multiply by -1 to get + 4.602

Example: Determination of pH from [OH1- ] using defintion pOH and equation pH + pOH = 14 

Calculate the pH of a solution that has a [OH1-

] = 1 X 10-5

M

Determine pOH = -log[OH1-

] = -log [1 X 10-5

] = 5

Use the relationship pH + pOH = 14

pH + 5 = 14

pH = 14 -5 = 9 

A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or

a weak base and its conjugate acid. It has the property that the pH of the solution changes very little when

a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at

a nearly constant value in a wide variety of chemical applications. Many life forms thrive only in arelatively small pH range; an example of a buffer solution is blood. 

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 25/31

[edit] Principles of Buffering

Simulated titration of an acidified solution of a weak acid (pKa = 4.7) with alkali.

Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between

the acid HA and its conjugate base A-.

HA H+

+ A- 

When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, theequilibrium is shifted to the left, in accordance with Le Chatelier's principle. Because of this, the

hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid

added. Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less

than the amount expected for the quantity of alkali added.

The effect is illustrated by the simulated titration of a weak acid with pKa = 4.7. The relative

concentration of undissociated acid is shown in blue and of its conjugate base in red. The pH changes

relatively slowly in the buffer region, pH = pKa ± 1, centered at pH = 4.7 where [HA] = [A-], but once the

acid is more than 95% deprotonated the pH rises much more rapidly.

[edit] Buffer capacity

Buffer capacity for a 0.1 M solution of an acid with pKa of 7

Buffer capacity, β, is a quantitative measure of the resistance of a buffer solution to pH change onaddition of hydroxide ions. It can be defined as follows.

where dn is an infinitesimal amount of added base and d(p[H+]) is the resulting infinitesimal change in the

cologarithm of the hydrogen ion concentration. With this definition the buffer capacity of a weak acid,with a dissociation constant Ka, can be expressed as

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 26/31

where CA is the analytical concentration of the acid.[1] pH is approximately equal to -log10[H+].

There are three regions of high buffer capacity.

  At very low p[H+] the first term predominates and β increases in proportion to the hydrogen ionconcentration. This is independent of the presence or absence of buffering agents and applies to

all solvents.

  In the region p[H+] = pKa ± 2 the second term becomes important and β rises to a maximum at

p[H+] = pKa. Buffer capacity is proportional to the concentration of the buffering agent, CA, sodilute solutions have little buffer capacity. It is also proportional to the acid dissociation constant,

Ka (not pKa); the weaker the acid the greater its buffering capacity.

  At very high p[H+] the third term predominates and β increases in proportion to the hydroxide ion

concentration. This is due to the self-ionization of water and is independent of the presence or

absence of buffering agents.

[edit] Calculating buffer pH (monoprotic acid)First write down the equilibrium expression.

HA A- + H+ 

The initial, change and equilibrium concentrations of these three components can be organized in an ICE

table. 

ICE table for a monoprotic acid

[HA] [A-] [H+]

I C0 0 0

C -x x x

E C0-x x x

The first row, labelled I, lists the initial conditions: the concentration of acid is C0, initially undissociated,

so the concentrations of A- and H+ are zero. The second row, labelled C for change, specifies the changes

that occur when the acid dissociates. The acid concentration decreases by an amount -x and the

concentrations of A- and H+ both increase by an amount +x. This follows from the equilibrium

expression. The third row, labelled E for equilibrium concentrations, adds together the first two rows and

shows the concentrations at equilibrium.

To find x, use the formula for the equilibrium constant:

Substitute the concentrations with the values found in the last row of the ICE table:

Simplify to:

 x2 + K a x − K aC 0 = 0

With specific values for C0 and Ka this equation can be solved for x. Assuming that pH = -log10[H+] the

pH can be calculated as pH = -log10x.

Note. When a pH meter is calibrated using known buffers, the reading gives the hydrogen ion activity

rather than its concentration. In this case the meter reading may differ from the value calculated as above.

For example, calculation of pH of phosphate-buffered saline would give the value of 7.96, whereas the

meter reading would be 7.4. The discrepancy arises when the acid dissociation constant value is specified

as a concentration quotient and would not occur if Ka were specified as a quotient of activities.

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 27/31

[edit] Calculating buffer pH (polyprotic acid)

% species formation calculated for a 10 millimolar solution of citric acid.

Polyprotic acids are acids that can lose more than one proton. The constant for dissociation of the first

proton may be denoted as K a1 and the constants for dissociation of successive protons as K a2, etc. Citric

acid, H3A, is an example of a polyprotic acid as it can lose three protons.

equilibrium p K a value

H3A H2A−

+ H+ pK a1 = 3.13

H2A−

HA2−

+ H+ pK a2 = 4.76

HA2−

A3−

+ H+ pK a3 = 6.40

When the difference between successive pK values is less than about three there is overlap between the

pH range of existence of the species in equilibrium. The smaller the difference, the more the overlap. In

the case of citric acid, the overlap is extensive and solutions of citric acid are buffered over the whole

range of pH 2.5 to 7.5. Calculation of the pH of a particular mixture requires a speciation calculation to be

performed.

[edit] ApplicationsBuffer solutions are necessary to keep the correct pH for enzymes in many organisms to work. Many

enzymes work only under very precise conditions; if the pH moves outside of a narrow range, theenzymes slow or stop working and can denature. In many cases denaturation can permanently disable

their catalytic activity.[2] A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3−) is present in blood

plasma, to maintain a pH between 7.35 and 7.45.

Industrially, buffer solutions are used in fermentation processes and in setting the correct conditions for

dyes used in colouring fabrics. They are also used in chemical analysis[1] and calibration of pH meters.

The majority of biological samples that are used in research are made in buffers, especially phosphate

buffered saline (PBS) at pH 7.4.

[edit] Useful buffer mixtures

Components pH range

HCl, Sodium citrate 1 - 5

Citric acid, Sodium citrate 2.5 - 5.6Acetic acid, Sodium acetate 3.7 - 5.6

K2HPO4, KH2PO4  5.8 - 8

Na2HPO4, NaH2PO4  6 - 7.5

CHES 8.6 – 10

Borax, Sodium hydroxide 9.2 - 11

[edit] "Universal" buffer mixtures

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 28/31

By combining substances with pK a values differing by only two or less and adjusting the pH a wide-range

of buffers can be obtained. Citric acid is a useful component of a buffer mixture because it has three pK a 

values, separated by less than two. The buffer range can be extended by adding other buffering agents.

The following two-component mixtures (McIlvaine's buffer solutions have a buffer range of pH 3 to 8.[3] 

0.2M Na2HPO4 /mL 0.1M Citric Acid /mL pH...

20.55 79.45 3.038.55 61.45 4.0

51.50 48.50 5.0

63.15 36.85 6.0

82.35 17.65 7.0

97.25 2.75 8.0

A mixture containing citric acid, potassium dihydrogen phosphate, boric acid, and diethyl barbituric acid

can be made to cover the pH range 2.6 to 12.[4]

 

Other universal buffers are Carmody buffer and Britton-Robinson buffer, developed in 1931.

[edit] Common buffer compounds used in biology

Common

Name

pKa at

25°C

Buffer

Range

Temp

Effect d pH/  d T

in (1/K)

**

Mol.

WeightFull Compound Name

TAPS 8.43 7.7 – 9.1 −0.018 243.33-{[tris(hydroxymethyl)methyl]amino}propanesulfonic

acid

Bicine 8.35 7.6 – 9.0 −0.018 163.2 N,N-bis(2-hydroxyethyl)glycine

Tris 8.06 7.5 – 9.0 −0.028 121.14 tris(hydroxymethyl)methylamine

Tricine 8.05 7.4 – 8.8 −0.021 179.2 N-tris(hydroxymethyl)methylglycine

TAPSO 7.635 7.0-8.2 259.33-[N-Tris(hydroxymethyl)methylamino]-2-

hydroxypropanesulfonic Acid

HEPES 7.48 6.8 – 8.2 −0.014 238.3 4-2-hydroxyethyl-1-piperazineethanesulfonic acid

TES 7.40 6.8 – 8.2 −0.020 229.202-{[tris(hydroxymethyl)methyl]amino}ethanesulfonic

acid

MOPS 7.20 6.5 – 7.9 −0.015 209.3 3-(N-morpholino)propanesulfonic acid

PIPES 6.76 6.1 – 7.5 −0.008 302.4 piperazine- N,N′-bis(2-ethanesulfonic acid)

Cacodylate 6.27 5.0 – 7.4 138.0 dimethylarsinic acid

SSC 7.0 6.5-7.5 189.1 saline sodium citrate

MES 6.15 5.5 – 6.7 −0.011 195.2 2-(N-morpholino)ethanesulfonic acid

** Values are approximate. [5]

 

What is a buffer?

A buffer is a molecule that tends to either bind or release hydrogen ions in order to maintain aparticular pH. You have seen that our blood, for example, needs to maintain the pH of 7.4. Buffers help

that occur. In order to explain this, let me use the bicarbonate ion that you are already familiar with as an

example. You have already seen that the bicarbonate ion form reversibly from carbonic acid, and can

reversibly form CO3-, the carbonate ion. Here is that figure again...

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 29/31

You can see that if the reactions go to the right, hydrogen ions are released, making the solution more

acidic, and if the reactions go to the left, hydrogen ions are sucked up, making the solution more basic.

The burning question that is probably eating away at you right now is: "how do I know which way the

reactions are going to go-- to the right or to the left?"

Good question! The bicarbonate ion isn't a very strong acid or base. It doesn't have to go one way or

the other all the time. Instead, the direction it goes depends on the solution it is in. You see, if it is in an

acidic solution, there are lots of hydrogen ions floating around; in this situation, the presence of tons of hydrogen ions will force the reaction to go to the left. Do you see why? Look at the reaction again. All

you need for the reaction to be able to go to the left is available hydrogen ions (see the H+ in the middle

that is needed for the reaction to go to the left?). Hydrogen ions are readily available in acidic solutions.

So if bicarbonate ions find themselves in acidic solutions, they tend to act like the weak base they can be,

and suck up the excess hydrogen ions.

To finish answering the questions, consider what will happen if bicarbonate is put into a basic

solution... A basic solution has very few hydrogen ions floating around. In this condition, the bicarbonate

reaction cannot proceed to the left, since no hydrogen ions are available. Instead, it will go to the right,

producing hydrogen ions. In this manner, bicarbonate ions will tend to act like the weak acid they can be,

and release hydrogen ions.

Now back to the general question about what buffers are. You see, a buffer can either accept or donate

hydrogen ions, depending on the solution they are in. Since the buffers will accept hydrogen ions in acidsand donate hydrogen ions in bases, there must be some in-between-pH where they hit an equilibrium

point and do not prefer to either accept or donate hydrogen ions. Right? That intermediate point, that

equilibrium, is the pH that the buffer tends to maintain. The bicarbonate buffer, so important in blood,

has its equilibrium right at the pH of 7.4. Pretty useful, huh?There are three important buffer systems in our bodies:

1.  bicarbonate buffer system

2.  phosphate buffer system

3.  protein buffer system

All three work similarly-- if they find themselves in a solution with a lot of free hydrogen ions floating

around (an acid), they act as bases and suck up the excess hydrogen ions. And if they find themselves in

a solution lacking free hydrogen ions (a base), they donate their hydrogen ions to the solution. Your book 

says the same thing, but uses different terminology. So your book says that these three all turn strongacids into weak acids and turn strong bases into weak bases. Let me try to explain...

If a buffer finds itself in a solution with a strong acid, what happens? The strong acid gives off its

hydrogen ions. The buffer grabs up many of the free hydrogen ions. In this way, even if 100 molecules

of strong acid were added to the solution, there will not be 100 hydrogen ions floating around. Instead,

there will fewer hydrogen ions floating around, maybe only 20. A yield of only 20 hydrogen ions from

100 molecules of acid is a 20% yield, and that is what you would expect of a weak acid. So, the buffer

changed the strong acid into a weak acid. Does that make sense?

I find it more difficult to think about buffers the way your book does, but you might find it easier... so

for that reason I have explained their description. I hope it makes some sense to you!

Now, it is time to go through the three buffer systems!

The bicarbonate buffer

systemWe have basically

finished this. Quick 

overview: carbon dioxide

is converted into

bicarbonate ions in the

blood. In this way,

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 30/31

bicarbonate ions float around in the blood, and serve to maintain a pH of 7.4 in the blood. I found this

picture on another web site, and have included it here just to have a different image of this process from

the one I have drawn.  Just ignore the stuff in blue. 

The phosphate buffer system

You learned that nucleicacid and phosphoprotein

break down yields

phosphoric acid.

Phosphoric acid is H3PO4,

and is shown at the far left

edge of this figure (also

taken from another web page). Phosphoric acid changes pretty quickly into dihydrogen phosphate, or

H2PO4-. This dihydrogen phosphate is an excellent buffer, since it can either grab up a hydrogen ion and

reform phosphoric acid, or it can give off another hydrogen ion and become monohydrogen phosphate, or

HPO42-

. This figure shows that in extremely basic conditions, monohydrogen phosphate can even give up

its remaining hydrogen ion, although your book doesn't show that.

If the H2PO4- is in an acidic solution, the reactions above go to the left, and it if the H2PO4- is in a basicsolution, the reactions above proceed to the right. Therefore, the phosphate buffer system can accept or

donate hydrogen ions depending on the solution it is in.

The protein buffer systemProteins themselves can act as buffers. You know this, because I told you that in blood, when

bicarbonate ions form, the hydrogen ions that are also products of bicarbonate production are absorbed by

the blood proteins. You also learned that this absorption of the hydrogen ions doesn't happen so well in

cerebrospinal fluid, because there are fewer proteins in CSF.

But how does this work? To understand this, you have

to think back to the structure of a protein. Proteins are

made up of amino acids. And you will have to remember

all about amino acids for this. Amino acids have a centralcarbon with four groups off of it:

1.  a carboxyl group (COOH)

2.  an amino group (NH2)

3.  a hydrogen atom

4.  an R group

The carboxyl and amino groups are what enable proteins to act as buffers. Let me explain.

The carboxyl groupThe carboxyl group is attached to the amino acid central carbon: C - COOH. In the figure above, you

can see the carboxyl group off to the left. You can see that the carboxyl group consists of a double bond

to one of the oxygens and a single bond to the hydroxyl group. The important part of the carboxyl group

for our purposes here is the hydrogen atom within the hydroxyl group.

At a near neutral pH, like the pH of blood, the carboxyl group is actually COO-

instead of COOH.Then, if a protein finds itself in a more acidic solution, the carboxyl group will be able to take on the extra

hydrogen ions and return to the COOH configuration.

Your book describes this by saying that if a protein is in a more basic solution, its carboxyl groups on

its amino acids change from COOH to COO-. This isn't exactly correct, because even at neutral pHs, the

carboxyl group is COO-. But it is true that if the carboxyl group is in the form of COOH, it will become

COO-if it finds itself in a more basic solution. OK?

The amino group

8/2/2019 Biochem Act 1

http://slidepdf.com/reader/full/biochem-act-1 31/31

The amino group is attached to the amino acid central carbon: C - NH2. The amino group is shown at

the right hand side of the diagram of the amino acid above. However, at a near neutral pH, like in blood,

the amino group is actually NH3+ rather than just NH2. It actually tends to carry an extra hydrogen ion on

it at a normal pH. Then, if a protein finds itself in a more basic environment, its amino groups on its

amino acids can actually release their hydrogen ions and return to NH2.

Again, your book describes this a little off... it says that when proteins are in acids, their amino groups

tend to become NH3+. In fact, the amino groups within protein amino acids are NH3+ even at neutralpHs. But when they get into basic pH solutions, they return to NH2.

Diagramatic Overview of amino acids providing buffering functions:

So, amino acids can accept or donate hydrogen ions, making them excellent buffers. And any given

protein typically has hundreds of amino acids. So, proteins make superb buffers. Remember, they are

found in very high concentration in intracellular solutions and in blood.