Ionisation of Water

We have learnt that when an acidic or a basic substance is dissolved in water, depending
upon its nature, it can either donate (or) accept a proton. In addition to that the pure water
itself has a little tendency to dissociate. i.e, one water molecule donates a proton to an
another water molecule. This is known as auto ionisation of water and it is represented as
below.

In the above ionisation, one water molecule acts as an acid while the another water
molecule acts as a base. The dissociation constant for the above ionisation is given by the
following expression

∵ Kw = [H3O+][OH–] …………. (8.4)

The concentration of pure liquid water is one. i.e, [H 2O]2 = 1

Here, Kw represents the ionic product (ionic product constant) of water.

It was experimentally found that the concentration of H 3O+ in pure water is 1 × 10-7 at 25°C.
Since the dissociation of water produces equal number of H 3O+ and OH–, the concentration of
OH– is also equal to 1 × 10-7 at 25°C.

Therefore, the ionic product of water at 25°C is

KW = [H3O]+[OH–] …………. (8.4)

KW = (1 × 10-7)(1 × 10-7)
= 1 × 10-14.

Like all equilibrium constants, Kw is also a constant at a particular temperature. The

dissociation of water is an endothermic reaction. With the increase in temperature, the
concentration of H3O+ and OH– also increases, and hence the ionic product also increases.

In neutral aqueous solution like NaCl solution, the concentration of H 3O+ is always equal to
the concentration of OH– whereas in case of an aqueous solution of a substance which may
behave as an acid (or) a base, the concentration of H 3O+ will not equal to
[OH–].
We can understand this by considering the aqueous HCl as an example. In addition to the
auto ionisation of water, the following equilibrium due to the dissociation of HCl can also
exist.

HCl + H2O ⇄ H3O+ + Cl–

In this case, in addition to the auto ionisation of water, HCl molecules also produces H 3O+ ion
by donating
a proton to water and hence [H3O+]>[OH–]. It means that the aqueous HCl solution is acidic.
Similarly, in basic solution such as aqueous NH 3, NaOH etc…. [OH–]>[H3O+].

Special Properties of Water

1. Water is polar :-
A single water molecule is held together by polar covalent bonds. Oxygen is more
electronegative than hydrogen, so the electrons in the H–O bond spend more time closer to the
oxygen atom. That side of the water molecule has a partial negative charge (shown with a
lowercase Greek letter delta minus: δ- ).The other side (next to the hydrogens) then has a partial
positive charge (delta plus: δ+) because the electrons spend less of their time there. The polar
nature of water influences all of its interactions with biochemical molecules. Nonpolar
(hydrophobic) molecules will act to exclude water, which is another kind of interaction in
itself.Water molecules form hydrogen bonds with each other
2. Pay special attention to this property–it explains virtually all of the unusual behaviors of water.
Water molecules are very likely to form hydrogen bonds between different water molecules
(not within a single water molecule, that’s a polar covalent bond).H bonds are weaker than
covalent bonds. Picture them blinking on and off–that’s why they’re shown as a dotted line
 rather than a solid line. This means water, as a material, is fairly interconnected to itself, which
leads to several interesting emerging properties
3. Water is cohesive
Because water is extensively hydrogen bonded, it behaves differently than materials where
their molecules are not stuck together. For instance, you can see a drop of water beading up
into a round shape (compared to, say, a drop of alcohol that lies flat). This is important when
you consider how body fluids move in plants and animals
4. Water has high surface tension
Because water is extensively hydrogen bonded (and is cohesive), it forms almost a skin on its
surface. High surface tension means water can support a surprising amount of weight before
something sinks. This provides a new place to support some forms of life–like a stick insect or
Jesus lizard, running across the surface of a body of water
5. Water adheres to surfaces
Because of its ability to form hydrogen bonds, water can stick to certain surfaces. This is
especially important for moving water (and watery liquids) through living tissues, like how water
moves through a plant through the process of transpiration.3 3 The combination of cohesion
and adhesion move water molecules in an unbroken column from the roots of a plant all the
way up through the leaves.
6. Water has high specific heat
Because of water’s extensive hydrogen bonding, it’s hard to raise its temperature (compared to
other liquids). You have to add enough heat energy to break the hydrogen bonds between
water molecules before they are free to start vibrating faster.4 Because cells are mostly 4 made
of water, this makes easier for living creatures to maintain a constant body temperature
7. Water has high heat of vaporization
Because of water’s extensive hydrogen bonding, it takes a lot of energy for water to evaporate.
Before it can change from liquid to gas, it has to break all those hydrogen bonds holding the
water molecules together. Then the individual water molecules can start vibrating with higher
kinetic energy, eventually breaking free and leaving as a gas (water vapor). Living organisms can
take advantage of this in the form of "evaporative cooling"–when water evaporates, it carries
away heat energy, leaving a cooler animal behind (think of how a dog pants to cool off).
8. Solid water is less dense than liquid water
When water changes from liquid to solid (ice), the molecules are held frozen in a fixed structure.
Ice has a lot of open space between water molecules, compared to how they are closer together
in liquid water. This is to say ice is less dense than liquid water , which means ice floats on top of
liquid water. Again, this provides additional habitats for living organisms (like polar bears). It also
insulates the liquid water underneath, keeping fish and other aquatic creatures alive in cold
weather

What is pH?
 pH is a measurement of the degree to which water is “acidic” (like lemon juice) or “basic” (like
bleach or soap). pH is measured on a scale that ranges from 0 (strongly acidic) to 14 (strongly
basic). In the middle is 7, where the pH is “neutral” (like in pure water).
In the term pH, the H refers to the element Hydrogen, specifically Hydrogen H+ . The plus means
that this Hydrogen is carrying a positive charge. Charged chemicals are also commonly called
“ions”, so H+ is a Hydrogen ion. The “p” in pH stands for “power” of the activity of the H+ ion
activity in the water.
The pH scale is based on the balance between H+ ions and negatively charged Hydroxide ions
(OH- ) in the water. When pH is low (e.g., pH = 2) there are a lot of H+ ions in the water. When
pH is high (e.g., pH = 13) there are a lot of OHions in the water. At the middle (pH=7), the two
ions balance each other out completely. Each whole pH value below 7 is ten times more acidic
than the next higher value. The same is true for whole pH values above 7, each is ten times
more basic than the value below it.
Most lakes are basic (also referred to as “alkaline”) when they first form. They become more
acidic as they age due to buildup of organic material from plant leaves and animals. As organic
substances decompose in the lake, carbon dioxide gas (CO2) is formed. This CO2 combines with
water to form a weak acid, called "carbonic" acid. This carbonic acid lowers the water’s pH over
time.
ACIDS AND BASES:-
Definitions
1. Arrhenius Definition – The first definition of acids and bases. Arrhenius defined an acid as a
species that produces H+ (a proton) in an aqueous solution and a base as a species that
produces OH- (a hydroxide ion) in a aqueous solution.
2. Bronsted-Lowry Definition – A Bronsted-Lowry acid donates protons, while a Bronsted-Lowry
base accepts protons. However, they can not be called Arrhenius bases since in aqueous
solution they do not dissociate to form OH-. The advantage of this definition is that it is not
limited to aqueous solutions. Bronsted-Lowry acids and bases always occur in pairs called
conjugate acidbase pairs. The pairs are related through the transfer of a proton.
3. Lewis Definition – Lewis defined an acid as an electron-pair acceptor and a base as a electron
pair donor. This definition is the most inclusive and encompasses species not included in the
Bronsted-Lowry definition.
Hydrogen Ion Equilibria Hydrogen ion concentration, [H+ ], is a measure of both Arrhenius and
Bronsted-Lowry defined acids. This is generally expressed as pH, where:
pH = -log[H+ ] = log(1/[H+ ])
Likewise, hydroxide ion concentration, [OH- ], is measured as pOH
. pOH = -log[OH- ] = log(1/[OH- ])
In any aqueous solution, the H2O solvent dissociates slightly:
H2O <===> H+ + OH
This dissociation is an equilibrium reaction and is therefore described by a constant, Kw, the
water dissociation constant. Kw = [H+ ][OH- ] = 10-14 Rewriting this equation using logarithms
show: pH + pOH = 14 In pure H20, [H+ ] is equal to [OH- ], since for everyone H20 that
dissociates, one H+ and OH- is made. A solution with equal H+ and OH- is neutral and has a pH of
 7. A pH below 7 indicates an excess of H+ and thus an acidic solution. A pH above 7 indicates a
relative excess of OH- ions and therefore a basic solution.
Strong Acids and Bases
Strong acids and bases are those that completely dissociate into their component ions in
aqueous solution. For example, NaOH is added to water, it dissociates completely:
NaOH + water ---> Na+ + OH
In a 1 M solution of NaOH, complete dissociation gives 1 M of OH- :
PH = 14 – (-log[OH- ]) = 14 + log[1] = 14
Weak Acids and Bases
Weak acids and bases are those that only partially dissociate in aqueous solution. A weak
monoprotic acid, HA in aqueous solution will achieve the follwing equilibrium after dissociation:
HA + H2O <===> H3O+ + A
The acid dissociation constant, Ka, is a measure of the how much the acid dissociates.
Ka = [H30+ ][A- ]/[HA]
Similarly, a weak monovalent base, BOH, dissociates to give B+ and OH- . The base dissociation
constant, Kb, is a measure of the degree that the base dissociates.
Kb = [B+ ][OH- ]/[BOH]
Thus Ka and Kb are inversely related. In other words, if Ka is large (the acid is trong), then Kb
(base is weak) will be small and vice versa. From this relationship, one can see that when a
conjugate acid/base pair form from a weak acid, the conjugate base is stronger than the acid.
Titration
Titration is a procedure used to determine the concentration of an acid or base. This is
accomplished by reacting a known volume of a solution of an unknown concentration with a
known volume of a solution of known concentration. When the number of acid molecules equal
the number of base molecules added (or vice versa), the equivalence point is reached. The
equivalence point in a titration is estimated in two common ways.: either by using a graphical
method, plotting the pH of the solution as a function of added titrant by using a pH meter or by
watching for a color change of an added indicator. Indicators are weak organic acids or bases
that have different colors in their undissociated and dissociated states. These states are
dependent on pH.

Henderson-Hasselbalch Equation

It is used to calculate the pH of a solution. Knowing the pH of a solution is very important for many chemical
reactions as well as for biological systems. The Henderson-Hasselbalch equation gives the approximate pH
value of a buffer solution. It provides a relationship between the pH of acids (in aqueous solutions) and their
pKa (acid dissociation constant).

In 1908, Lawrence Henderson derived an equation to calculate the pH of a solution that resists changing its pH value
on dilution or with the addition of a small volume of acid or alkali.

This equation was later re-expressed in logarithmic terms by Karl Albert Hasselbalch in 1917. The resulting equation
was named as Henderson-Hasselbalch Equation and is written as:

pH = pKa + log10 ([A–]/[HA])

 Where, pKa represents the acid dissociation constant, [A –] represents the molar concentration of the conjugate base and
[HA] represents the molar concentration of the weak acid.
The equation can also be written as-

We can Understand the Henderson-Hasselbalch Equation with the help of a real-life situation. Suppose, Ravi is a physically fit guy
with no major health issues. What do you think will be the pH of a blood sample taken from Ravi’s body?

As Ravi is healthy, his pH should be near 7.4. Blood has buffers that resist small changes in pH. The maintenance of blood's pH is
of utmost necessary for cellular functions. Blood is buffered by plasma proteins, hemoglobin, phosphates and bicarbonate. A blood
pH should not go above 7.8 or fall below 6.8.

Henderson Hasselbalch Equation Formula

The Henderson Hasselbalch equation shows the approximate pH value of a Buffer solution. This equation represents the
relationship between the pH or pOH of an aqueous solution and the acid dissociation constant and the ratio of the concentrations of
the dissociated chemical species.

Henderson Hasselbalch Equation can only be used if the acid dissociation constant is known. Henderson Hasselbalch Equation
formula is written as follows:

pH=pKa+log10([A−]/[HA])

Where, pKa is the acid dissociation constant, [A−] is the molar concentration of the conjugate base, and [HA]is the molar
concentration of the weak acid.

Derivation of Henderson Hasselbalch Equation

The ionization constants of strong acids and strong bases can be calculated easily with the direct formula. But it is difficult to find the
ionization constants of the weak acids and bases with the same methods as the extent of ionization of these acids and bases are
very low. So, in such a situation, Henderson-Hasselbalch Equation is very helpful.

There are two ways to derive Henderson Hasselbalch Equation. The first derivation of the Henderson Hasselbalch Equation is for
Base and the second derivation is for acid.

There are 3 main Assumptions that we have to make before starting the derivation:

 −1 < log ([A−]/[HA]) < 1

 The self-ionization energy of water can be ignored.
 The acid used is monobasic.
 Derivation 1)For the derivation of Henderson Hasselbach Equation for acid, take an example of ionization of weak acid
HA:

In general, the equilibrium shown by a weak acid HA is:

HA + H2O ↔ H+ + A-

The dissociation constant of the reaction is

Ka = [H+] [A−] / [HA]

Taking the negative logarithm on both sides,

-log Ka = - log[H+] [A−] / [HA]

-log Ka = - log [ H+] - log [A−] / [HA]

We know that, -log [ H+]= pH and -log Ka = pKa, substituting these in Equation

We get the following equation,

pKa = pH -log [A−] / [HA]

After rearranging the equation, we get,

pH = pKa + log [A−] / [HA]

This is the Henderson Hasselbalch equation for acid.

Now, if [A−] = [HA]

We get log [A−] / [HA] = 0

Therefore, we get pH = pKa, which means that both the species are the same and the acid will be half dissociated.

Derivation 2) For the derivation of Henderson Hasselbalch Equation, take an example of ionization of a base:

B + H2O ↔ OH+ + HB-

The first step is to use the formula for acid dissociation constant, Ka

Kb = [BH+] [OH−] / [B]

Taking negative logs on both sides,

-log Kb = −log [BH+] [OH−] / [B]

-log Kb= -log [OH-] - log [BH+] / [B]

We know that - log [ OH-] = pOH and -log Kb= pKb

We get the following equation,

pKb= pOH -log [BH+] / [B]

After rearranging the equation, we get,

pOH = pKb+ log [BH+] / [B]

This is the Henderson Hasselbalch equation for Base.

Laws of Thermodynamics as Related to Biology

Definition: The laws of thermodynamics are important unifying principles of biology. These
principles govern the chemical processes (metabolism) in all biological organisms. The First Law
of Thermodynamics, also known as the law of conservation of energy, states that energy can
neither be created nor destroyed. It may change from one form to another, but the energy in a
closed system remains constant. The Second Law of Thermodynamics states that when energy is
transferred, there will be less energy available at the end of the transfer process than at the
beginning. Due to entropy, which is the measure of disorder in a closed system, all of the
available energy will not be useful to the organism. Entropy increases as energy is transferred. In
addition to the laws of thermodynamics, the cell theory, gene theory, evolution, and
homeostasis form the basic principles that are the foundation for the study of life.

First Law of Thermodynamics in Biological Systems

All biological organisms require energy to survive. In a closed system, such as the universe, this
energy is not consumed but transformed from one form to another. Cells, for example, perform
a number of important processes. These processes require energy. In photosynthesis, the
energy is supplied by the sun. Light energy is absorbed by cells in plant leaves and converted to
chemical energy. The chemical energy is stored in the form of glucose, which is used to form
complex carbohydrates necessary to build plant mass. The energy stored in glucose can also be
released through cellular respiration. This process allows plant and animal organisms to access
the energy stored in carbohydrates, lipids, and other macromolecules through the production of
ATP. This energy is needed to perform cell functions such as DNA replication, mitosis, meiosis,
cell movement, endocytosis, exocytosis, and apoptosis.

Second Law of Thermodynamics in Biological Systems

As with other biological processes, the transfer of energy is not 100% efficient. In
photosynthesis, for example, not all of the light energy is absorbed by the plant. Some energy is
reflected and some is lost as heat. The loss of energy to the surrounding environment results in
an increase of disorder or entropy. Unlike plants and other photosynthetic organisms, animals
cannot generate energy directly from the sunlight. They must consume plants or other animal
organisms for energy. The higher up an organism is on the food chain, the less available energy
it receives from its food sources. Much of this energy is lost during metabolic processes
performed by the producers and primary consumers that are eaten. Therefore, much less
energy is available for organisms in higher trophic levels. The lower the available energy, the less
number of organisms can be supported. This is why there are more producers than consumers
in an ecosystem. Living systems require constant energy input to maintain their highly ordered
state. Cells, for example, are highly ordered and have low entropy. In the process of maintaining
this order, some energy is lost to the surroundings or transformed. So while cells are ordered,
the processes performed to maintain that order result in an increase in entropy in the
cell's/organism's surroundings. The transfer of energy causes entropy in the universe to
increase.