Acids, Bases and Salts

Why everything tastes so differently? Why lemon is tangy and mango

is sweet? This is because of different percentage of acids, bases, and

salts in their chemical composition. Let’s learn how these can be

characterized, understanding the concepts by Arrhenius,

Bronsted-Lowry and Lewis etc.

Experimental Definitions

Earlier, acids, bases, and salts were characterized by the experimental

testing of their aqueous solutions. An acid is defined as a substance

whose water solution tastes sour, turns blue litmus red and neutralizes

bases. A substance is called base if its aqueous solution tastes bitter,

turns red litmus blue or neutralizes acids.

Salt is a neutral substance whose aqueous solution does not affect

litmus. According to Faraday: acids, bases, and salts are termed as

electrolytes. Further, Liebig proposed that acids are compounds which

contain hydrogen that can be replaced by metals.

Acids

Acidity is a characteristic property of acids. Acidic substances are

usually very sour. Apart from hydrochloric acid, there are many other

types of acids around us. Citrus fruits like lemons and oranges contain

citric and ascorbic acids while tamarind paste contains tartaric acid.

In fact, the word ‘acid’ and ‘acidity’ are derived from the Latin word

‘acidus’ which means sour. If you dip a blue litmus paper into an acid,

it will turn red while a red litmus paper will not change colour. Acids

also liberate dihydrogen when they react with some metals.

Bases

Bases turn red litmus paper blue while the blue litmus paper stays

blue. They taste bitter and also feel soapy. Some other common

examples of bases include sodium bicarbonate that is used in cooking

and household bleach.

Image: Litmus paper test. [Source: Wikimedia Commons]

Salts

Apart from sodium chloride, other common salts are sodium nitrate,

barium sulfate etc. Sodium chloride or common salt is a product of the

reaction between the hydrochloric acid (acid) and sodium hydroxide

(base). Solid sodium chloride is made of a cluster of positively

charged sodium ions and negatively charged chloride ions held

together by electrostatic forces.

Electrostatic forces between opposite charges are inversely

proportional to the dielectric constant of the medium. In other words,

we can say that a compound that has acidity in its nature and a

compound that has basicity as its nature, may yield salts when

combined together.

The universal solvent, water, has a dielectric constant of 80.

Therefore, when sodium chloride is dissolved in water, the dielectric

constant of water reduces the electrostatic force, allowing the ions to

move freely in the solution. They are also well-separated due to

hydration with water molecules.

Image: Dissolution of sodium chloride in water [Source: Wikimedia Commons]

Ionization And Dissociation

Dissociation is the separation of ions from an ionic crystal when a

solid ionic compound dissolves in water. On the other hand, ionization

is the process where a neutral molecule breaks into charged ions when

dissolved in a solution. The extent of ionization depends on the

strength of the bonds between ions and the extent of solvation of ions.

The three most important modern concepts of acids and bases are:

Arrhenius Concept

According to Arrhenius concept, Substances which produce H+ ions

when dissolved in water are called acids while those which ionize in

water to produce OH– ions are called bases.

HA → H+ + A– (Acid)

BOH → B+ + OH– (Base)

Arrhenius proposed that acid-base reactions are characterized by acids

if they dissociate in aqueous solution to form hydrogen ions (H+) and

bases if they form hydroxide (OH–) ions in aqueous solution.

Limitations of Arrhenius Concept

● The presence of water is absolutely necessary for acids and

bases. Dry HCl can’t act as an acid. HCl acts as an acid in

water only and not any other solvent.

● The concept does not explain the acidic and basic character of

substances in non-aqueous solvents.

● The neutralization process is only possible for reactions which

can occur in aqueous solutions, although reactions involving

salt formation can occur in the absence of a solvent.

● The acidic character of some salts such as AlCl3 in aqueous

solution can’t be explained.

● An extended as well as artificial explanation is needed to define

the basic nature of NH3.

Bronsted-Lowry Concept

Bronsted and Lowry in 1923 independently proposed a more general

definition of acids and bases. According to them, an acid is defined as

any hydrogen-containing material (molecule, anion or cation) which

can donate a proton to other substance and a Base is any

substance(molecule, cation or anion) that can accept a proton from any

other substance. Therefore, acids are proton donor whereas bases are

proton acceptor.

Conjugate Acid-Base Pairs

Consider a reaction

Acid1 + Base2 → Acid2 + Base1

H2O + HCl ⇔ H3O+ + Cl–

In this reaction, HCl donates a proton to H2O and is, therefore an acid.

Water, on the other hand, accepts a proton from HCl, and is, therefore,

a base. In the reverse reaction which at equilibrium proceeds at the

same rate as the forward reaction, the H3O+ ions donate a proton to

Cl– ion, hence H3O+, an ion is an acid. Cl– ion, because it accepts a

proton fromH3O+ ion, is a base.

Acid-base pairs in which the members of reaction can be formed from

each other by the gain or loss of protons are called conjugate acid-base

pairs.

Limitations of Bronsted Lowry Concept

● Bronsted Lowry could not explain the reaction occurring in the

non-protonic solvent like COCl3, SO2, N2O4, etc.

● It cannot explain the reactions between acidic oxides like etc

and the basic oxides like etc which can easily take place in the

absence of solvent as well e.g. (No proton transfer)

● Substances like BF3, AlCl3 etc, do not contain hydrogen which

means they can’t donate a proton, still they behave as acids.

Lewis Concept

According to Lewis theory of acid-base reactions, bases donate pairs

of electrons and acids accept pairs of electrons. Thus, it can be said

that a Lewis acid is electron-pair acceptor.

The advantage of the Lewis theory is that complements the model of

oxidation-reduction reactions. Oxidation-reduction reactions take

place on 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 further suggested that acids react with bases and

share a pair of electrons but there is no change in the oxidation

numbers of any atoms. Either an electron is transferred from one atom

to another, or the atoms come together to share a pair of electrons.

Al(OH)3 + 3H+ → Al3+ + 3H2O (Aluminium hydroxide is acting as a

base)

Al(OH)3 + OH– → Al(OH)4- (Aluminium hydroxide is acting as an

acid)

These reactions are showing clearly: When Aluminium hydroxide

accepts protons, it acts as a base. When it accepts electrons, it acts as

an acid. This Lewis acid-base theory also explains why non-metal

oxides such as carbon dioxide dissolve in H2O to form acids, such as

carbonic acid H2CO3.

CO2(g) + H2O(l) → H2CO3(aq)

Limitations of Lewis Concept

● Lewis concept gave a generalized idea including all

coordination reactions and compounds. This is not true always.

● An idea about the relative strength of acids and bases is not

provided by Lewis concept.

● Lewis concept is not in line with the acid-base reaction

concept.

● Lewis concept has not discussed the behaviour of protonic

acids like HCl.

Solved Example for You

Question: Whether the following ions or molecules can act as Lewis

acid or a Lewis base?

● Ag+

● NH3

Solution:

● A silver cation is Lewis acid

● Ammonia is Lewis base

Buffer Solutions

What do you think will happen if the pH of our blood changes

drastically from its normal pH of 7.35? Yes, the cells of our body will

not function properly and our body systems will fail! Human blood

contains a ‘buffer’ that allows it to maintain its pH at 7.35 to ensure

normal functioning of cells. Buffer solutions are also important in

chemical and biochemical processes where the control of pH is very

important. Let’s understand buffer solutions in more detail.

Buffer Solutions

Buffers are solutions that resist a change in pH on dilution or on

addition of small amounts of acids or alkali.

A lot of biological and chemical reactions need a constant pH for the

reaction to proceed. Buffers are extremely useful in these systems to

maintain the pH at a constant value. This does not mean that the pH of

buffers does not change. It only means that the change in pH is not as

much as it would be with a solution that is not a buffer.

Browse more Topics under Equilibrium

● Acids, Bases and Salts

● Equilibrium in Chemical Processes

● Equilibrium in Physical Processes

● Factors Affecting Equilibria

● Ionization of Acids and Bases

● Law of Chemical Equilibrium and Equilibrium Constant

● Solubility Equilibria

Learn more about pH Scale here in more detail.

Types of Buffer Solutions

Buffers are broadly divided into two types – acidic and alkaline buffer

solutions. Acidic buffers are solutions that have a pH below 7 and

contain a weak acid and one of its salts. For example, a mixture of

acetic acid and sodium acetate acts as a buffer solution with a pH of

about 4.75.

Alkaline buffers, on the other hand, have a pH above 7 and contain a

weak base and one of its salts. For example, a mixture of ammonium

chloride and ammonium hydroxide acts as a buffer solution with a pH

of about 9.25. Buffer solutions help maintain the pH of many different

things as shown in the image below.

Preparation of a Buffer Solution

If you know the pKa (acid dissociation constant) of the acid and pKb

(base dissociation constant) of the base, then you can make a buffer of

known pH by controlling the ratio of salt and acid or salt and base.

Buffers can either be prepared by mixing a weak acid with its

conjugate base or a weak base with its conjugate acid.

For example, phosphate buffer, a commonly used buffer in research

labs, consists of a weak base (HPO42-) and its conjugate acid

(H2PO4–). Its pH is usually maintained at 7.4.

Understand the Concept of Equilibrium in Chemical Process in detail

here.

Buffer Action

So, how does a buffer work? Let’s take the example of a mixture of

acetic acid (CH3COOH) and sodium acetate (CH3COONa). Here,

acetic acid is weakly ionized while sodium acetate is almost

completely ionized. The equations are given as follows:

CH3COOH H+ + CH3COO–

CH3COONa Na+ + CH3COC–

To this, if you add a drop of a strong acid like HCl, the H+ ions from

HCl combine with CH3COO– to give feebly ionized CH3COOH.

Thus, there is a very slight change in the pH value. Now, if you add a

drop of NaOH, the OH– ions react with the free acid to give

undissociated water molecules.

CH3COOH + OH– CH3COO– + H2O

In this way, the OH– ions of NaOH are removed and the pH is almost

unaltered.

Learn more about ionization of Acids and Bases here.

Solved Example for You

Question: Which of the following statement/s is false about buffer

solutions?

a. The pH of a buffer solution does not change on dilution.

b. Buffer solutions do not have a definite pH.

c. The pH of a buffer solution changes slightly on the addition of

a small amount of acid or base.

d. The pH of buffer solution does not change on standing for long.

Solution: The option ‘b’ is false. Buffer solutions have a definite pH.

Equilibrium in Physical Processes

Just the way chemical reactions attain a state of equilibrium, there

exists equilibrium in physical processes too. This refers to the

equilibrium that develops between different states or phases of a

substance such as solid, liquid and gas. Let’s try and understand

equilibrium in physical processes in more detail. Substances undergo

different phase transformation processes such as –

solid ⇌ liquid

liquid ⇌ gas

solid ⇌ gas

Let’s understand how they attain equilibrium during each of these

transformations.

Solid-Liquid Equilibrium

What happens if you keep ice and water in a perfectly insulated

manner, such as in a thermos flask at a temperature of 273K and

atmospheric pressure? We see that the mass of ice and water do not

change and that the temperature remains constant, indicating a state of

equilibrium.

However, the equilibrium is not static because there is intense activity

at the boundary between ice and water. Some ice molecules escape

into liquid water and some molecules of water collide with ice and

adhere to it. Despite this exchange, there is no change in mass of ice

and water. This is because the rates of transfer of ice molecules to

water and the reverse process are equal at 273K and atmospheric

pressure.

It is evident that ice and water are in equilibrium only at a particular

pressure and temperature. Therefore, for any pure substance at

atmospheric pressure, the temperature at which the solid and liquid

phases are at equilibrium is called the normal melting point or normal

freezing point of the substance.

The system of ice and water is in dynamic equilibrium and we can

conclude the following –

● Both opposing processes occur at the same time.

● The two processes occur at the same rate such that the amount

of ice and water remain constant.

Browse more Topics under Equilibrium

● Acids, Bases and Salts

● Buffer Solutions

● Equilibrium in Chemical Processes

● Factors Affecting Equilibria

● Ionization of Acids and Bases

● Law of Chemical Equilibrium and Equilibrium Constant

● Solubility Equilibria

Liquid-vapour Equilibrium

Equilibrium Vapor Pressure

To understand this concept, let’s perform the following experiment.

Experiment:

● Place a drying agent like anhydrous calcium chloride for a few

hours in a transparent box with a U-tube containing mercury

i.e. manometer. This will soak up all the moisture in the box.

● Remove the drying agent by tilting the box to one side and

quickly place a petri dish containing water.

Observations:

● The mercury in the manometer rises slowly and then attains a

constant value. This is because the pressure inside the

manometer increases due to the addition of water molecules

into the gaseous phase.

● Initially, there is no water vapour in the box. As the water in

the petri dish evaporates, the volume of water in the petri dish

decreases and the pressure in the box increases.

● The rate of increase in pressure decreases with time because of

condensation of vapour into water.

● Finally, it reaches an equilibrium where there is no net

evaporation or condensation.

Conclusion: Equilibrium is reached when the

rate of evaporation = rate of condensation

H2O(l) ⇌ H2O (vap)

At equilibrium, the pressure that the water molecules exert at a given

temperature remains constant and is called the equilibrium vapour

pressure of water. The vapour pressure of water increases with

temperature.

Boiling Point

At the same temperature, different liquids have different equilibrium

vapour pressures. The liquid with a higher vapour pressure is more

volatile and has a lower boiling point. Let’s understand this concept

with the following experiment.

Experiment:

● Expose three Petri dishes containing 1ml each of acetone, water

and ethyl alcohol to the atmosphere.

● Repeat the experiment with different liquid volumes in a

warmer room.

Observations:

● In all cases, the liquid eventually disappears.

● The time taken for complete evaporation of each liquid differs.

Conclusions:

● The time taken for complete evaporation of the liquid depends

on – the nature of the liquid, the amount of liquid and the

temperature.

● In an open system i.e. when the petri dish is kept open, the rate

of evaporation remains constant but the molecules of the liquid

are dispersed into a larger volume of the room. Consequently,

the rate of evaporation is much higher than the rate of

condensation from vapour to liquid. Therefore, open systems

do not reach an equilibrium.

On the other hand, in a closed vessel or system, water and water

vapour are in equilibrium at atmospheric pressure (1.013 bar) and

100°C.

For any pure liquid at one atmospheric pressure, the temperature at

which the liquid and vapour are at equilibrium is called the normal

boiling point of the liquid.

For water, the boiling point is 100°C at atmospheric pressure. The

boiling point of liquids depends on atmospheric pressure i.e. the

altitude of the place. Boiling point decreases at higher altitudes.

Solid-Vapor Equilibrium

Have you ever observed what happens to solid iodine placed in a

closed jar? The jar gets filled up with violet coloured vapour and the

colour intensity increases with time! The intensity of the colour

becomes constant after a certain time i.e. equilibrium is reached. In

this way, solid iodine sublimes to give iodine vapour while the vapour

condenses to form solid iodine. The equilibrium in this process is

given as –

I2(solid) ⇌ I2(vapour)

Sublimation of Iodine [Source: Wikimedia Commons]

Other substances that show this kind of equilibrium are –

Camphor (solid) ⇌ Camphor (vapour)

NH4Cl (solid) ⇌ NH4Cl (vapour)

Equilibrium Involving Dissolution of Solid or Gases in Liquids

Solids In Liquids

What happens when you make a thick sugar solution by dissolving

sugar at high temperature, then allow it to cool at room temperature.

Yes, the sugar crystals separate out.

In this case, the thick sugar solution is a saturated solution because no

more solute i.e. sugar can be dissolved in it at a given temperature.

The concentration of solute in a saturated solution depends on the

temperature. A dynamic equilibrium exists between the solute

molecules in the solid state and in solution in a saturated solution.

Sugar (solution) ⇌ Sugar (solid)

Also, the rate of dissolution of sugar = rate of crystallization of sugar.

Let’s understand this further, using an example. What happens when

you add radioactive sugar to a saturated solution of non-radioactive

sugar? You will see radioactivity both in the solution and solid sugar

after some time. Initially, there are no radioactive sugar molecules in

the solution.

But, due to the dynamic nature of equilibrium, there is an exchange

between the radioactive and non-radioactive sugar molecules from the

two phases. Thus, the ratio of radioactive to non-radioactive sugar

molecules in the solution increases till it reaches a constant value.

Gases In Liquids

Why do we see fizz and hear a sound when we open soda bottles?

This happens because some of the CO2 dissolved in it fizzes out

rapidly due to the difference in solubility of CO2 at different

pressures. The equilibrium between the CO2 molecules in the gaseous

state and those dissolved in liquid under pressure is given as –

CO2(gas) ⇌ CO2(in solution)

This equilibrium is governed by Henry’s law. It states that the mass of

a gas dissolved in a given mass of a solvent at any temperature is

proportional to the pressure of the gas above the solvent. This amount

decreases with increase in temperature.

The soda bottle is sealed under the pressure of the gas where its

solubility in water is high. When the bottle is opened, some of the CO2

escapes trying to reach a new equilibrium or its partial pressure in the

atmosphere. This is why soda water turns flat when the bottle is left

open for too long.

Fizz in soda water [Source: pxhere]

Features of Equilibrium in Physical Processes

Process Conclusion

Solid ⇌ Liquid

H2O(s) ⇌ H2O(l) Melting point is fixed at constant pressure.

Liquid⇌ Vapour

H2O(l)⇌ H2O(g) pH2O constant at given temperature.

Solute(s) ⇌ Solute (solution)

Sugar(s) ⇌ Sugar (solution)

Concentration of solute in solution is constant at a given temperature.

Gas(g) ⇌ Gas (aq)

CO2(g) ⇌ CO2(aq)

[gas(aq)]/[gas(g)] is constant at a given temperature.

[CO2(aq)]/[CO2(g)] is constant at a given

temperature.

General Characteristics Of Equilibrium In Physical Processes

The following characteristics are common to the state of equilibrium

in physical processes.

● At a given temperature, equilibrium in physical processes is

achieved only in a closed system.

● The opposing processes occur at the same rate and there exists

a dynamic but stable condition during equilibrium in physical

processes.

● All properties of the system that are measurable remain

constant.

● Equilibrium in physical processes is characterized by a constant

value of one of its parameters at a given temperature.

● The extent to which a physical process has progressed before

reaching equilibrium is indicated by the magnitude of the

abovementioned parameter at any stage.

Solved Example For You

Question: The fizz observed when you open a bottle of soda water is

governed by which of the following laws?

a. Murphy’s law

b. Henry’s law

c. Raoult’s law

d. Avogadro law

Solution: The answer is option ‘b’. Henry’s law explains the

phenomenon of fizz i.e. the escape of some CO2 molecules due to the

difference in solubility of CO2 at different pressures.

Factors Affecting Equilibria

Chemical equilibria in a chemical reaction define the state in which

there is no further change in the concentration of the reactants and

products. In numerous biological and environmental processes, these

chemical equilibria play an important role.

One example we can consider for chemical equilibria is that it

involves molecules of O2 and the protein haemoglobin play a crucial

role in the transportation and delivery of O2 from our lungs to our

muscles. Similarly, the equilibria that involve the CO molecules and

haemoglobin that accounts and leads to the toxicity of CO.

Factors Affecting Equilibria

According to the Le Chatelier’s Principle, states that if a system under

equilibrium is subjected to a change in pressure, temperature or

concentration, in this case, the equilibrium shifts further reducing as

well as to counteract the effect of the change. The factors that affect

equilibria are:

Effect of Pressure Change

There is no effect of pressure if the number of moles of gaseous

reactant and products is equalised. However, it is different for the total

number of moles of gaseous reactants and a total number of moles of

gaseous products. On increasing the pressure, the total number of

moles per unit volume also increases leading to shifting in the

equilibrium direction wherein the number of moles per unit volume

will be less.

If the total number of moles of products are more than the total

number of moles of reactants, in this case, the low pressure will also

favour forward reaction. If the number of moles of reactants is more

than that of products, high pressure would be favourable to forward

reaction.

Effect of Change of Concentration

The equilibrium changes when the concentration of any reactants or

products in the reaction changes. It further leads to minimising its

effect.

Effect of Inert Gas Addition

After the addition of an inert gas and with the volume kept constant,

there is no effect on the equilibrium. This is because, at constant

volume, the addition of an inert gas does not change partial pressure or

molar concentration.

Effect of Temperature Change

The equilibrium shifts in opposite direction when there is a change of

increase or decrease in the system of temperature. This takes effect in

order to neutralise the change in effect. In an exothermic reaction, the

low temperature favours forwards reaction –

For example. N2(g) + 3H2(G) ——— 2NH3 (g)

∆H = -92.38 kJ mol -1

Low temperature slows down the reaction for which we are using the

catalyst. In case of an endothermic reaction, the increase in

temperature will shift the equilibrium in direction of the endothermic

reaction.

Effect of a Catalyst

There is no effect on the equilibrium composition of a reaction

mixture. This is because catalyst increases the speed of both forward

and backward reactions to the same extent in a reversible reaction.

Solved Examples for You

Question: In the reaction A + B -> C + D, what will happen to the

equilibrium if the concentration of A is increased?

Solution: The reaction will increase in the forward direction.

Ionization of Acids and Bases

The process in which neutral molecules get splits up into charged ions

when exposed in a solution is referred to as the ionization of a

compound. According to the Arrhenius theory, the acids are the

compounds that dissociate in the aqueous medium in order to generate

the hydrogen ions, H+ in the aqueous medium. Find interesting? Let’s

learn more about the ionization of acids and bases in this section.

Ionization of a Compound

While the bases are those compounds that furnish the hydroxyl ions,

OH– in the aqueous medium. The degree of ionization of the acids and

bases helps determine its strength. On the basis of different acidic and

basic compounds, the degree of ionization may differ.

Browse more Topics Under Equilibrium

● Acids, Bases and Salts

● Buffer Solutions

● Equilibrium in Chemical Processes

● Equilibrium in Physical Processes

● Factors Affecting Equilibria

● Ionization of Acids and Bases

● Law of Chemical Equilibrium and Equilibrium Constant

● Solubility Equilibria

Ionization of Acids

The degree of Ionisation refers to the strength of an acid or a base. A

strong acid is said to completely ionize in water whereas a weak acid

is said to only ionise partially. As there are different degrees of

ionization of acids, there are also different levels of weakness for

which there is a simple quantitative way to express.

Since the ionization of a weak acid is an equilibrium, the chemical

equation and an equilibrium constant expression can be stated as :

HA ( aq ) + H2O ( l ) H3O+ ( aq ) + A–

Ka = [ H3O+ ] [A–] / [HA]

Equilibrium Constant for ionisation of an acid defines its Acid

Ionisation Constant (Ka). However, the stronger the acid, the larger

will be the acid ionisation constant (Ka). This means that a strong acid

is a better proton donor. As a result of the concentration of the product

in the numerator of the Ka, the stronger the acid the larger is the acid

ionisation constant (Ka).

Ionization of Bases

Some bases like lithium hydroxide or sodium hydroxide get

completely dissociated into their ions in an aqueous medium which is

referred to as strong bases. Therefore, the ionisation of these bases

yields hydrochloric ions, as (OH–). A similar expression for the bases

is:

A + H2O OH– + HA+

Kb = [ OH– ] [ HA+ ] / [ A ]

The base ionisation constant i.e Kb refers to the equilibrium constant

for the ionisation of a base. Therefore, we can say that a strong base

implies a good proton acceptor while a strong acid implies a good

proton donor. The dissociation of weak acids or weak bases in water

is:

CH3COOH + H2O ⇔ CH3COO‾ + H3O+

NH3 + H2O ⇔ NH4+ ( aq ) + OH‾( aq )

Solved Examples for You

Question: A 0.500 M solution of formic acid has a pH value of 2.04.

Determine the Ka for formic acid.

Solution. Initial [HCOOH] = 0.500 M and pH = 2.04. Unknown, Ka =

?

Concentrations [HCOOH] [H + ] [HCOO − ]

Initial 0.500 0 0

Change -9.12 × 10 -3 +9.12 × 10 -3 +9.12 × 10 -3

Equilibrium 0.491 9.12 × 10 -3 9.12 × 10 -3

Now substituting into the Ka expression gives:

The value of Ka is consistent with that of a weak acid. Now, following

the same steps, we find the value of Kb of the base. If 0.750 M

solution of the weak base ethylamine (C2H5 NH2) has a pH of 12.31.

The pOH is 14 – 12.31 = 1.69. The [OH − ] is found from 10 -1.69 =

2.04 × 10 -2 M. The ICE table shall be then as shown below.

Concentrations [C2H5NH2] [C2H5NH3+] [OH−]

Initial 0.750 0 0

Change -2.04 × 10 -2 +2.04 × 10 -2 +2.04 × 10 -2

Equilibrium 0.730 2.04 × 10 -2 2.04 × 10 -2

Substituting the value of Kb it yields the Kb for Ethylamine.

Law of Chemical Equilibrium and Equilibrium Constant

Chemical equilibrium refers to the state wherein both the

reactants and the products present in the concentration have no

tendency to change with the period of time during a chemical

reaction. A chemical reaction achieves chemical equilibrium

when the rate of forward reaction and that of the reverse

reaction is same. Also, since the rates are equal and there is no

net change in the concentrations of the reactants and the

products – the state is referred to as a dynamic equilibrium and

the rate constant is known as equilibrium constant. Let’s find

out more.

Law of Chemical Equilibrium

Representation of the attainment of chemical equilibrium is –

The equilibrium constant is defined as the product of the molar

concentration of the products which is each raised to the power

equal to its stoichiometric coefficients divided by the product of

the molar concentrations of the reactants, each raised to the

power equal to its stoichiometric coefficients is constant at

constant temperature. T

his equilibrium constant can be simply expressed in terms of the

partial pressures of the reactants and the products. However, if

it is expressed in terms of the partial pressure, it is denoted by

Kp.

Browse more Topics under Equilibrium

● Acids, Bases and Salts

● Buffer Solutions

● Equilibrium in Chemical Processes

● Equilibrium in Physical Processes

● Factors Affecting Equilibria

● Ionization of Acids and Bases

● Solubility Equilibria

Equilibrium Constant

Equilibrium Constant Units and Formula

Law of mass action also forms the basis which states that the

rate of a chemical reaction is directly proportional to the

product of the concentrations of the reactants raised to their

respective stoichiometric coefficients. Therefore, given the

reaction –

aA(g) + bB(g) ⇔ cC(g) + dD(g)

By using the law of mass action here,

● The forward reaction rate would be k+ [A]a[B]b

● The backward reaction rate would be k– [C]c[D]d

where, [A], [B], [C] and [D] being the active masses and k+ and

k− are rate constants of forward and backward reactions, also

the a, b, c, d are the stoichiometric coefficients related to A, B, C

and D respectively. However, at the equilibrium – the forward

and the backward rates are equal, stating –

Rate of forward reaction = Rate of backward reaction

or,

or,

where,

Kc is the equilibrium constant expressed in terms of the molar

concentrations. The equation Kc = [ C ]c·[ D ]d / [ A ]a·[ B ]b

or, Kc = Kf / Kb is the Law of Chemical Equilibrium. The

equilibrium constant is therefore related to the standard Gibbs free

energy change for the reaction which is stated by the equation –

§Gº= -RT ln Keq

where T states the temperature, R is the universal gas constant and Keq

is the equilibrium constant.

Solved Examples for You

Question: Write the equilibrium constant expression for the

reaction equation:

NH3 + HOAc = NH4+ + OAc–

Solution:

[NH4+] [OAc–] ——————— = K (unitless constant) [NH3]

[HOAc]

Question: A closed container has N2O4 and NO2 gases in it. It

has been placed in the lab for many days. What would you

consider the container and the gases to be?

a. an open system

b. a closed system

c. not a system

Solution: A closed system since it has been in the lab, the

temperature of the system is the same as its environment.

Solubility Equilibria

The equilibrium constant that represents the maximum amount of

solid which can be dissolved in an aqueous solution is defined as the

solubility product. The Solubility Equilibria are based on the

assumption that the solids dissolve in water in order to give the basic

particles from which they are formed.

Solubility Equilibria

Each of the molecular solids dissolves to give an individual aqueous

molecule such as

H2O

C12H22O11(S) —————> C12H22O11 (aq)

The ionic solids dissociate to give their respective positive and

negative ions :

H2O

NaCl (s) ————————-> Na + (aq) + Cl – (aq )

The ions formed from the dissociation of the ionic solids can also

carry an electrical current which makes the salt solutions good

conductors of electricity. However, molecular solids do not dissociate

in water to give ions so as no electrical current can be carried.

Browse more Topics under Equilibrium

● cids, Bases and Salts

● Buffer Solutions

● Equilibrium in Chemical Processes

● Equilibrium in Physical Processes

● Factors Affecting Equilibria

● Ionization of Acids and Bases

● Law of Chemical Equilibrium and Equilibrium Constant

Solubility

The ratio of the maximum amount of solute to the volume of the

solvent in which the solute can dissolve. This is generally expressed in

two ways i.e –

a. Grams of solute per 100 g of water

b. Moles of solute per litre of solution

A salt refers to be soluble if it dissolves in water to give a solution

along with the concentration of at least 0.1 m at the room temperature.

A salt is also considered to be insoluble if the concentration of an

aqueous solution is less than 0.0001 m at the room temperature. Salts

are considered to be slightly soluble; those between 0.0001 m and 0.1

m.

According to the principles of solubility equilibria, the salts have low

solubilities in water. The reaction that can be considered for the

dissociation of the salt AgCl is –

AgCl(s) ————— Ag+(aq) + Cl– (aq)

However, the reverse reaction for the dissolving of the salt would be

the precipitation of the ions to form a solid –

Ag+(aq) + Cl– (aq) —————– AgCl(s)

When the rate at which AgCl dissolves is equal to the rate at which

AgCl precipitates, the system has reached the solubility equilibria.

Solubility Product Equilibrium Constant (Ksp)

The product of the equilibrium concentrations of the ions in a

saturated solution of a salt. Each concentration in this is raised to the

power of the respective coefficient of ion in the balanced equation.

For example, the solubility product equilibrium constant for the

dissociation of AgCl is:

Another example of a solubility product equilibrium constant where

we can consider the reaction for the dissociation of CaF2 in water is :

The solubility product equilibrium constant for this reaction would be

the product of the concentration of Ca2+ ion and the concentration of

the F– ion raised to the second power (squared):

Solved Examples for You

Question: In the atmosphere, SO2 and NO are oxidised to SO3 and

NO2, respectively, which react with water to give H2SO4 and HNO3.

The resultant solution is called acid rain. SO2 dissolves in water to

form diprotic acid.

SO2(g)+H2O(l)⇔HSO⊝3+H⊕; Ka1=10−2

HSO⊝3⇔SO2−3+H⊕; Ka2=10−7

and for equilibrium,

SO2(aq)+H2O(l)⇔SO2−3(aq)+2H⊕(aq)

Ka=Ka1×Ka2=10-9 at 300K.

A. H2SO3 is less acidic than H2SO4

B. HNO3 is less acidic than HNO2

C. SO2(g) is reduced in the atmosphere during a thunderstorm

D. CO2 gas develop more acidity in water than SO2

Solution: Option A. Sulphurous acid H2SO3 is less acidic than

sulphuric acid H2SO4. When a proton is lost by sulphurous acid, the

negative charge is delocalized on 3 O atoms and when a proton is lost

by sulphuric acid, the negative charge is delocalized by 4 O atoms.

Since the extent of delocalization in HSO-1 ion is less than the extent

of delocalization in HSO-4, the stability of anion and the acidity of the

acid in sulphurous acid is lower than in sulphuric acid.