About me

Dr. Abu Bin Imran

Fulbright fellow (USA), NEDO Fellow (Japan)
D.Eng: (Nagoya University, Nagoya, Japan)
M.Sc & B.Sc: SUST
Professor, Department of Chemistry, BUET

Office: OAB 258

Email: abimran@chem.buet.ac.bd

Telephone: PABX Ext 7969 (office)

Mobile: 01712762745
Web: abimran.buet.ac.bd
 Chem 113: Chemistry
3.00 credits, 3 hours/week

Dr. Abu Bin Imran

Properties of solutions, Colloid and Nanochemsitry,
Phase rule and phase diagram;
Energy and chemistry,
Introduction to computational chemistry;
Design of new molecules, materials and drug.

Dr. Ayesha Akter

Quantum concept in atomic structure, VSEPRT; molecular geometry,
Quantum concept in bonding; VBT and MOT, Frontier MOT and
electronic transition, Silicon chemistry, Chemistry of biodegradable
and conductive polymer; LED, LCD/touch screen, Chemistry of
proteins, nucleic acids (DNA, RNA), carbohydrates and lipids;
Electrochemistry; electrolytic conduction, corrosion, devices for
energy storage,
 References

1. Chemistry by Raymond Chang, Kenneth A.

Goldsby
2. General Chemistry
By Darrell D. Ebbing and Steven D. Gammon
3. Principles of Physical Chemistry by Kindle edition
by Maron, S. H., Prutton, C.F.
Solution
A solution is a homogeneous mixture of two or more substances.

focus involving at least one liquid component—that is, gas-liquid,

liquid-liquid, and solid-liquid solutions.
A saturated solution contains the maximum
amount of a solute that will dissolve in a given
solvent at a specific temperature.

An unsaturated solution contains less solute than

it has the capacity to dissolve.

A supersaturated solution, contains more solute

than is present in a saturated solution.
Supersaturated solutions are not very stable. In time, some
of the solute will come out of a supersaturated solution as
crystals.

Crystallization is the process in which dissolved solute

comes out of solution and forms crystals.
 A Molecular View of the Solution Process

The intermolecular attractions that hold molecules together

in liquids and solids also play a central role in the formation
of solutions. When one substance (the solute) dissolves in
another (the solvent), particles of the solute disperse
throughout the solvent. The solute particles occupy
positions that are normally taken by solvent molecules.

The ease with which a solute particle replaces a solvent

molecule depends on the relative strengths of three types of
interactions:
• solvent-solvent interaction
• solute-solute interaction
• solvent-solute interaction
 A Molecular View of the Solution Process

Steps 1 and 2 require energy input to break attractive

intermolecular forces; therefore, they are endothermic.
In step 3 the solvent and solute molecules mix. This
process can be exothermic or endothermic.
 A Molecular View of the Solution Process

If the solute-solvent attraction is stronger than the solvent-solvent

attraction and solute-solute attraction, the solution process is favorable,
or exothermic (ΔHsoln<0).

If the solute-solvent interaction is weaker than the solvent-solvent and

solute-solute interactions, then the solution process is endothermic
(ΔHsoln>0).

The solution process is governed by two factors. One is energy, which

determines whether a solution process is exothermic or endothermic.

The second factor is disorder.

when solute and solvent molecules mix to form a solution, there is an
increase in randomness, or disorder. In the pure state, the solvent and
solute possess a fair degree of order. Much of this order is destroyed when
the solute dissolves in the solvent.
 Solubility

Solubility is a measure of how much solute will dissolve in a

solvent at a specific temperature.

The saying “like dissolves like” is helpful in predicting the

solubility of a substance in a given solvent.

Carbon tetrachloride (CCl4) and benzene (C6H6) are nonpolar

liquids. The only intermolecular forces present in these
substances are dispersion forces.
When these two liquids are mixed, they readily dissolve in each
other, because the attraction between CCl4 and C6H6 molecules
is comparable in magnitude to the forces between CCl4
molecules and between C6H6 molecules. Two liquids are said to
be miscible if they are completely soluble in each other in all
proportions.
 Solubility

Alcohols such as methanol, ethanol, and 1,2-ethylene glycol are

miscible with water because they can form hydrogen bonds with
water molecules:

When sodium chloride dissolves in water, the ions are stabilized in

solution by hydration, which involves ion-dipole interaction. In general,
we predict that ionic compounds should be much more soluble in polar
solvents, such as water, liquid ammonia, and liquid hydrogen fluoride,
than in nonpolar solvents, such as benzene and carbon tetrachloride.
Because the molecules of nonpolar solvents lack a dipole moment, they
cannot effectively solvate the Na+ and Cl- ions.

Solvation is the process in which an ion or a molecule is surrounded by

solvent molecules arranged in a specific manner.

The process is called hydration when the solvent is water.

Predict the relative solubilities in the following cases: (a) Bromine (Br2)
in benzene (C6H6, µ=0 D) and in water (µ=1.87 D), (b) KCl in carbon
tetrachloride (CCl4, µ= 0 D) and in liquid ammonia (NH3, µ= 1.46 D), (c)
formaldehyde (CH2O) in carbon disulfide (CS2, µ= 0 D) and in water.
 Concentration Units
Percent by Mass
The percent by mass (also called percent by weight or
weight percent) is the ratio of the mass of a solute to the
mass of the solution, multiplied by 100 percent:

The percent by mass is a unitless number because it is a

ratio of two similar quantities.

independent of temperature
 Concentration Units

Mole Fraction (X)

The mole fraction of a component of a solution, say,
component A, is written XA and is defined as

The mole fraction is also unitless

 Concentration Units

Molarity (M)
molarity is the number of moles of solute in 1 L of
solution;
the units of molarity are mol/L.
 Concentration Units
Molality (m)
Molality is the number of moles of solute dissolved in 1
kg (1000 g) of solvent

molality is independent of temperature, because the

concentration is expressed in number of moles of solute and
mass of solvent. The volume of a solution typically increases
with increasing temperature, so that a solution that is 1.0 M
at 25°C may become 0.97 M at 45°C because of the increase in
volume on warming. This concentration dependence on
temperature can significantly affect the accuracy of an
experiment. Therefore, it is sometimes preferable to use
molality instead of molarity.
Calculate the molality of a sulfuric acid solution containing 24.4 g
of sulfuric acid in 198 g of water. The molar mass of sulfuric acid is
98.09 g.
Calculate the concentration of a 0.396 m glucose (C6H12O6) solution in
molarity. Molar mass of glucose =180.2 g and the density of the solution
=1.16g/mL
Soln
The density of a 2.45 M aqueous solution of methanol (CH3OH) is 0.976
g/mL. What is the molality of the solution? The molar mass of methanol
is 32.04 g.
Calculate the molality of a 35.4 percent (by mass) aqueous
solution of phosphoric acid (H3PO4). The molar mass of
phosphoric acid is 97.99 g.
 The Effect of Temperature on Solubility

Solubility is defined as the maximum amount of a solute that will dissolve in a

given quantity of solvent at a specific temperature.

Solid Solubility and Temperature

the solubility of a solid substance

increases with temperature.

there is no clear correlation between

the sign of ∆Hsoln and the variation
of solubility with temperature.

the solution process of CaCl2 is

exothermic, and that of NH4NO3 is
endothermic
The Effect of Temperature on Solubility
Solid Solubility and Temperature

Fractional Crystallization
Fractional crystallization is the separation of a
mixture of substances into pure components on
the basis of their differing solubilities.

Many of the solid inorganic and organic

compounds that are used in the laboratory
were purified by fractional crystallization.
The Effect of Temperature on Solubility
Gas Solubility and Temperature

The solubility of gases in water usually

decreases with increasing temperature

When water is heated in a beaker, you can see

bubbles of air forming on the side of the glass
before the water boils. As the temperature
rises, the dissolved air molecules begin to
“boil out” of the solution long before the
water itself boils.
The Effect of Temperature on Solubility
Gas Solubility and Temperature

Thermal pollution
Every year in the United States some 100,000
billion gallons of water are used for
industrial cooling, mostly in electric power
and nuclear power production. This process
heats the water, which is then returned to the
rivers and lakes from which it was taken.

An increase in water temperature accelerates their

rate of metabolism, which generally doubles with
each 10°C rise.

The speedup of metabolism increases the

fish’s need for oxygen at the same time that the
supply of oxygen decreases because of its lower
solubility in heated water.
 The Effect of Temperature on Solubility
Gas Solubility and Temperature
Fishing on a hot summer day

an experienced fisherman usually picks a deep spot in the river or lake to cast the
bait. Because the oxygen content is greater in the deeper, cooler region, most fish
will be found there.
 The Effect of Pressure on the Solubility of Gases
External pressure has no influence on the solubilities of liquids and solids,
but it does greatly affect the solubility of gases. The quantitative relationship
between gas solubility and pressure is given by Henry’s† law, which states that
the solubility of a gas in a liquid is proportional to the pressure of the gas over
the solution:
William Henry (1775–1836). English chemist

c is the molar concentration (mol/L) of the dissolved gas;

P is the pressure (in atm) of the gas over the solution at
equilibrium;
k is a constant that depends only on temperature. The
constant k has the units mol/L.atm.
 The Effect of Pressure on the Solubility of Gases

A molecular interpretation of Henry’s law. When the

partial pressure of the gas over the solution increases
from (a) to (b), the concentration of the dissolved gas
also increases

The amount of gas that will dissolve in a solvent depends on how

frequently the gas molecules collide with the liquid surface and become
trapped by the condensed phase.
The solubility of nitrogen gas at 25°C and 1 atm is 6.8 × 10-4mol/L. What is the
concentration (in molarity) of nitrogen dissolved in water under atmospheric
conditions? The partial pressure of nitrogen gas in the atmosphere is 0.78 atm.

Solution

The decrease in solubility is the result of lowering the pressure from 1 atm to 0.78 atm.
 The Effect of Pressure on the Solubility of Gases

Exception of Henry’s law:

 The solubility of ammonia is much higher than expected because of the reaction

 Carbon dioxide also reacts with water, as follows:

 Oxygen gas is only sparingly soluble in water. However, its solubility in blood is
dramatically greater because of the high content of hemoglobin (Hb) molecules.
Each hemoglobin molecule can bind up to four oxygen molecules, which are
eventually delivered to the tissues for use in metabolism:
 The Effect of Pressure on the Solubility of Gases
 On August 21, 1986, Lake Nyos in Cameroon, a
small nation on the west coast of Africa,
suddenly belched a dense cloud of carbon
dioxide. Speeding down a river valley, the
cloud asphyxiated over 1700 people and many
livestock.

 Lake Nyos is stratified into layers that do not

mix. A boundary separates the freshwater at
the surface from the deeper, denser solution
containing dissolved minerals and gases,
including CO2. The CO2 gas comes from
springs of carbonated groundwater that
percolate upward into the bottom.

 earthquake, landslide, or even strong winds

may have upset the delicate balance within the
Deep waters in Lake Nyos are
lake, creating waves that overturned the water
pumped to the surface to
remove dissolved CO2 gas. layers.
Colligative Properties of Nonelectrolyte Solutions
Colligative properties (or collective properties) are properties that depend
only on
the number of solute particles in solution and not on the nature of the
solute particles.

regardless of whether they are atoms, ions, or molecules.

The colligative properties are

 vapor-pressure lowering,
 boiling-point elevation,
 freezing-point depression,
 and osmotic pressure.

Relatively dilute solutions whose concentrations are ≤ 0.2 M.

Colligative Properties of Nonelectrolyte Solutions

Vapor-Pressure Lowering
If a solute is nonvolatile (that is, it does not have a measurable vapor
pressure),

the vapor pressure of its solution is always less than that of the pure
solvent.

Thus, the relationship between solution vapor pressure and solvent vapor
pressure depends on the concentration of the solute in the solution.

This relationship is expressed by Raoult’s law

Colligative Properties of Nonelectrolyte Solutions
Vapor-Pressure Lowering
Raoult’s law: It states that the vapor pressure of a solvent over a
solution, P1, is given by the vapor pressure of the pure solvent, P°1,
times the mole fraction of the solvent in the solution, X1:

In a solution containing only one solute, X1 = 1 - X2, where X2 is the mole

fraction of the solute.

the decrease in vapor pressure, ∆P, is directly proportional to the solute

concentration (measured in mole fraction).
 Colligative Properties of Nonelectrolyte Solutions
Vapor-Pressure Lowering
Why is the vapor pressure of a solution less than that of the pure solvent?

One driving force in physical and chemical processes is an increase in disorder—

the greater the disorder, the more favorable the process. Vaporization increases the
disorder of a system because molecules in a vapor have less order than those in a
liquid. Because a solution is more disordered than a pure solvent, the difference
in disorder between a solution and a vapor is less than that between a pure solvent
and a vapor.

The presence of solute particles disrupts the intermolecular forces between

the solvent molecules, making it harder for them to escape into the vapor
phase.

the vapor pressure of the solution will be lower than that of the pure solvent
due to the reduction in the number of solvent molecules at the surface
available to evaporate.
 Colligative Properties of Nonelectrolyte Solutions
Vapor-Pressure Lowering
If both components of a solution are volatile (that is, have
measurable vapor pressure), the vapor pressure of the solution is
the sum of the individual partial pressures.

Raoult’s law holds equally well in this case:

where PA and PB are the partial pressures over the solution for
components A and B; PA 0 and PB 0 are the vapor pressures of the pure
substances; and XA and XB are their mole fractions.
The total pressure is given by Dalton’s law of partial pressure
 Colligative Properties of Nonelectrolyte Solutions
Vapor-Pressure Lowering
If both components of a solution are volatile
In a solution of benzene and toluene, the
vapor pressure of each component obeys
Raoult’s law.

The benzene-toluene solution is one of

the few examples of an ideal solution,
which is any solution that obeys Raoult’s The dependence of the partial
pressures of benzene and toluene
law. on their mole fractions in a
benzene-toluene solution (Xtoluene
=1 - Xbenzene ) at 80°C. This
One characteristic of an ideal solution is solution is said to be ideal
that the heat of solution, ΔHsoln, is zero because the vapor pressures
obey Raoult’s law.
 Colligative Properties of Nonelectrolyte Solutions
Vapor-Pressure Lowering

Nonideal solutions.
(a) Positive deviation occurs when PT is greater than that predicted
by Raoult’s law (the solid black line).
(b) Negative deviation. Here, PT is less than that predicted by
Raoult’s law (the solid black line).
Calculate the vapor pressure of a solution made by dissolving 218 g of
glucose (molar mass=180.2 g/mol) in 460 mL of water at 30°C. What is
the vapor-pressure lowering? The vapor pressure of pure water at 30°C
is 31.82 mmHg. Assume the density of the solution is 1.00 g/mL.

The mole fraction of

water, X1, is given by
 Colligative Properties of Nonelectrolyte Solutions
Boiling-Point Elevation
The boiling point of a solution is the temperature at which its vapor
pressure equals the external atmospheric pressure. Because the presence
of a nonvolatile solute lowers the vapor pressure of a solution, it must also
affect the boiling point of the solution.

The boiling point elevation (∆Tb ) is defined as the boiling point of the
solution (Tb) minus the boiling point of the pure solvent (T°b):

The value of ∆Tb is proportional to the vapor-pressure lowering, and so it is

also proportional to the concentration (molality) of the solution. That is,

where m is the molality of the solution and Kb is the molal boiling-

point elevation constant. The units of Kb are °C/m.
 Colligative Properties of Nonelectrolyte Solutions
Freezing-Point Depression
The freezing point depression (∆Tf ) is defined as the freezing point of the
pure solvent (T°f) minus the freezing point of the solution (Tf ):

where m is the concentration of the solute in molality units, and Kf is the

molal freezing-point depression constant. Like Kb, Kf has the units °C/m.
 Colligative Properties of Nonelectrolyte Solutions
Freezing-Point Depression

Freezing involves a transition from the disordered state to

the ordered state.

For this to happen, energy must be removed from the

system.

Because a solution has greater disorder than the solvent,

more energy needs to be removed from it to create order
than in the case of a pure solvent.

Therefore, the solution has a lower freezing point than its

solvent.
 Colligative Properties of Nonelectrolyte Solutions
Freezing-Point Depression

Ice on frozen roads and sidewalks melts when sprinkled with salts such as
NaCl or CaCl2. This method of thawing succeeds because it depresses the
freezing point of water.

Sprinkling salt over ice

 Phase diagram

Phase diagram illustrating the boiling-point elevation and freezing-

point depression of aqueous solutions. The dashed curves pertain to
the solution, and the solid curves to the pure solvent.
boiling point of the solution is higher than that of water, and
the freezing point of the solution is lower than that of water.
Calculate the freezing point of a solution containing 651 g of ethylene glycol in
2505 g of water. Would you keep this substance in your car radiator during the
summer? The molar mass of ethylene glycol is 62.01 g and kf= 1.86°C/m, Kb=
0.52 °C/m).

Because pure water freezes at 0°C, the solution will freeze at (0 -7.79)°C or -7.79°C.
 Colligative Properties of Nonelectrolyte Solutions
Osmotic Pressure

The selective passage of solvent molecules through a porous membrane

from a dilute solution to a more concentrated one. a semipermeable
membrane, which allows the passage of solvent molecules but blocks the
passage of solute molecules.
The osmotic pressure (p) of a solution is the pressure required to stop
osmosis. The osmotic pressure of a solution is given by
 Colligative Properties of Nonelectrolyte Solutions
Osmotic Pressure

Osmotic pressure is directly proportional to the concentration

of solution.

If two solutions are of equal concentration and, hence,

have the same osmotic pressure, they are said to be
isotonic.

If two solutions are of unequal osmotic pressures, the

more concentrated solution is said to be hypertonic

and the more dilute solution is described as hypotonic

 Colligative Properties of Nonelectrolyte Solutions
Osmotic Pressure Hemolysis

To study the contents of red

blood cells, which are
protected from the external
environment by a
semipermeable membrane,
biochemists use a technique
called hemolysis. The red
blood cells are placed in a
hypotonic solution. Because
the hypotonic solution is less
concentrated than the interior
of the cell, water moves into
the cells. The cells swell and
 Colligative Properties of Nonelectrolyte Solutions
Osmotic Pressure Preservation of Jam and Jelly

A large quantity of sugar is actually essential to the preservation process

because the sugar helps to kill bacteria that may cause botulism. When a
bacterial cell is in a hypertonic (high-concentration) sugar solution, the
intracellular water tends to move out of the bacterial cell to the more
concentrated solution by osmosis. This process, known as crenation, causes
the cell to shrink and, eventually, to cease functioning.
 Colligative Properties of Nonelectrolyte Solutions
Osmotic Pressure Transporting water upward in plants

Osmotic pressure also is the major mechanism for transporting water

upward in plants. Because leaves constantly lose water to the air, in a
process called transpiration, the solute concentrations in leaf fluids
increase. Water is pulled up through the trunk, branches, and stems of
trees by osmotic pressure. Up to 10 to 15 atm pressure is necessary to
transport water to the leaves at the tops of California’s redwoods, which
reach about 120 m in height.
The formula for low-molecular-mass starch is (C6H10O5)n, where n
averages 200. When 0.798 g of starch is dissolved in 100.0 mL of
water solution, what is the osmotic pressure, in mmHg, at 25oC?
Using Colligative Properties to Determine Molar Mass
Camphor is a white solid that melts at 179.5oC and freezing-point-
depression constant is 40oC/m. a. A 1.07-mg sample of a compound was
dissolved in 78.1 mg of camphor. The solution melted at 176.0 oC. What
is the molecular mass of the compound? b. If the empirical formula of
the compound is CH, what is the molecular formula?

moles of the compound that are dissolved in 78.1 mg of camphor

 Colligative Properties of Electrolyte Solutions
Electrolytes dissociate into ions in solution, and so one unit of an
electrolyte compound separates into two or more particles when it
dissolves. total number of solute particles that determines the
colligative properties of a solution. 0.1 m CaCl2 solution to depress
the freezing point by three times as much as a 0.1 m sucrose solution
because each CaCl2 produces three ions. To account for this effect we
define a quantity called the van’t Hoff factor, given by

Consequently, the equations for colligative properties must be modified as

Colligative Properties of Electrolyte Solutions
In reality, the colligative properties of electrolyte solutions are usually
smaller than anticipated because at higher concentrations,
electrostatic forces come into play and bring about the formation of
ion pairs. An ion pair is made up of one or more cations and one or
more anions held together by electrostatic forces. The presence of an
ion pair reduces the number of particles in solution, causing a
reduction in the colligative properties.

Electrolytes containing multicharged ions such as Mg2+, Al3+, SO4 2-,

and PO4 3- have a greater tendency to form ion pairs than electrolytes
such as NaCl and KNO3, which are made up of singly charged ions.
Estimate the freezing point of a 0.010 m aqueous solution of aluminum
sulfate, Al2(SO4)3. Assume the value of i based on the formula of the
compound.
 Colloids
A colloid is a dispersion of particles of one substance (the
dispersed phase) throughout another substance or solution
(the continuous phase).

Fog is an example of a colloid: it consists of very small water

droplets (dispersed phase) in air (continuous phase).

A colloid differs from a true solution in that the dispersed

particles are larger than normal molecules, though they are
too small to be seen with a microscope.

The particles range from about 1×103 pm to about 2 × 105 pm

in size.
 Tyndall Effect
Although a colloid appears to be homogeneous because the dispersed
particles are quite small, it can be distinguished from a true solution by its
ability to scatter light. The scattering of light by colloidal-size particles is
known as the Tyndall effect.

For example, the atmosphere appears to be a clear gas, but a ray of

sunshine against a dark background shows up many fine dust particles by
light scattering.

A light beam is visible perpendicular to its path only if light is

scattered toward the viewer. The vessel on the left contains a colloid,
which scatters light. The vessel on the right contains a true solution,
which scatters negligible light.
LYOPHILIC AND LYOPHOBIC SOLS OR COLLOIDS

Sols are colloidal systems in which a solid is dispersed in a

liquid.
These can be subdivided into two classes :
(a) Lyophilic sols (solvent-loving)
(b) Lyophobic sols (solvent-hating)

Lyophilic sols are those in which the dispersed phase exhibits a

definite affinity for the medium or the solvent.
The examples of lyophilic sols are dispersions of starch, gum, and
protein in water.

Lyophobic sols are those in which the dispersed phase has no

attraction for the medium or the solvent.
The examples of lyophobic sols are dispersion of gold, iron (III)
hydroxide and sulphur in water.
 Hydrophilic and Hydrophobic Colloids
Colloids in which the continuous phase is water are divided into two
major classes:
hydrophilic colloids and hydrophobic colloids.

A hydrophilic colloid is a colloid in which there is a strong attraction

between the dispersed phase and the continuous phase (water). Protein
solutions, such as gelatin in water, are hydrophilic colloids.

A hydrophobic colloid is a colloid in which there is a lack of attraction

between the dispersed phase and the continuous phase (water).
Hydrophobic colloids are basically unstable.

Given sufficient time, the dispersed phase aggregates into larger

particles.

A colloid of gold particles in water prepared by Michael Faraday in 1857

is still preserved in the British Museum in London. This colloid is
hydrophobic as well as a sol (solid particles dispersed in water).
 Coagulation

Coagulation is the process by which the dispersed phase of a colloid is

made to aggregate and thereby separate from the continuous phase.
A positively charged colloidal particle of iron(III) hydroxide gathers a
layer of anions around it. Phosphate ions, for example, gather more
closely to the positively charged colloidal particles than do chloride
ions. If the ion layer is gathered close to the colloidal particle, the
overall charge is effectively neutralized. In that case, two colloidal
particles can approach close enough to aggregate.
 Coagulation

The curdling of milk when it sours

is another example of coagulation.
Milk is a colloidal suspension in
which the particles are prevented
from aggregating because they have
electric charges of the same sign.
The ions responsible for the
coagulation (curdling) are formed
when lactose (milk sugar) ferments
to lactic acid.

A third example is the coagulation

of a colloidal suspension of soil in
river water when the water meets
the concentrated ionic solution of
an ocean. The Mississippi Delta was
formed in this way.
 Association Colloids
When molecules or ions that have both a hydrophobic and a hydrophilic end are
dispersed in water, they associate, or aggregate, to form colloidal-sized particles,
or micelles.
A micelle is a colloidal-sized particle formed in water by the association of
molecules or ions that each have a hydrophobic end and a hydrophilic end. The
hydrophobic ends point inward toward one another, and the hydrophilic ends are
on the outside of the micelle facing the water.
A colloid in which the dispersed phase consists of micelles is called an association
colloid.

Ordinary soap in water provides an example of an association colloid. Soap

consists of compounds such as sodium stearate, C17H35COONa. The stearate ion
has a long hydrocarbon end that is hydrophobic (because it is nonpolar) and a
carboxyl group at the other end that is hydrophilic (because it is ionic).
 Cleansing action of soap

In water solution, the stearate ions associate into micelles in which the
hydrocarbon ends point inward toward one another and away from the
water, and ionic carboxyl groups are on the outside of the micelle facing
the water.
 Cleansing action of soap

The cleansing action of soap occurs because oil and grease can be absorbed into
the hydrophobic centers of soap micelles and washed away.
 Association Colloids

Synthetic detergents also form association colloids. Sodium lauryl sulfate is

a synthetic detergent present in toothpastes and shampoos. It has a
hydrophilic sulfate group (—OSO3-) and a hydrophobic dodecyl group
(C12H25—, the hydrocarbon end).
 Association Colloids

Cationic detergents have a positive charge at the hydrophilic end.