write an assignment on solution and colloids

Chapter 12: Solutions and Colloids

Key Concepts

• 12.1 The Dissolution Process
• 12.2 Electrolytes
• 12.3 Solubility
• 12.4 Colligative Properties
• 12.5 Colloids

Coral reefs are home to about 25% of all marine species. They are being threatened by climate change, oceanic acidification, and water pollution, all of which change the composition of the solution we know as seawater. Dissolved oxygen in seawater is critical for sea creatures, but as the oceans warm, oxygen becomes less soluble. As the concentration of carbon dioxide in the atmosphere increases, the concentration of carbon dioxide in the oceans increases, contributing to oceanic acidification. Coral reefs are particularly sensitive to the acidification of the ocean, since the exoskeletons of the coral polyps are soluble in acidic solutions. Humans contribute to the changing of seawater composition by allowing agricultural runoff and other forms of pollution to affect our oceans.

Solutions are crucial to the processes that sustain life and to many other processes involving chemical reactions. In this module, we will consider the nature of solutions, and examine factors that determine whether a solution will form and what properties it may have. In addition, we will discuss colloids—systems that resemble solutions but consist of dispersions of particles somewhat larger than ordinary molecules or ions.

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Properties and Applications

Classification, colloid examples, preparations.

Colloids are mixtures of two or more substances where one substance is dispersed within another. The dispersed substance is referred to as the “dispersed phase”. The medium in which it is dispersed is known as the “dispersion medium”. [1 – 4]

A colloidal solution typically consists of particles ranging in size from 1 nanometer to 1 micrometer. These particles can be solid, liquid, or gas. Unlike true solutions where solute particles are dissolved at a molecular level, colloidal particles remain suspended in the dispersion medium without settling down due to gravity, thus forming a heterogeneous mixture.

The properties of colloids differ from those observed in true solutions or suspensions. For instance, they often display unique optical effects, such as the Tyndall effect, when light passes through them. This phenomenon occurs because the dispersed particles scatter light rays and make them visible. [1]

Colloids find numerous applications across various industries. They are utilized in areas such as food science (e.g., emulsions), pharmaceuticals (e.g., drug delivery systems), cosmetics (e.g., creams), and even environmental engineering (e.g., wastewater treatment). [1,6]

Colloids are a type of mixture, with one substance dispersed as particles in another substance. These mixtures can be classified based on the particle size and the nature of the dispersed phase. The different categories of colloids include emulsions, foams, aerosols, sols, and gels. [1,2,6]

Emulsions are colloidal systems where liquid droplets are dispersed in another immiscible liquid.

Foams consist of gas bubbles dispersed in a liquid or solid medium. They can be formed by mechanical agitation or trapping gas during formation.

Aerosols refer to colloidal systems where solid or liquid particles are suspended in a gas medium.

Sols are colloidal systems where solid particles are dispersed in a liquid medium. Due to Brownian motion, sols have smaller particle sizes that do not settle over time.

Gels consist of a continuous solid network throughout a liquid medium. The gel network traps the liquid within its structure, giving it a semi-solid consistency.

The table below lists a few examples of colloids. [1,3,6]

Colloids are created through various techniques, two prominent ones being the condensation method and the dispersion method. [2,6]

Condensation Method

The condensation method initiates particle growth from smaller entities. In this method, molecules or ions in a solution gather and agglomerate, forming larger particles. For instance, the precipitation of silver chloride in a silver nitrate and sodium chloride solution exemplifies the condensation technique. The ions collide, creating tiny solid particles that settle, leaving a dispersed colloid in the solution. Similarly, in aerosols, condensation occurs when vapor condenses into small droplets, suspending in the air as mist or fog.

Dispersion Method

The dispersion method fragments larger particles into smaller ones. It resembles sculpting, where larger chunks are chiseled into intricate shapes. This approach involves breaking down bulkier substances into colloidal particles. An example of the dispersion method is the creation of colloids by mechanical means, like milling or grinding. Breaking down a solid into fine particles and dispersing them in a liquid medium exemplifies this method. Another way to form colloids is by electrical means, like the dispersal of droplets in an emulsion by applying an electric field.

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Solutions, Suspensions, Colloids, and Dispersions

The Distinguishing Characteristics That Set These Similar Things Apart

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Solutions, suspensions, colloids, and other dispersions are similar but have characteristics that set each one apart from the others.

A solution is a homogeneous mixture of two or more components. The dissolving agent is the solvent. The substance that is dissolved is the solute. The components of a solution are atoms, ions, or molecules, making them 10 -9 m or smaller in diameter.

Example: Sugar and water

Suspensions

The particles in suspensions are larger than those found in solutions. Components of a suspension can be evenly distributed by mechanical means, like by shaking the contents but the components will eventually settle out.

Example: Oil and water

Particles intermediate in size between those found in solutions and suspensions can be mixed in such a way that they remain evenly distributed without settling out. These particles range in size from 10 -8 to 10 -6 m in size and are termed colloidal particles or colloids. The mixture they form is called a colloidal dispersion . A colloidal dispersion consists of colloids in a dispersing medium.

Example: Milk

Other Dispersions

Liquids, solids, and gasses all may be mixed to form colloidal dispersions.

Aerosols : Solid or liquid particles in a gas Examples: Smoke is solid in a gas. Fog is a liquid in a gas.

Sols : Solid particles in a liquid Example: Milk of Magnesia is a sol with solid magnesium hydroxide in water.

Emulsions : Liquid particles in a liquid Example: Mayonnaise is oil in water .

Gels : Liquids in solid Examples: Gelatin is protein in water. Quicksand is sand in water.

Telling Them Apart

You can tell suspensions from colloids and solutions because the components of suspensions will eventually separate. Colloids can be distinguished from solutions using the Tyndall effect . A beam of light passing through a true solution, such as air, is not visible. Light passing through a colloidal dispersion, such as smoky or foggy air, will be reflected by the larger particles and the light beam will be visible.

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• Colloidal Solution

What are Colloids or Colloidal Solution?

Also known as colloids or colloidal suspension, the colloidal solution can be defined as a mixture of particles of substances. These particles are microscopically dispersed and soluble/insoluble which are suspended in a fluid regularly.

They generally represent a solution system in which the particles comprising that system have a particle size intermediate that of a true solution and a coarse dispersion, roughly ranging between 1nm to 500 nm (or 1nm to 0.5µm). A colloidal solution may be considered as a two-phase (heterogeneous) system under some circumstances, while it may be considered as a one-phase (homogeneous) system under other circumstances. 

Examples of Colloidal Solution

Not all mixtures are known as colloids. The mixtures where the suspended particles don't settle down at the button and get evenly dispersed into another substance are called colloids. Some examples of colloidal solutions are as follows:

Whipped cream

Fire retardant

Perfume 

Classification of Colloids

The colloids are classified based on the following:

1. Based on their Physical State 

Aerosol (air as the dispersion medium), Gels (solid dispersion medium) and Emulsion (liquid-liquid solutions in which the dispersed phase is liquid)

2. Based on their Dispersion Medium

Hydrosol (water acts as a dispersion medium), Alcosol (alcohol acts as a dispersion medium and Acrosol (contains a dispersed phase particle in the air).

3. Based on Interaction Forces

The types of colloidal solutions based on the interaction between the forces of the dispersion medium and dispersed phase are discussed below:

Lyophilic Colloids

The colloidal systems in which the colloidal particles interact to an appreciable extent with the dispersion medium are referred to as the lyophilic colloids. The term lyophilic means solvent loving. Owing to their affinity for the dispersion medium, such materials form colloidal sols. 

The lyophilic colloidal sols are usually obtained by simply dissolving the required material (whose sol is to be prepared) into the solvent that is being used. The most common examples of the formation of sols are dissolving acacia in water , dissolving gelatin in water or dissolving celluloid in amyl acetate. The various properties of this class of colloids are due to the attraction between the dispersed phase and the dispersion medium, which leads to salvation; the attachment of solvent molecules to the molecules of the dispersed phase. If water is taken as the dispersion medium, the colloids prepared are known as hydrophilic colloids. Most lyophilic colloids are organic molecules, for example, gelatin, acacia, insulin, albumin, rubber, and polystyrene. Of these, insulin, albumin, gelatin and acacia produce lyophilic or hydrophilic sols. Rubber and polystyrene form lyophilic colloids in non aqueous, organic solvents. These materials accordingly are referred to as lipophilic colloids. These examples illustrate the important point that the term lyophilic has meaning only when applied to the material dispersed in a specific dispersion medium. A material that forms a lyophilic colloidal system in one liquid (e.g., water) may not do so in another liquid (e.g., benzene). 

Lyophobic Colloids

Lyophobic colloids are composed of substances which have very little attraction, if existing, for the dispersion medium. These are the lyophobic (solvent-hating) colloids and, predictably, their properties differ from those of the lyophilic colloids. This is primarily due to the absence of a solvent sheath around the particle. These types of colloids are generally constituted when inorganic particles are dispersed in water. Examples of such materials are gold, silver, sulfur, arsenious sulfide, and silver iodide. Unlike lyophilic colloids, lyophobic colloids require special methods of preparation. These include two types of methods. First, dispersion methods, in which size reduction of coarse particles is done, and second, condensation method, which requires the aggregation of small-sized particles to form bigger particles which lie within colloidal size range.

Association Colloids

Association or amphiphilic colloids are the third type of colloidal systems. In these types of colloids, certain molecules or ions, termed amphiphiles or surface-active agents, are characterized by having two distinct regions of opposing solution affinities within the same molecule or ion. They have one polar region which is attracted towards a polar solvent and within the same molecule; they have a no-polar region which is attracted towards the non-polar solvent. These amphiphiles can arrange themselves according to the type of solution (polar or nonpolar) they are put into. When they are put into a polar solution, they expose their polar regions towards the solvent while covering their non-polar regions towards the inner core, and vice versa. When present in a liquid medium at low concentrations, the amphiphiles exist separately and are of such a size as to be sub colloidal. As the concentration of amphiphiles increases, they start to aggregate faster. These aggregates may comprise of 50 or more amphiphiles and are called micelles. 

Preparation of Colloidal Solution

There are mainly two major ways for preparation of colloidal solution, i.e., by condensation method (chemical techniques) and by dispersion method (physical techniques).

1. Condensation Method : Preparation of colloidal solution by condensation method uses the following chemical techniques:

Double decomposition

Excessive cooling

Exchange of solvent

Change of physical state

2. Dispersion Method: The dispersion method for preparation of colloids mainly includes the following physical methods:

Mechanical dispersion

Bredig’s Arc Method or by Electrical Dispersion

Peptization

Properties of Colloidal Solutions

The colloidal solution exhibit a wide range of properties which are classified into three broad types discussed below:

1. Optical Properties of Colloidal Solutions

The Faraday-Tyndall Effect: When a strong beam of light is passed through a colloidal sol, a visible cone, resulting from the scattering of light by the colloidal particles, is formed. This is the Faraday–Tyndall effect.

Elicitation in Electron Microscope: The electron microscope, capable of yielding pictures of the actual particles, even those approaching molecular dimensions, is now widely used to observe the size, shape, and structure of colloidal particles. The success of the electron microscope is due to its high resolving power, which can be defined in terms of ‘d’, the smallest distance by which two objects are separated and yet remain distinguishable. The smaller the wavelength of the radiation used, the smaller is ‘d’ and the greater is the resolving power. The source of radiation for the optical microscope is visible light which can resolve only two particles at a time of about 20 nm (200 Å). The radiation source of the electron microscope is a beam of high energy electrons having wavelengths in the region of 0.01 nm (0.1 Å).

Light Scattering: Light scattering property of the Colloidal solution particles is based on the Faraday-Tyndall Effect, discussed above. A perfect example of this is the blue color of the sky which is visible to our eyes due to the scattering of the light of blue wavelength by the colloidal particles present in the atmosphere. This property of colloidal particles is used to determine their molecular weight.

2. Kinetic Properties of Colloidal Solutions

Brownian Motion: Brownian motion describes the random movement of colloidal particles. The erratic motion, which may be observed with particles as large as about 5 µm, was explained as resulting from the bombardment of the particles by the molecules of the dispersion medium. The motion of the molecules cannot be observed because they are too small to see. The velocity of the particles increases with decreasing particle size. Increasing the viscosity of the medium, which may be accomplished by the addition of glycerin, decreases and finally stops the Brownian movement.

Diffusion: Colloidal particles diffuse spontaneously from a region of higher concentration to one of lower concentration until the concentration of the system is uniform throughout. Diffusion is a direct result of the Brownian movement. Diffusion of the colloidal particles are governed by a law known as Fick’s first law of diffusion which states that the amount of substance diffusing at a particular time across a plane of area is directly proportional to the change of concentration across both sides.

Osmotic Pressure: The osmotic pressure of the colloidal particles is described by the Van’t Hoff Equation, π = cRT, where π is the osmotic pressure, c is the concentration of the solute in the system, R is the universal gas constant and T is the temperature. According to this equation, the osmotic pressure of the colloidal particles is directly proportional to all these components.

Sedimentation: The colloidal particles do not have any tendency to sediment because the particles are constantly in Brownian motion, as already discussed. This Brownian motion in the colloidal particles is enough to combat the gravitational force applied on them. Hence, a stronger force must be applied to bring about the sedimentation of colloidal particles in a quantitative and measurable manner. This is accomplished by use of the ultracentrifuge which can produce a force one million times that of gravity.

Viscosity: Viscosity is an expression of resistance to the flow of a system under an applied stress. If a liquid is more viscous, a greater amount of force is required to initiate its flow and regulate it at a particular rate. The viscosity of the colloidal solution is given by an equation developed by Einstein, η = ηo(1 + 2.5ϕ) where y, ηo is the viscosity of the dispersion medium, η is the viscosity of the dispersion and φ is the volume fraction.

3. Electrical Properties of Colloidal Solutions

Electrokinetic Phenomena: The movement of a charged surface with respect to an adjacent liquid phase is the basic principle underlying four electro-kinetic phenomena: electrophoresis , electro-osmosis, sedimentation potential, and streaming potential. Electrophoresis is a phenomenon of the movement of charged particles in a liquid medium upon application of a potential difference. Electro-osmosis is a phenomenon in which the application of a potential causes a charged particle to move relative to the liquid, which is stationary. Sedimentation potential is the production of a potential difference when charged particles undergo sedimentation. The streaming potential differs from electro-osmosis in that forcing a liquid to flow through a plug or bed of particles creates the potential.

Donnan Membrane Equilibrium: If sodium chloride is placed in a solution on one side of a semipermeable membrane and a negatively charged colloid together with its counter ions R-Na+ is placed on the other side, the sodium and chloride ions can pass freely across the barrier but not the colloidal anionic particles.

Important Questions

1. What are Foams? Give Examples. 

Ans: Foam is a gas-liquid solution where the dispersed medium is the gas. Example- shaving cream, whipped cream.

2. What is the Difference Between Lyophilic Colloids and Lyophobic Colloids?

Ans:  Lyophilic colloids are reversible solutions with a strong interaction between the dispersed phase and dispersion medium. They have high stability and are resistant to coagulation. Whereas, Lyophobic colloids are irreversible solutions.They are unstable and have weak Van Der Waals forces of attraction between dispersed phase and dispersion medium. As a result,they are easy to coagulate. 

3. What are Gels? Give an Example.

Ans: Gels are a type of sols consisting of two or more phases, where the solid is dispersed into the liquid medium.

4. Which Physical Method is Used for Preparation of Metallic Sols? 

Ans: Bredig's dispersion method is used for the preparation of metallic sols such as gold sol where the gold particles are broken down so that it acquires the size of sol particles. Those particles are then immersed in the required dispersion medium for formation of sols.

5. What Causes an Emulsion?

Ans: When  two insoluble liquids are mixed together (in the form of drops) to disperse one liquid into the other it leads to the formation of emulsion. They can be oil-in-water or water-in-oil depending on the continuous phase. 

FAQs on Colloidal Solution

1. Elaborate in brief about Interaction between particles.

In the interaction of colloid particles, the following forces are important: Excluded volume repulsion: This relates to the inability of any solid particles to collide. Interaction between electrostatic charges: Colloidal particles frequently have an electrical charge, which causes them to attract or repel one another. This interaction is influenced by the charge of both the constant and scattered stages, as well as the movement of the stages. Forces of van der Waals: This is owing to the interplay of two dipoles, one of which is stable and the other is generated. Even though the particles do not have a persistent dipole, variations in the electron density cause the particle to form a transient dipole. This brief dipole causes a dipole in adjacent particles. The induced dipoles and the transient dipoles are then attracted to one another. 

2. What is the stability of a colloidal system?

A colloidal system's stability is characterized by particles staying stable in solution and is determined by particle contact forces. Electrostatic contacts and van der Waals forces are examples of this. They both add to the system's total potential energy. The interaction energy owing to positive forces between colloidal particles is less than kT, where k is the Boltzmann constant and T is the ultimate temperature. If this is the situation, the colloidal particles will resist or only faintly attract one another, leaving the item suspended. The favorable forces will dominate if the contact energy is greater than kT, and the colloidal particles will begin to cluster around. Aggregation is the usual term for this phenomenon, however, it is also known as flocculation, coagulation, or precipitation. 

3. How can destabilization be accomplished? 

Destabilization can be achieved in a variety of ways: 

The electrical boundary that inhibits particle aggregation is removed. This can be performed by reducing the particle's Debye screening length by adding salt to the suspension. This can also be accomplished by adjusting the pH of a suspension. This efficiently neutralizes the surface charge of the suspended particles. This eliminates the repulsive forces that hold colloidal particles apart, allowing van der Waals forces to cause aggregation. Minor pH changes can cause a considerable change in the zeta potential. 

Rapid coagulation or aggregation tends to happen when the zeta potential falls under a particular threshold, often about 5mV. A charged polymer flocculant is added. Individual colloidal particles can be bridged by attractive electrostatic interactions in polymer flocculants. 

The inclusion of a positively charged polymer, for instance, can flocculate negatively charged colloidal silica or clay particles. Depletants are non-adsorbed polymers that cause aggregation due to the entropic effect.

4. How to monitor stability? 

Multiple light scattering mixed with vertical scanning is the most extensively applied method for monitoring a product's dispersion status. This also helps in identifying and quantifying the destabilizing processes.  Turbidimetry is a technique that measures the fraction of light that is backscattered by colloidal particles after passing through the sample. The average particle size and volume fraction of the dispersed stage are directly related to the dispersion intensity. As a result, local variations in concentration caused by sedimentation or creaming, as well as aggregation-induced particle clumping, are identified and monitored. These occurrences are linked to colloids that are volatile.

5. What is a colloidal crystal?

A colloidal crystal is a finely organized array of particles that can be created across a large distance and that resemble their atomic or molecular equivalents. One of the best natural illustrations of this ordering phenomenon may be seen in precious opal, where close-packed domains of amorphous colloidal spheres of silicon dioxide produce bright regions of pure spectrum hue. After years of deposition and pressure under hydrostatic and gravitational forces, these spherical particles precipitate in extremely siliceous pools in Australia and elsewhere, forming these highly orderly arrays. This operates as a natural diffraction grating for visible light waves, especially when the interstitial spacing is on par with the incident light wave.

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• Properties of Colloidal Solutions

A colloid is a mixture in which one substance of microscopically dispersed insoluble particles is suspended throughout another substance. Owing to this peculiar structure of colloid, it has varied physical and chemical properties. Let us explore more about the physical, chemical, optical as well as electrical properties of colloidal solutions.

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Physical Properties of Colloidal Solutions

• Stability : Colloids are relatively stable in nature. The particles of the dispersed phase are in a state of continuous motion and remain suspended in the solution.
• Filterability :   Colloids require specialized filters known as ultrafilters for filtration. They readily pass through ordinary filter papers without yielding any residue.
• Heterogenous nature : Since colloids consist of two phases, the dispersed phase as well as the dispersion medium, they are known as heterogeneous in nature.
• Homogenous appearance : Even though colloids have suspended particles and are heterogeneous in nature, they appear as if it is a homogenous solution. This is because the suspended particles are so tiny that they are not visible by the naked eye.

Browse more Topics under Surface Chemistry

• Adsorption Isotherm
• Classification of Colloids
• Preparation of Colloids
• Shape-selective Catalysis by Zeolites

Colligative Properties

The particles of the dispersed phase come together to form associate molecules. The formation of these associate molecules renders the solution certain special properties such as

• a decrease in vapor pressure
• elevation in boiling point
• depression in freezing point
• a decrease in osmotic pressure

Read the Classification of Colloids here .

Optical Properties of Colloidal Solutions: Tyndall Effect

Colloids exhibit a phenomenon known as the Tyndall effect observed by Tyndall in 1869. When we pass an intense converging beam of light through a colloidal solution kept in dark, the path of the beam gets illuminated with a bluish light. This phenomenon of scattering of light by colloidal particles is called the Tyndall effect and the illuminated path is known as the Tyndall cone. The dispersed colloidal particles scatter the light falling on them resulting in emissions that are comparable to ultraviolet and visible radiations. These scattered radiations get illuminated.

The zone of scattered light is observed to be much larger than the particle itself. This makes the colloidal particles to appear as tiny bright spots when viewed under a microscope. This has to be done at right angles to the beam of light.

True solutions do not exhibit a Tyndall effect. This is because the size of particles (ions or molecules) present in a true solution are too small to scatter light. Thus, the Tyndall effect can be used to distinguish a colloidal solution from a true solution.

Learn different types of Emulsion and its properties here.

Mechanical Properties of Colloidal Particles: Brownian Movement

The dispersed particles present in a colloidal solution exhibit a very important property called the Brownian movement. When a colloidal solution is viewed under an ultramicroscope, the colloidal particles are seen continuously moving in a zigzag path.

There is a continuous bombardment of the moving molecules of the dispersion medium on the colloidal particles from all directions. This imparts a momentum to the particles to move in a forward direction where again it collides with another particle. These collisions result in the random zigzag movement of the colloidal particle.

The Brownian movement imparts stability to the sol. It opposes the gravitational force acting on colloidal particles and prevents them from settling down thus maintaining the stability of the sol.

Learn about Chemical Adsorption and its significance here.

Electrical Properties of Colloidal Solutions

The particles of the colloidal solution carry the same type of charge, while the dispersion medium carries an equal and opposite charge. The charge on the dispersion medium balances the charge on dispersed particles and the solution as a whole is electrically neutral.

The dispersed particles of a colloid repel each other since they carry similar charges and this prevents them from settling down thus maintaining the stability of the sol. Based on the nature of the charge, the colloidal sols may be classified as positively charged and negatively charged sols.

Solved Questions For You

Que: Which of these is not a property of colloids?

• Tyndall effect
• Brownian motion
• Heterogenous nature
• High instability

Ans: The correct option is “D”. High instability is not a property of colloids. Rather, colloids are quite stable.

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4.5: Colloids and their Uses

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• Page ID 92719

• Stephen Lower
• Simon Fraser University

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Learning Objectives

• Summarize the principal distinguishing properties of solutions, colloidal dispersions, and suspensions.
• For the various dispersion types (emulsion, gel, sol, foam, etc.), name the type (gas, liquid, or solid) of both the dispersed phase and the dispersions phase.
• Describe the origins of Brownian motion and how it can observed.
• Describe the electric double layer that surrounds many colloidal particles.
• Explain the mechanisms responsible for the stability of lyophilic and lyophobic colloidal dispersions.
• Define: surfactant, detergent, emulsifier, micelle.
• Give some examples of how colloidal dispersions can be made.
• Explain why freezing or addition of an electrolyte can result in the coagulation of an emulsion.
• Describe some of the colloid-related principles involved in food chemistry, such as the stabilization of milk and mayonaisse, the preparation of butter, and the various ways of cooking eggs.
• Describe the role of colloids in wastewater treatment.

Sand, salt, and chalk dust are made up of chunks of solid particles, each containing huge numbers of molecules. You can usually see the individual particles directly, although the smallest ones might require some magnification. At the opposite end of the size scale, we have individual molecules which dissolve in liquids to form homogeneous solutions. There is, however, a vast but largely hidden world in between: particles so tiny that they cannot be resolved by an optical microscope, or molecules so large that they begin to constitute a phase of their own when they are suspended in a liquid. This is the world of colloids which we will survey in this lesson. As you will see, we encounter colloids in the food we eat, the consumer products we buy... and we ourselves are built largely of colloidal materials.

Introducing Colloids

Colloids occupy an intermediate place between [particulate] suspensions and solutions, both in terms of their observable properties and particle size. In a sense, they bridge the microscopic and the macroscopic. As such, they possess some of the properties of both, which makes colloidal matter highly adaptable to specific uses and functions. Colloid science is central to biology, food science and numerous consumer products.

• Solutions are homogeneous mixtures whose component particles are individual molecules whose smallest dimension is generally less than 1 nm. Within this size range, thermal motions maintain homogeneity by overcoming the effects of gravitational attraction.
• Colloidal dispersions appear to be homogeneous, and the colloidal particles they contain are small enough (generally between 1-1000 nm) to exhibit Brownian motion, cannot be separated by filtration, and do not readily settle out. But these dispersions are inherently unstable and under certain circumstances, most colloidal dispersions can be "broken" and will "flocculate" or settle out.
• Suspensions are heterogeneous mixtures in which the suspended particles are sufficiently large (> 1000 nm in their smallest dimension) to settle out under the influence of gravity or centrifugal force. The particles that form suspensions are sometimes classified into various size ranges.

Colloidal particles need not fall within the indicated size range in all three dimensions; thus fibrous colloids such as many biopolymers may be very extended sizes along one direction.

The nature of colloidal particles

To begin, you need to recall two important definitions:

• a phase is defined as a region of matter in which the composition and physical properties are uniform. Thus ice and liquid water, although two forms of the single substance H 2 O, constitute two separate phases within a heterogeneous mixture .
• A solution is a homogeneous mixture of two or more substances consisting of a single phase. (Think of sugar dissolved in water).

But imagine that you are able to shrink your view of a solution of sugar in water down to the sub-microscopic level at which individual molecules can be resolved: you would see some regions of space occupied by H 2 O molecules, others by sugar molecules, and likely still others in which sugar and H 2 O molecules are momentarily linked together by hydrogen bonding— not to mention the void spaces that are continually appearing and disappearing between molecules as they are jumbled about by thermal motions. As with so many simple definitions, the concept of homogeneity (and thus of a solution ) breaks down as we move from the macro-scale into the molecular scale. And it is the region in between these two extremes that constitutes the realm of the colloid.

Smaller is bigger

What makes colloidal particles so special is not so much their sizes as it is the manner in which their surface areas increase as their sizes decrease. If we take a sample of matter and cut it up into smaller and smaller chunks, the total surface area will increase very rapidly. Although mass is conserved, surface area is not; as a solid is sliced up into smaller bits, more surfaces are created. These new surfaces are smaller, but there are many more of them; the ratio of surface area to mass can become extremely large.

The total surface area increases as the inverse cube of the the face length, so as we make our slices still smaller, the total surface area grows rapidly. In practical situations with real colloids, surface areas can reach hectares (or acres) per mole!

Why do we focus so much attention on surface area? The general answer is that surfaces (or more generally, interfaces between phases) possess physical and chemical properties of their own. In particular,

• Surfaces can exert van der Waals attractive forces on other molecules near them, and thus loosely bind other particles by adsorption
• Interfaces between different phases usually give rise to imbalances in electrical charge which can cause them to interact with nearby ions.
• The surfaces of many solids present "broken bonds" which are chemically active.

In normal "bulk" matter, these properties are mostly hidden from us owing to the small amount of surface area in relation to the quantity of matter. But as the particle size diminishes, surface phenomena begin to dominate their properties. The small sizes of colloidal solids allows the properties of their surfaces to dominate their behavior.

Colloidal Dispersions

Colloidal matter commonly exists in the form of colloidal-sized phases of solids, liquids, or gases that are uniformly dispersed in a separate medium (sometimes called the dispersions phase ) which may itself be a solid, liquid, or gas. Colloids are often classified and given special names according to the particular kinds of phases involved.

Notes on this table:

• Pumice is a volcanic rock formed by the rapid depressurization and cooling of molten lava. The sudden release of pressure as the lava is ejected from the volcano allows dissolved gases to expand, producing tiny bubbles that get frozen into the matrix. Pumice is distinguished from other rocks by its very low density.
• Aerogels are manufactured rigid solids made by removing the liquid from gels, leaving a solid, porous matrix that can have remarkable and useful physical properties. Aerogels based on silica, carbon, alumina and other substances are available.
• Milk is basically an emulsion of butterfat droplets dispersed in an aqueous solution of carbohydrates.
• Opal consists of droplets of liquid water dispersed in a silica (SiO 2 ) matrix.

Large molecules can behave as colloids

Very large polymeric molecules such as proteins, starches and other biological polymers, as well as many natural polymers, exhibit colloidal behavior. There is no clear point at which a molecule becomes sufficiently large to behave as a colloidal particle.

Macroscopic or microscopic?

Colloidal dispersions behave very much like solutions in that they appear to be homogeneous on a macroscopic scale. They are often said to be microheterogeneous . The most important feature that distinguishes them from other particulate matter is that:

Colloids dispersed in liquids or gases are sufficiently small that they do not settle out under the influence of gravity. This, together with the their small sizes which allows them to pass through most filters, makes it difficult to separate colloidal matter from the phase in which it is dispersed.

Optical properties of colloidal dispersions

Colloidal dispersions are distinguished from true solutions by their light-scattering properties. The nature of this scattering depends on the ratio of the particle size in the medium to the wavelength of the light. A collimated beam of light passing through a solution composed of ordinary molecules ( r ) tends retain its shape. When such a beam is directed through a colloidal dispersion, it spreads out ( left container ).→

John Tyndall discovered this effect in 1869. Tyndall scattering (as it is commonly known) scatters all wavelengths equally. This is in contrast to Rayleigh scattering , which scatters shorter wavelengths more, bringing us blue skies and red sunsets. Tyndall scattering can be seen even in dispersions that are transparent. As the density of particles (or the particle size) increases, the light scattering may become great enough to produce a "cloudy" effect, as in this image of a smoke-filled room. This is the reason that milk, fog, and clouds themselves appear to be white. The individual water droplets in clouds (or the butterfat droplets in milk) are actually transparent, but the intense light scattering disperses the light in all directions, preventing us from seeing through them.

The Ultramicroscope

Colloidal particles are, like molecules, too small to be visible though an ordinary optical microscope. However, if one looks in a direction perpendicular to the light beam, a colloidal particle will "appear" over a dark background as a tiny speck due to the Tyndall scattering. A microscope specially designed for this application is known as an ultramicroscope . Bear in mind that the ultramicroscope (invented in Austria in 1902) does not really allow us to "see" the particle; the scattered light merely indicates where it is at any given instant.

Brownian motion

If you observe a single colloidal particle through the ultramicroscope, you will notice that it is continually jumping around in an irregular manner. These movements are known as Brownian motion. Scottish botanist Robert Brown discovered this effect in 1827 when observing pollen particles floating in water through a microscope. (Pollen particles are larger than colloids, but they are still small enough to exhibit some Brownian motion.)

It is worth noting that Albert Einstein's analysis of Brownian motion in 1901 constituted the first proof of the molecular theory of matter. Brownian motion arises from collisions of the liquid molecules with the solid particle. For large particles, the millions of collisions from different directions cancel out, so they remain stationary. The smaller the particle, the smaller the number of surrounding molecules able to collide with it, and more likely that random fluctuations will occur in the number of collisions from different sides. Simple statistics predicts that every once in a while, the imbalance in collisions from different directions will become great enough to give the particle a real kick!

Electrical Properties of Colloids

In general, differences in electric potential exist between all phase boundaries. If you have studied electrochemistry, you will know that two dissimilar metals in contact exhibit a "contact potential", and that similar potential differences exist between a metal and a solution in which it is immersed. But this principle extends well beyond ordinary electrochemistry; there are small potential differences even at the water-glass interface in a drinking glass, and the water-air interface above it.

Colloids are no exception to this rule; there is always a difference in electric potential between the colloid "phase" and that of the surrounding liquid. Even if the liquid consists of pure water, the polar H 2 O molecules at the colloid's surface are likely to be predominantly oriented with either their oxygen (negative) or hydrogen (positive) ends facing the interface, depending on the electrical properties of the colloid particle itself.

Interfacial electrical potential differences can have a variety of origins:

• Particles composed of ionic or ionizable substances usually have surface charges due to adsorption of an ion (usually an anion) from the solution, or to selective loss of one kinds of ion from the crystal surface. For example, Ag+ ions on the surface of a silver iodide crystal go into solution more readily than the Br- ions, leaving a negatively-charged surface.
• The charges of amphiprotic groups such as those on the surfaces of metal oxides and hydroxides will vary with the pH of the aqueous medium. Thus a particle of a metal oxide M–O will become positive in acidic solution due to formation of M–OH + , while that of a sparingly soluble hydroxide M–OH will become negative at low pH as it changes to M–O – . Colloidal-sized protein molecule can behave in a similar manner owing to the behavior of amphiprotic carboxylate-, amino- and sulfhydryl groups.
• Non-ionic particles or droplets such as oils or latex will tend to selectively adsorb positive or negative ions present in solution, thus "coating themselves" with electrical charge.
• In clays and other complex structures, isomorphous replacement of one ion by another having a different charge will leave a net electric charge on the particle. Thus particles of kaolinite clay become negatively charged due to replacement of some of the Si 4 + ions by Al 3 + .

Charged colloidal particles will attract an excess of oppositely-charged counter-ions to their vicinity from the bulk solution, forming a localized "cloud" of compensating charge around each particle. The entire assembly is called an electric double layer. Electric double layers of one kind or another exist at all phase boundaries, but those associated with colloids are especially important.

Stability of colloidal dispersions

What keeps the colloidal particles suspended in the dispersion medium? How can we force the particles to settle out? These are very important practical matters:

• Colloidal products such as paints and many foods (e.g., milk) must remain in dispersed form if they are to be useful;
• Other dispersions, often those formed as by-products of operations such as mining, water treatment, paper-manufacture, or combustion are environmental nuisances. The only practical way of disposing them is to separate the colloidal material from the much greater volume of the dispersion medium (most commonly water). Simple evaporation of the water is usually not a practical option; it is generally too slow, or too expensive if forced by heating.

You will recall that weak attractive forces act between matter of all kinds. These are known generally as van der Waals and dispersion forces, and they only "take hold" at very close distances. Countering these is the universal repulsive force that acts at even shorter distances, but is far stronger; it is the basic reason why two atoms cannot occupy the same space. For very small atomic and molecular sized particles, another thing that keeps them apart is thermal motion. Thus when two molecules in a gas collide, they do so with more than enough kinetic energy to overcome the weak attractive forces between them. As the temperature of the gas is reduced, so is the collisional energy; below its boiling point, the attractive forces dominate and the gas condenses into a liquid.

Electrical forces help keep colloids dispersed

When particles of colloidal dimension suspended in a liquid collide with each other, they do so with much smaller kinetic energies than is the case for gases, so in the absence of any compensating repulsion forces, we might expect van der Waals or dispersion attractions to win out. This would quickly result in the growth of aggregates sufficiently large to exceed colloidal size and to fall to the bottom of the container. This process is called coagulation .

So how do stable dispersions such as sols manage to survive? In the preceding section, we saw that each particle with its double layer is more or less electrically neutral. However, when two particles approach each other, each one "sees" mainly the outer part [shown here in blue] of the double layer of the other. These will always have the same charge sign (which depends on the type of colloid and the nature of the medium), so there will be an electrostatic repulsive force that opposes the dispersion force attractions.

Electrostatic (coulombic) forces have a strong advantage in this respect because they act over much greater distances do van der Waals forces. But as we will see further on, electrostatic repulsion can lose its effectiveness if the ionic concentration of the medium is too great, or if the medium freezes. Under these conditions, there are other mechanisms that can stabilize colloidal dispersions.

Interactions with the solvent

Colloids can be divided into two general classes according to how the particles interact with the dispersions medium (often referred to as the "solvent").

Lyophilic colloids

In one class of colloids, called lyophilic ("solvent loving") colloids, the particles contain chemical groups that interact strongly with the solvent, creating a sheath of solvent molecules that physically prevent the particles from coming together. Ordinary gelatine is a common example of a lyophilic colloid. It is in fact hydrophilic , since it forms strong hydrogen bonds with water. When you mix Jell-O or tapioca powder to make a gelatine dessert, the material takes up water and forms a stable colloidal gel. Lyophilic (hydrophilic) colloids are very common in biological systems and in foods.

Lyophobic colloids

Most of the colloids in manufactured products exhibit very little attraction to water: think of oil emulsions or glacially-produced rock dust in river water. These colloids are said to be lyophobic . Lyophobic colloids are all inherently unstable; they will eventually coagulate . However, "eventually" can be a very long time (the settling time for some clay colloids in the ocean is 200-600 years!).

For systems in which coagulation proceeds too rapidly, the process can be slowed down by adding a stabilizer. Stabilizers can act by coating the particles with a protective layer such as a polymer as described immediately below, or by providing an ion that is selectively adsorbed by the particle, thereby surrounding it with a charged sheath that will repel similar particles it collides with. Dispersions of these colloids are stabilized by electrostatic repulsion between the electric double layers surrounding the particles which we discussed in the preceding section.

Stabilization by cloaking

"Stabilization by stealth" has unwittingly been employed since ancient times through the use of natural gums to stabilize pigment particles in inks, paints, and pottery glazes. These gums are also widely used to stabilize foods and personal care products. A lyophobic colloid can be made to masquerade as lyophilic by coating it with something that itself possesses suitable lyophilic properties.

Steric stabilization

Alternatively, attaching a lyophobic material to a colloid of any type can surround the particles with a protective shield that physically prevents the particles from approaching close enough to join together. This method usually employs synthetic polymers and is often referred to as steric stabilization .

Synthetic polymers , which can be tailor-made for specific applications, are now widely employed for both purposes. The polymer can be attached to the central particle either by simple adsorption or by chemical bond formation.

Surfactants and micelle formation

Surfactants and detergents are basically the same thing. Surfactants that serve as cleaning agents are commonly called detergents (from L. detergere "to wipe away, cleanse"). Surfactants are molecules consisting of a hydrophylic "head" connected to a hydrophobic chain. Because such molecules can interact with both "oil" and water phases, they are often said to be amphiphilic . Typical of these is the well known cleaning detergent sodium dodecyl sulfonate ("sodium laurel sulfate") CH 3 (CH 2 ) 11 OSO 3 – Na + .

Amphiphiles possess the very important property of being able to span an oil-water interface. By doing so, they can stabilize emulsions of both the water-in-oil and oil-in-water types. Such molecules are essential components of the lipid bilayers that surround the cells and cellular organelles of living organisms.

Emulsions are inherently unstable; left alone, they tend to separate into "oil" and "water" phases. Think of a simple salad dressing made by shaking vegetable oil and vinegar. When a detergent-like molecule is employed to stabilize an emulsion, it is often referred to as an emulsifier . The resulting structure (left) is known as a micelle .

Emulsifiers are essential components of many foods. They are widely employed in pharmaceuticals, consumer goods such as lotions and other personal care products, paints and printing inks, and numerous industrial processes.

How detergents remove "dirt"

The "dirt" we are trying to remove consists of oily or greasy materials whose hydrophobic nature makes them resistant to the action of pure water. If the water contains amphiphilic molecules such as soaps or cleaning detergents that can embed their hydrophobic ends in the particles, the latter will present a hydrophilic interface to the water and will thus become "solubilized".

Soaps and detergents can also disrupt the cell membranes of many types of bacteria, for which they serve as disinfectants . However, they are generally ineffective against viruses, which do not possess cell membranes.

Bile: your body's own detergent

Oils and fats are important components of our diets, but being insoluble in water, they are unable to mix intimately with the aqueous fluid in the digestive tract in which the digestive enzymes are dissolved. In order to enable the lipase enzymes (produced by the pancreas) to break down these lipids into their component fatty acids, our livers produce a mixture of surfactants known as bile . The great surface area of the micelles in the resulting emulsion enables efficient contact between the lipase enzymes and the lipid materials.

The liver of the average adult produces about 500 mL of bile per day. Most of this is stored in the gall bladder, where it is concentrated five-fold by removal of water. As partially-digested material exits the stomach, the gall bladder squeezes bile into the top of the small intestine (the duodenum ).

In addition to its action as a detergent (which also aids in the destruction of bacteria that may have survived the high acidity of the gastric fluid), the alkaline nature of the bile salts neutralizes the acidity of the stomach exudate. The bile itself consists of of salts of a variety of bile acids, all of which are derived from cholesterol. The cholesterol-like part of the structure is hydrophobic, while the charged end of the salt is hydrophilic.

Microemulsions

Ordinary emulsions are inherently unstable; they do not form spontaneously, and once formed, the drop sizes are sufficiently large to scatter light, producing a milky appearance. As time passes, the average drop size tends to increase, eventually resulting in gravitational separation of the phases.

Microemulsions, in contrast, are thermodynamically stable and can form spontaneously. The drop radii are at the very low end of the colloidal scale, often 100 nm or smaller. This is too small to appreciably scatter visible light, so microemulsions appear visually to be homogenous systems.

Microemulsions require the presence of one or more surfactants which increase the flexibility and stability of the boundary regions. This allows them to vary form smaller micelles than surface tension forces would ordinarily allow; in some cases they can form sponge-like bicontinuous mixtures in which "oil" and "water" phases extend throughout the mixture, affording more contact area between the phases.

The uses of microemulsions are quite wide-ranging, with drug delivery, polymer synthesis, enzyme-assisted synthesis, coatings, and enhanced oil recovery being especially prominent.

4 Making and breaking colloidal dispersions

Particles of colloidal size can be made in two general ways:

• Start with larger particles and break them down into smaller ones ( Dispersion ).
• Build up molecular-sized particles (atoms, ions, or small molecules) into aggregates within the colloidal size range. ( Condensation )

Dispersion processes all require an input of energy as new surfaces are created. For solid particles , this is usually accomplished by some kind of grinding process such as in a ball- or roller-mill. Solids and liquids can also be broken into colloidal dimensions by injecting them into the narrow space between a rapidly revolving shaft and its enclosure, thus subjecting them to a strong shearing force that tends to pull the two sides of a particle in opposite directions.

The application of ultrasound (at about 20 kHz) to a mixture of two immiscible liquids can create liquid-in-liquid dispersions; the process is comparable to what we do when we shake a vinegar-and-oil salad dressing in order to create a more uniform distribution of the two liquids.

How Dispersions are Broken

That oil-in-vinegar salad dressing you served at dinner the other day has now mostly separated into two layers, with unsightly globs of one phase floating in the other. This is surface chemistry in action! Emulsions are fundamentally unstable because molecules near surfaces (i.e., interfaces between phases) are no longer surrounded by their own kind on all sides. The resulting repulsions between like and unlike exact an energetic cost that must eventually be repaid through processes that reduce the interfacial area.

The consequent breakup of the emulsion can proceed through various stages:

• Coalescence - smaller drops join together to form larger ones;
• Flocculation - the small drops stick together without fully coalescing;
• Creamin g - Most oils have lower densities than water, so the drops float to the surface, but may not completely coalesce;
• Breaking - the ultimate thermodynamic fate and end result of the preceding steps.

The time required for these processes to take place is highly variable, and can be extended by the presence of stabilizer substances. Thus milk , an emulsion of butterfat in water, is stabilized by some of its natural components.

Coagulation and flocculation

The processes described above that allow colloids to remain suspended sometimes fail when conditions change, or equally troublesome, they work entirely too well and make it impossible to separate the colloidal particles from the medium; this is an especially serious problem in wastewater settling basins associated with sewage treatment and operations such as mining and the pulp-and-paper industries.

Coagulation is the general term that refers to the "breaking" of dispersions so that the colloidal particles can be collected, usually by settling out. The term Flocculation is often used as a synonym for coagulation, but it is more properly reserved for a special method of effecting coagulation which is described further on. Most coagulation processes act by disrupting the outer (diffuse) part of the electric double layer that gives rise to the electrostatic repulsion between them.

"Do not freeze"

Have you ever encountered milk that had previously been frozen? Not likely something you would want to drink! You will see "Do not freeze" labels on many foodstuffs and on colloidal consumer products such as latex house paint. Freezing disrupts the double layer by causing the ions within it to combine into neutral species so that the particles can now approach closely enough for attractive forces to take over, and once they do so, they never let go: coagulation is definitely an irreversible process!

Addition of an electrolyte

Coagulation of water-suspended dispersions can be brought about by raising the ionic concentration of the medium. The added ions will migrate to the oppositely-charged regions of the double layer, thus neutralizing its charges; this effectively reduces the thickness of the double layer, eventually allowing the attractive forces to prevail.

The coagulated clay accumulates as sediments which eventually form a geographical feature called a river delta.

A liquid phase dispersed in a solid medium is known as a gel , but this formal definition does not always convey the full sense of the nature of the "solid". The latter may start out as a powdery or granulated material such as natural gelatin or a hydrophilic polymer, but once the gel has formed, the "solid" part is less a "phase" than a cross-linked network that extends throughout the volume of the liquid, whose quantity largely defines the volume of the entire gel.

Hydrogels can contain up to 90% water by weight

Most of the gels we commonly encounter have water as the liquid phase, and thus are called hydrogels ; ordinary gelatin deserts are well known examples.

The "solid" components of hydrogels are usually polymeric materials that have an abundance of hydrophilic groups such as hydroxyl (–OH) that readily hydrogen-bond to water and also to each other, creating an irregular, flexible, and greatly-extendable network. These polymers are sometimes synthesized for this purpose, but are more commonly formed by processing natural materials, including natural polymers such as cellulose.

• Gelatine is a protein-like material made by breaking down the connective tissues of animal skins, organs, and bones. The many polar groups on the resulting protein fragments bind them together, along with water molecules, to form a gel.
• A number of so-called super-absorbant polymers derived from cellulose, polyvinyl alcohol and other materials can absorb huge quantities of water, and have found uses for products such as disposable diapers, environmental spill control, water retention media for plants, surgical pads and wound dressings, and protective inner coatings and water-blockers in fiber optics and electrical cables.

Gels are essential components of a huge variety of consumer products ranging from thickening agents in foods and personal care products to cushioning agents in running shoes.

Gels can be fragile!

You may have noticed that a newly-opened container of yogurt or sour cream appears to be smooth and firm, but once some of the material has been spooned out, little puddles of liquid appear in the hollowed-out depressions.

As the spoon is plunged into the material, it pulls the nearby layers of the gel along with it, creating a shearing action that breaks it apart, releasing the liquid. Anyone who has attacked an egg yolk with a cook's whisk, written with a ball-point pen, or spread latex paint on a wall has made use of this phenomenon which is known as shear thinning .

Our bodies are mostly gels

Embedded within the cytosol is the filament-like cytoskeleton which controls the overall shape of the cell and holds the organelles in place.

(In free-living cells such as the amoeba, changes in the cytoskeleton enable the organism to alter its shape and move around to engulf food particles.)

Be thankful for the gels in your body; without them, you would be little more than a bag of gunge-filled liquid, likely to end up as a puddle on the floor!

The individual cells are bound into tissues by the extracellular matrix (ECM) which — on a much larger scale, holds us together and confers an overall structure and shape to the body. The ECM is made of a variety of structural fibers (collagens, elastins) embedded in a gel-like matrix.

6 Applications of colloids

Thickening agents.

The usefulness of many industrial and consumer products is strongly dependent on their viscosity and flow properties. Toothpastes, lotions, lubricants, coatings are common examples. Most of the additives that confer desirable flow properties on these products are colloidal in nature; in many cases, they also provide stabilization and prevent phase separation. Since ancient times, various natural gums have been employed for such purposes, and many remain in use today.

More recently, manufactured materials whose properties can be tailored for specific applications have become widely available. Examples are colloidal microcrystalline cellulose, carboxymethyl cellulose, and fumed silica.

Fumed silica is a fine (5-50 nm), powdery form of SiO 2 of exceptionally low bulk density (as little as 0.002 g cm –3 ); the total surface area of one Kg can be as great as 60 hectares (148 acres). It is made by spraying SiCl 4 (a liquid) into a flame. It is used as a filler, for viscosity and flow control, a gelling agent, and as an additive for strenghthening concrete.

Food colloids

Most of the foods we eat are largely colloidal in nature. The function of food colloids generally has less to do with nutritional value than appearance, texture, and "mouth feel". The latter two terms relate to the flow properties of the material, such as spreadability and the ability to "melt" (transform from gel to liquid emulsion) on contact with the warmth of the mouth.

Dairy products

Milk is basically an emulsion of lipid oils ("butterfat") dispersed in water and stabilized by phospholipids and proteins. Most of the protein content of milk consists of a group known as caseins which aggregate into a complex micellar structure which is bound together by calcium phosphate units.

Homogenizer

The stabilizers present in fresh milk will maintain its uniformity for 12-24 hours, but after this time the butterfat globules begin to coalesce and float to the top ("creaming"). In order to retard this process, most milk sold after the early 1940's undergoes homogenization in which the oil particles are forced through a narrow space under high pressure. This breaks up the oil droplets into much smaller ones which remain suspended for the useful shelf life of the milk.

Before homogenization become common, milk bottles commonly had enlarged tops ↑ to make it easier to skim off the cream that would separate out.

The structures of cream, yogurt and ice cream are dominated by the casein aggregates mentioned above.

Ice cream is a complex mixture of several colloid types:

• an emulsion (of butterfat globules in a highly viscous aquatic phase);
• a semisolid foam consisting of small (100 μ) air bubbles which are beat into the mixture as it is frozen. Without these bubbles, the frozen mixture would be too hard to conveniently eat;
• a gel in which a network of tiny (50 μ) ice crystals are dispersed in a semi-glassy aqueous phase containing sugars and dissolved macromolecules.

Whereas milk is an oil (butterfat)-in-water dispersion, butter and margarine have a "reversed" (water-in-oil) arrangement. This transformation is accomplished by subjecting the butterfat droplets in cream to violent agitation ( churning ) which forces the droplets to coalesce into a semisolid mass within which remnants of the water phase are embedded. The greater part of this phase ends up as the by-product buttermilk .

Eggs: colloids for breakfast, lunch, and dessert

A detailed study of eggs and their many roles in cooking can amount to a mini-course in colloid chemistry in itself. There is something almost magical in the way that the clear, viscous "white" of the egg can be transformed into a white, opaque semi-solid by brief heating, or rendered into more intricate forms by poaching, frying, scrambling, or baking into custards, soufflés, and meringues, not to mention tasty omelettes, quiches, and more exotic delights such as the eggah (Arabic) and kuku (Persian) dishes of the Middle-East.

The raw eggwhite is basically a colloidal sol of long-chain protein molecules, all curled up into compact folded forms due to hydrogen bonding between different parts of the same molecule. Upon heating, these bonds are broken, allowing the proteins to unfold. The denuded chains can now tangle and bind to each other, transforming the sol into a cross-linked hydrogel, now so dense that scattered light changes its appearance to opaque white.

What happens next depends very much on the skill of the cook. The idea is to drive out enough of the water entrapped within the gel network to achieve the desired density while retaining enough gel structure to prevent it from forming a rubbery mass, as usually happens with hard-boiled eggs. This is especially important when the egg structure is to be incorporated into other food components as in baked dishes.

The key to all this is temperature control; the eggwhite proteins begin to coagulate at 65°C and if yolk proteins are present, the mixture is nicely set at about 73°; by 80° the principal (albumin) protein has set, and at much above this the gel network will collapse into an overcooked mass. The temperature limit required to avoid this disaster can be raised by adding milk or sugar; the water part of the milk dilutes the proteins, while sugar molecules hydrogen-bond to them, forming a protective shield that keeps the proton strand separated. This is essential when baking custards, but incorporating a bit of cream into scrambled eggs can similarly help them retain their softness.

Whipped cream and meringues

The other colloidal personalities eggs can display are liquid and solid foams. Instead of applying heat to unfold the proteins, we "beat" them; the shearing force of a whisk or egg-beater helps pull them apart, and the air bubbles that get entrapped in the mixture attract the hydrophobic parts of the unfolded proteins and help hold them in place. Sugar will stabilize the foam by raising its viscosity, but will interfere with protein folding if added before the foam is fully formed. Sugar also binds the residual water during cooking, retarding its evaporation until after the proteins not broken up by beating can be thermally coagulated.

Paints and inks

Paints have been used since ancient times for both protective and decorative purposes. They consist basically of pigment particles dispersed in vehicle — a liquid capable for forming a stable solid film as the paint "dries".

The earliest protective coatings were made by dissolving plant-derived natural polymers (resins) in an oil such as that of linseed. The double-bonds in these oils tends to oxidize when exposed to air, causing it to polymerize into an impervious film. The colloidal pigments were stabilized with naturally-occurring surfactants such as polysaccharide gums.

Present-day paints are highly-engineered products specialized for particular industrial or architectural coatings and for marine or domestic use. For environmental reasons, water-based ("latex") vehicles are now preferred.

The most critical properties of inks relate to their drying and surface properties; they must be able to flow properly and attach to the surface without penetrating it — the latter is especially critical when printing on a porous material such as paper.

Many inks consist of organic dyes dissolved in a water-based solvent, and are not colloidal at all. The ink used in printing newspapers employs colloidal carbon black dispersed in an oil vehicle. The pressure applied by the printing press forces the vehicle into the pores of the paper, leaving most of the pigment particles on the surface.

The inks employed in ball-point pens are gels, formulated in such a way that the ink will only flow over the ball and onto the paper when the shearing action of the ball (which rotates as it moves across the paper) "breaks" the gel into a liquid; the resulting liquid coats the ball and is transferred to the paper. As in conventional printing, the pigment particles remain on the paper surface, while the liquid is pressed into the pores and gradually evaporates.

Water and wastewater treatment

Turbidities of 5, 50, and 500 units. [WikiMedia]

Water, whether intended specifically for drinking, or wastewaters such as sewage or from industrial operations such as from pulp-and-paper manufacture (most of which are likely to end up being re-used elsewhere) usually contains colloidal matter that cannot be removed by ordinary sand filters, as evidenced by its turbidity. Even "pristine" surface waters often contain suspended soil sediments that can harbor infectious organisms and may provide them with partial protection from standard disinfection treatments.

The sulfates of aluminum (alum) and of iron(III) have long been widely employed for this purpose. Synthetic polymers tailored specifically for these applications have more recently come into use.

The usual method of removing turbidity is to add a flocculating agent (flocculant). These are most often metallic salts that can form gel-like hydroxide precipitates, often with the aid of added calcium hydroxide (quicklime) if pH of the water must be raised.

The flocculant salts neutralize the surface charges of the colloids, thus enabling them to coagulate; these are engulfed and trapped by fragments of gelatinous precipitate, which are drawn together into larger aggregates by gentle agitation until they become sufficiently large to form flocs which can be separated by settling or filtration.

Soil colloids

The four major components of soils are mineral sediments, organic matter, water, and air. The water is primarily adsorbed to the mineral and organic materials, but may also share pore spaces with air; pore spaces constitute about half the bulk volume of typical solid.

The principal colloidal components of soils are mineral sediments in the form of clays, and the humic materials in the organic matter. In addition to influencing the consistency of soil by binding water molecules, soil colloids play an essential role in storing and exchanging the mineral ions required by plants.

Most soil colloids are negatively charged, and therefor attract cations such as Ca 2 + , Mg 2 + , and K + into the outer parts of their double layers. Because these ions are loosely bound, they constitute a source from which plant roots can draw these essential nutrients. Conversely, they can serve as a sink for these same ions when they are released after the plant dies.

These are layered structures based on alumino-silicates or hydrous oxides, mostly of iron or aluminum. Each layer is built of two or three sheets of extended silica or alumina structures linked together by shared oxygen atoms. These layers generally have an overall negative charge owing to the occasional replacement of a Si 4 + ion by one of Al 3 + .

Adjacent layers are separated by a region of adsorbed cations (to neutralize the negative charges) and water molecules, and thus are held together relatively loosely. It is these interlayer regions that enable clays to work their magic by exchanging ions with both the soil water and the roots of plants.

Humic substances

The principal organic components of soil are complex substances of indeterminate structure that present –OH and –COOH groups which become increasingly dissociated as the pH increases. This allows them to bind and exchange cations in much the same way as described above.

Module 11: Solutions and Colloids

Assignment: 11.

• Olive oil (non polar)
• Ethanol or grain alcohol ( polar)
• Table Salt or NaCl (ionic)
• Sugar (Polar covalent)
• 2) Compare the processes that occur when potassium chloride, barium hydroxide and hydrobromic acid (HBr) dissolve in water. Write out the equation and make a sketch of what is happening.
• Acetic acid (Vinegar)
• Butane (C 4 H 8 )
• Ammonia (NH 3 )
• 4) At standard pressure (1 atm) and 0°C, diatomic oxygen (O 2 ) can dissolve in water at a rate of 0.70 g/ 1 L of water. At the same temperature, but at a new pressure of 0.25 atm, how much oxygen can dissolve in water?
• 120 g of NH 4 NO 3 in 300 g of water
• 25g of Br 2 in 150 g of ethanol, C 2 H 5 OH
• 25g of NaCl in dichloromethane, CH 2 Cl 2
• Methanol, Ethanol and water in a solution that is 35% methanol, 40% Ethanol and 15% water
• 875 g of H 2 SO 4 in 2kg of water
• Cyanic acid (HCN) in a solution that is 74% HCN by mass
• A 26% solution (mass) of potassium carbonate (K 2 CO 3 ) with a density of 1.09 g/cm3.
• 8) If a solution of MgBr 2 in water freezes at -1.28°C, what is the boiling point of the solution?. Assume ideal solution behavior.
• 9) Explain two major differences between colloids and solutions.
• 10) How would you prepare a 4.5 m aqueous solution of glycerol (C 3 H 8 O 3 )? Calculate freezing point for this solution. This is often used by scientists to freeze bacteria- why might this molecule be a good choice?

SOLUTIONS , SUSPENSIONS AND COLLOIDS -- SUMMARY TABLES

==>> For  more on Mixtures (Solutions, Suspensions, Emulsions, Colloids  )

In summary:

TRY THE ONLINE ASSESSMENT QUIZ ON SOLUTIONS SUSPENSIONS AND COLLOIDS

• Science of Fluids
• What are Fluids ?
• What is Pressure?
• What is Hydrostatic Pressure?
• Surface Tension and Capillary Action
• Pascals Principle
• Archimedes Principle
• What is Viscosity?
• Bernouili's Principle

Related Sites

• Mass, Volume, Density Page
• Hydrocolloids in Cooking
• What are Polysaccharides?
• What are the components of blood?
• K-12 - NGSS Activities
• EDinformatics Science Challenge

• Chemistry Articles
• Classification Of Colloids

Classification of Colloids

What is colloid.

A colloid is primarily a heterogeneous mixture in which the minute particles of one substance are dispersed in another substance, called the dispersion medium.

The minute particles here are 1 to 1000 nanometers in diameter but they still remain suspended and do not settle at the bottom of the mixture. They are visible under an optical or an electron (smaller particles) microscope.

Table of Contents

• Dispersed Phase and Dispersion Medium
• Recommended Video
• Types of Colloid
• Multimolecular Colloids
• Macromolecular Colloids
• Associated Colloids

Examples of Colloids

• Lyophilic colloids
• Lyophobic colloids

Dispersed Phase and Dispersion Medium

A colloid is a mixture in which one substance which has fine particles (dispersed phase)  mixed into another substance (dispersion medium). The particles of the colloids have a range from 1 to 1000 nm in diameter. The solution is called colloidal dispersion because the particles of solutions do not mix or settle down. They are dispersed in the solution.

The substances which are dispersed in the solution are called the dispersed phase, and the solution in which it is dispersed is called dispersion medium.

Recommended Videos

Types of Colloids

Colloids can be classified according to different properties of the dispersed phase and medium.

Firstly, based on the types of particles of the dispersed phase, colloids can be classified as:

• Multimolecular colloids
• Macromolecular colloids
• Associated colloids

1. Multimolecular Colloids

When the dissolution of smaller molecules of substance or many atoms takes place, they combine to form a species whose size is in the range of colloidal size. The species formed is known as the multimolecular colloid. For example, the Sulphur solution contains particles which have thousands of S 8 .

2. Macromolecular Colloids

In this type of colloid, the macromolecules form a solution with a suitable solvent. The size of the particles of this macromolecular solution lies in the range of colloidal particle size. Thus, this solution is also known as the macromolecular colloids. The colloids formed here are similar to that of the actual solution in many respects and are very stable. Example: Starch, proteins, enzymes, and cellulose are the naturally occurring macromolecular colloids whereas polyethene, synthetic rubber, etc. are the synthetic macromolecules.

3. Associated Colloids

Some substances act as a strong electrolyte when they are in low concentrations, but they react as colloidal sols when they are in high concentration. In higher concentration, particles aggregate showing colloidal behaviour. These aggregated particles are known as the micelles. They are also known as the associated colloids. The formation of the micelles occurs above a particular temperature called the Kraft temperature (T k ) and also above a specific concentration called the critical micelle concentration. These colloids can be reverted by diluting it. Examples of some associated colloids are soaps and synthetic detergents .

Based on the physical state of the dispersion medium and of the dispersed phase, colloids can be classified into:

• Solid Aerosol

The following table describes the types along with examples:

Based on the nature of the interaction between the dispersion medium and the dispersed phase, colloids can be classified into lyophilic and lyophobic.

If the dispersed phase has an affinity for the dispersion medium, the colloid is called a lyophilic colloid. The words lyo and philic mean ‘ liquid ’ and ‘ loving ’ respectively. Thus, even if the dispersed phase is separated from the dispersion medium, they can readily be reconstituted by simply mixing them. Moreover, they are difficult to coagulate due to their stable nature. They are also known as intrinsic colloids . Examples are starch, rubber, protein, etc.

If the dispersed phase has little or no affinity for the dispersion medium, the colloid is called a lyophobic colloid. The words lyo and phobic mean ‘ liquid ’ and ‘ fearing ’ respectively. Hence, they are liquid-hating. They are difficult to prepare because the dispersed phase does not readily form a colloid with the dispersion medium; they require some special methods. They are unstable and require stabilising agents for their preservation. They are also known as extrinsic colloids . Examples are sols of metals like silver and gold, sols of metallic hydroxides, etc.

After reading this article, we now know about the different types of colloids when classified from types of particles used in the dispersed phase. Learn more about the colloids, register with BYJU’S & download BYJU’S – The Learning App.

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8th-graders given hitler-themed assignment to rate nazi monster as a ‘solution seeker,’ ‘ethical decision-maker’.

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An Adolf Hitler-themed question-and-answer assignment given to students at a private school in Atlanta has sparked outrage among parents over its suspected antisemitic nature.

Eighth-grade students at the Mount Vernon School in Atlanta were given a series of questions asking them to rate some of the characteristics of Adolf Hitler — the dictator of Nazi Germany from 1933 to 1945, whose antisemitic ideology fueled the Holocaust — as a leader, according to Fox 5 Atlanta . 

One question posed to students asked, “According to the Mount Vernon Mindset rubric, how would you rate Adolf Hitler as a ‘solution seeker’?” 

A second question asked how students would “rate Adolf Hitler as an ethical decision-maker?”

For both questions, the students were given the option of selecting “Lacks Evidence,” “Approaching Expectations,” “Meets Expectations” or “Exceeds Expectations” to describe the ruthless dictator. 

The bizarre questions ignited outrage among parents — many of whom were concerned the queries were antisemitic by nature, according to the outlet. 

Students at the private school also had issues with the questions, with one telling the outlet the assignment was “troubling” and could be seen as glorifying the warmongering totalitarian leader. 

“Obviously, that looks horrible in the current context,” another student told the outlet. “Knowing Mount Vernon, we do things a little odd around here.”

The student added that the school is known to “try to think outside the box” but shared that “oftentimes that doesn’t work.”

Several former students told Fox 5 that those questions weren’t given to them during eighth grade.

While many parents and students were shaken over the assignment, one student believes the school attempted to pose a historically provocative question that required students to use their critical thinking skills. 

“I can definitely see why they’d be upset, but overall, I think it’s important to look at both sides of the coin in every situation, and I think it’s important to be able to compare and contrast everything that’s happened in our world history, whether it’s been good or bad,” said the student.

Upon learning the phrasing of the questions in the assignment, Mount Vernon officials said they had removed it from the school’s curriculum. 

The principal of Mount Vernon, Kristy Lundstrom, wrote in a statement that the assignment was “an exploration of World War II designed to boost student knowledge of factual events and understand the manipulation of fear leveraged by Adolf Hitler in connection to the Treaty of Versailles.” 

“Immediately following this incident, I met with the School’s Chief of Inclusion, Diversity, Equality, and Action, Head of Middle School, and a concerned Rabbi and friend of the School who shared the perspective of some of our families and supported us in a thorough review of the assignment and community impact.”

“Adolf Hitler and the events of the time period are difficult and traumatic to discuss.”

The private school, about 16 miles outside downtown Atlanta, is a “co-educational day school for more than 1200 students in Preschool through Grade 12,” according to the institution’s  website . 

“We are a school of inquiry, innovation, and impact. Grounded in Christian values, we prepare all students to be college ready, globally competitive, and engaged citizen leaders,” its mission statement reads.

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11.E: Solutions and Colloids (Exercises)

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11.2: The Dissolution Process

How do solutions differ from compounds? From other mixtures?

A solution can vary in composition, while a compound cannot vary in composition. Solutions are homogeneous at the molecular level, while other mixtures are heterogeneous.

Which of the principal characteristics of solutions can we see in the solutions of \(\ce{K2Cr2O7}\) shown in

Figure: When potassium dichromate (\(\ce{K2Cr2O7}\)) is mixed with water, it forms a homogeneous orange solution. (credit: modification of work by Mark Ott)

The solutions are the same throughout (the color is constant throughout), and the composition of a solution of K 2 Cr 2 O 7 in water can vary.

When KNO 3 is dissolved in water, the resulting solution is significantly colder than the water was originally.

• Is the dissolution of KNO 3 an endothermic or an exothermic process?
• What conclusions can you draw about the intermolecular attractions involved in the process?
• Is the resulting solution an ideal solution?

(a) The process is endothermic as the solution is consuming heat. (b) Attraction between the K + and \(\ce{NO3-}\) ions is stronger than between the ions and water molecules (the ion-ion interactions have a lower, more negative energy). Therefore, the dissolution process increases the energy of the molecular interactions, and it consumes the thermal energy of the solution to make up for the difference. (c) No, an ideal solution is formed with no appreciable heat release or consumption.

Give an example of each of the following types of solutions:

• a gas in a liquid
• a gas in a gas
• a solid in a solid

(a) CO 2 in water; (b) O 2 in N 2 (air); (c) bronze (solution of tin or other metals in copper)

Indicate the most important types of intermolecular attractions in each of the following solutions:

• The solution in Figure .
• NO( l ) in CO( l )
• Cl 2 ( g ) in Br 2 ( l )
• HCl( aq ) in benzene C 6 H 6 ( l )
• Methanol CH 3 OH( l ) in H 2 O( l )

(a) ion-dipole forces; (b) dipole-dipole forces; (c) dispersion forces; (d) dispersion forces; (e) hydrogen bonding

Predict whether each of the following substances would be more soluble in water (polar solvent) or in a hydrocarbon such as heptane (C 7 H 16 , nonpolar solvent):

• vegetable oil (nonpolar)
• isopropyl alcohol (polar)
• potassium bromide (ionic)

(a) heptane; (b) water; (c) water

Heat is released when some solutions form; heat is absorbed when other solutions form. Provide a molecular explanation for the difference between these two types of spontaneous processes.

Heat is released when the total intermolecular forces (IMFs) between the solute and solvent molecules are stronger than the total IMFs in the pure solute and in the pure solvent: Breaking weaker IMFs and forming stronger IMFs releases heat. Heat is absorbed when the total IMFs in the solution are weaker than the total of those in the pure solute and in the pure solvent: Breaking stronger IMFs and forming weaker IMFs absorbs heat.

Solutions of hydrogen in palladium may be formed by exposing Pd metal to H 2 gas. The concentration of hydrogen in the palladium depends on the pressure of H 2 gas applied, but in a more complex fashion than can be described by Henry’s law. Under certain conditions, 0.94 g of hydrogen gas is dissolved in 215 g of palladium metal.

• Determine the molarity of this solution (solution density = 1.8 g/cm 3 ).
• Determine the molality of this solution (solution density = 1.8 g/cm 3 ).
• Determine the percent by mass of hydrogen atoms in this solution (solution density = 1.8 g/cm 3 ).

http://cnx.org/contents/ mH6aqegx @2/The-Dissolution-Process

11.3: Electrolytes

Explain why the ions Na + and Cl − are strongly solvated in water but not in hexane, a solvent composed of nonpolar molecules.

Crystals of NaCl dissolve in water, a polar liquid with a very large dipole moment, and the individual ions become strongly solvated. Hexane is a nonpolar liquid with a dipole moment of zero and, therefore, does not significantly interact with the ions of the NaCl crystals.

Explain why solutions of HBr in benzene (a nonpolar solvent) are nonconductive, while solutions in water (a polar solvent) are conductive.

HBr is an acid and so its molecules react with water molecules to form H 3 O + and Br − ions that provide conductivity to the solution. Though HBr is soluble in benzene, it does not react chemically but remains dissolved as neutral HBr molecules. With no ions present in the benzene solution, it is electrically nonconductive.

Consider the solutions presented:

(a) Which of the following sketches best represents the ions in a solution of Fe(NO 3 ) 3 ( aq )?

(b) Write a balanced chemical equation showing the products of the dissolution of Fe(NO 3 ) 3 .

(a) Fe(NO 3 ) 3 is a strong electrolyte, thus it should completely dissociate into Fe 3+ and \(\ce{(NO3- )}\) ions. Therefore, (z) best represents the solution. (b) \(\ce{Fe(NO3)3}(s)⟶\ce{Fe^3+}(aq)+\ce{3NO3- }(aq)\)

Compare the processes that occur when methanol (CH 3 OH), hydrogen chloride (HCl), and sodium hydroxide (NaOH) dissolve in water. Write equations and prepare sketches showing the form in which each of these compounds is present in its respective solution.

Methanol, \(CH_3OH\), dissolves in water in all proportions, interacting via hydrogen bonding.

\[CH_3OH_{(l)}+H_2O_{(l)}⟶CH_3OH_{(aq)}\]

Hydrogen chloride, HCl, dissolves in and reacts with water to yield hydronium cations and chloride anions that are solvated by strong ion-dipole interactions.

Hydrogen chloride:

\[HCl{(g)}+H_2O_{(l)} \rightarrow H_3O^+_{(aq)}+Cl^−_{(aq)}\]

Sodium hydroxide, NaOH, dissolves in water and dissociates to yield sodium cations and hydroxide anions that are strongly solvated by ion-dipole interactions and hydrogen bonding, respectively.

Sodium hydroxide:

\[NaOH_{(s)} \rightarrow Na^+_{(aq)} + OH^−_{(aq)}\]

What is the expected electrical conductivity of the following solutions?

• C 6 H 12 O 6 ( aq ) (glucose)

(a) high conductivity (solute is an ionic compound that will dissociate when dissolved); (b) high conductivity (solute is a strong acid and will ionize completely when dissolved); (c) nonconductive (solute is a covalent compound, neither acid nor base, unreactive towards water); (d) low conductivity (solute is a weak base and will partially ionize when dissolved)

Why are most solid ionic compounds electrically nonconductive, whereas aqueous solutions of ionic compounds are good conductors? Would you expect a liquid (molten) ionic compound to be electrically conductive or nonconductive? Explain.

A medium must contain freely mobile, charged entities to be electrically conductive. The ions present in a typical ionic solid are immobilized in a crystalline lattice and so the solid is not able to support an electrical current. When the ions are mobilized, either by melting the solid or dissolving it in water to dissociate the ions, current may flow and these forms of the ionic compound are conductive.

Indicate the most important type of intermolecular attraction responsible for solvation in each of the following solutions:

• the solutions in Figure
• methanol, CH 3 OH, dissolved in ethanol, C 2 H 5 OH
• methane, CH 4 , dissolved in benzene, C 6 H 6
• the polar halocarbon CF 2 Cl 2 dissolved in the polar halocarbon CF 2 ClCFCl 2
• O 2 ( l ) in N 2 ( l )

(a) ion-dipole; (b) hydrogen bonds; (c) dispersion forces; (d) dipole-dipole attractions; (e) dispersion forces

11.4: Solubility

Suppose you are presented with a clear solution of sodium thiosulfate, Na 2 S 2 O 3 . How could you determine whether the solution is unsaturated, saturated, or supersaturated?

Add a small crystal of \(Na_2S_2O_3\). It will dissolve in an unsaturated solution, remain apparently unchanged in a saturated solution, or initiate precipitation in a supersaturated solution.

Supersaturated solutions of most solids in water are prepared by cooling saturated solutions. Supersaturated solutions of most gases in water are prepared by heating saturated solutions. Explain the reasons for the difference in the two procedures.

The solubility of solids usually decreases upon cooling a solution, while the solubility of gases usually decreases upon heating.

Suggest an explanation for the observations that ethanol, C 2 H 5 OH, is completely miscible with water and that ethanethiol, C 2 H 5 SH, is soluble only to the extent of 1.5 g per 100 mL of water.

The hydrogen bonds between water and C 2 H 5 OH are much stronger than the intermolecular attractions between water and C 2 H 5 SH.

Calculate the percent by mass of KBr in a saturated solution of KBr in water at 10 °C using the following figure for useful data, and report the computed percentage to one significant digit.

At 10 °C, the solubility of KBr in water is approximately 60 g per 100 g of water.

\[\%\; KBr =\dfrac{60\; g\; KBr}{(60+100)\;g\; solution} = 40\%\]

Which of the following gases is expected to be most soluble in water? Explain your reasoning.

(c) CHCl 3 is expected to be most soluble in water. Of the three gases, only this one is polar and thus capable of experiencing relatively strong dipole-dipole attraction to water molecules.

At 0 °C and 1.00 atm, as much as 0.70 g of O 2 can dissolve in 1 L of water. At 0 °C and 4.00 atm, how many grams of O 2 dissolve in 1 L of water?

This problem requires the application of Henry’s law. The governing equation is \(C_g = kP_g\).

\[k=\dfrac{C_g}{P_g}=\dfrac{0.70\;g}{1.00\; atm} =0.70\;g\; atm^{−1}\]

Under the new conditions, \(C_g=0.70\;g\;atm^{−1} \times 4.00\; atm = 2.80\; g\).

Refer to following figure for the following three questions:

• How did the concentration of dissolved CO 2 in the beverage change when the bottle was opened?
• What caused this change?
• Is the beverage unsaturated, saturated, or supersaturated with CO 2 ?

(a) It decreased as some of the CO 2 gas left the solution (evidenced by effervescence). (b) Opening the bottle released the high-pressure CO 2 gas above the beverage. The reduced CO 2 gas pressure, per Henry’s law, lowers the solubility for CO 2 . (c) The dissolved CO 2 concentration will continue to slowly decrease until equilibrium is reestablished between the beverage and the very low CO 2 gas pressure in the opened bottle. Immediately after opening, the beverage, therefore, contains dissolved CO 2 at a concentration greater than its solubility, a nonequilibrium condition, and is said to be supersaturated.

The Henry’s law constant for CO 2 is \(3.4 \times 10^{−2}\; M/atm\) at 25 °C. What pressure of carbon dioxide is needed to maintain a CO 2 concentration of 0.10 M in a can of lemon-lime soda?

\[P_g=\dfrac{C_g}{k}=\dfrac{0.10\; M}{3.4 \times 10^{−2}\;M/atm} =2.9\; atm\]

The Henry’s law constant for O 2 is \(1.3\times 10^{−3}\; M/atm\) at 25 °C. What mass of oxygen would be dissolved in a 40-L aquarium at 25 °C, assuming an atmospheric pressure of 1.00 atm, and that the partial pressure of O 2 is 0.21 atm?

Start with Henry's law

\[C_g=kP_g\]

and apply it to \(O_2\)

\[C(O_2)=(1.3 \times 10^{−3}\; M/atm) (0.21\;atm)=2.7 \times 10^{−4}\;mol/L\]

The total amount is \((2.7 \times 10^{−4}\; mol/L)(40\;L=1.08 \times 10^{−2} \;mol\]

The mass of oxygen is \((1.08 \times 10^{−2}\; mol)(32.0\; g/mol)=0.346\;g\)

or, using two significant figures, \(0.35\; g\).

How many liters of HCl gas, measured at 30.0 °C and 745 torr, are required to prepare 1.25 L of a 3.20- M solution of hydrochloric acid?

First, calculate the moles of HCl needed. Then use the ideal gas law to find the volume required.

M = mol L−1

3.20M=xmol1.25L

x = 4.00 mol HCl

Before using the ideal gas law, change pressure to atmospheres and convert temperature from °C to kelvin .

\[1\;atmx = 760torr745torr

x = 0.9803 atm

V= nRTP =(4.000molHCl)(0.08206LatmK−1mol−1)(303.15K)0.9803atm=102 L HCl

more http://cnx.org/contents/ 2488fW6W @2/Solubility

11.5: Colligative Properties

Which is/are part of the macroscopic domain of solutions and which is/are part of the microscopic domain: boiling point elevation, Henry’s law, hydrogen bond, ion-dipole attraction, molarity, nonelectrolyte, nonstoichiometric compound, osmosis, solvated ion?

What is the microscopic explanation for the macroscopic behavior illustrated in [link] ?

The strength of the bonds between like molecules is stronger than the strength between unlike molecules. Therefore, some regions will exist in which the water molecules will exclude oil molecules and other regions will exist in which oil molecules will exclude water molecules, forming a heterogeneous region.

Sketch a qualitative graph of the pressure versus time for water vapor above a sample of pure water and a sugar solution, as the liquids evaporate to half their original volume.

A solution of potassium nitrate, an electrolyte, and a solution of glycerin (C 3 H 5 (OH) 3 ), a nonelectrolyte, both boil at 100.3 °C. What other physical properties of the two solutions are identical?

Both form homogeneous solutions; their boiling point elevations are the same, as are their lowering of vapor pressures. Osmotic pressure and the lowering of the freezing point are also the same for both solutions.

What are the mole fractions of H 3 PO 4 and water in a solution of 14.5 g of H 3 PO 4 in 125 g of water?

What are the mole fractions of HNO 3 and water in a concentrated solution of nitric acid (68.0% HNO 3 by mass)?

• Find number of moles of HNO 3 and H 2 O in 100 g of the solution. Find the mole fractions for the components.
• The mole fraction of HNO 3 is 0.378. The mole fraction of H 2 O is 0.622.

Calculate the mole fraction of each solute and solvent:

• 583 g of H 2 SO 4 in 1.50 kg of water—the acid solution used in an automobile battery
• 0.86 g of NaCl in 1.00 × 10 2 g of water—a solution of sodium chloride for intravenous injection
• 46.85 g of codeine, C 18 H 21 NO 3 , in 125.5 g of ethanol, C 2 H 5 OH
• 25 g of I 2 in 125 g of ethanol, C 2 H 5 OH

a. \(583\:g\:\ce{H2SO4}\times\dfrac{1\:mole\:\ce{H2SO4}}{98.08\:g\:\ce{H2SO4}}=5.94\:mole\:\ce{H2SO4}\) \(\:\:\:\:\:\:\:\:\:\:\:\:\:\:\:\:\:\) \(1.50\:kg\:\ce{H2O}\times\dfrac{1000\:g}{1\:kg}\times\dfrac{1\:mole\:\ce{H2O}}{18.02\:g\:\ce{H2O}}=83.2\:moles\:\ce{H2O}\)

• 0.710 kg of sodium carbonate (washing soda), Na 2 CO 3 , in 10.0 kg of water—a saturated solution at 0 °C
• 125 g of NH 4 NO 3 in 275 g of water—a mixture used to make an instant ice pack
• 25 g of Cl 2 in 125 g of dichloromethane, CH 2 Cl 2
• 0.372 g of histamine, C 5 H 9 N, in 125 g of chloroform, CHCl 3
• \(X_\mathrm{Na_2CO_3}=0.0119\); \(X_\mathrm{H_2O}=0.988\);
• \(X_\mathrm{NH_4NO_3}=0.9927\); \(X_\mathrm{H_2O}=0.907\);
• \(X_\mathrm{Cl_2}=0.192\); \(X_\mathrm{CH_2CI_2}=0.808\);
• \(X_\mathrm{C_5H_9N}=0.00426\); \(X_\mathrm{CHCl_3}=0.997\)

Calculate the mole fractions of methanol, CH 3 OH; ethanol, C 2 H 5 OH; and water in a solution that is 40% methanol, 40% ethanol, and 20% water by mass. (Assume the data are good to two significant figures.)

What is the difference between a 1 M solution and a 1 m solution?

In a 1 M solution, the mole is contained in exactly 1 L of solution. In a 1 m solution, the mole is contained in exactly 1 kg of solvent.

What is the molality of phosphoric acid, H 3 PO 4 , in a solution of 14.5 g of H 3 PO 4 in 125 g of water?

What is the molality of nitric acid in a concentrated solution of nitric acid (68.0% HNO 3 by mass)?

(a) Determine the molar mass of HNO 3 . Determine the number of moles of acid in the solution. From the number of moles and the mass of solvent, determine the molality. (b) 33.7 m

Calculate the molality of each of the following solutions:

• 0.710 kg of sodium carbonate (washing soda), Na 2 CO 3 , in 10.0 kg of water—a saturated solution at 0°C

(a) 6.70 × 10 −1 m ; (b) 5.67 m ; (c) 2.8 m ; (d) 0.0358 m

The concentration of glucose, C 6 H 12 O 6 , in normal spinal fluid is \(\mathrm{\dfrac{75\:mg}{100\:g}}\). What is the molality of the solution?

A 13.0% solution of K 2 CO 3 by mass has a density of 1.09 g/cm 3 . Calculate the molality of the solution.

• Why does 1 mol of sodium chloride depress the freezing point of 1 kg of water almost twice as much as 1 mol of glycerin?
• What is the boiling point of a solution of 115.0 g of sucrose, C 12 H 22 O 11 , in 350.0 g of water?
• Determine the molar mass of sucrose; determine the number of moles of sucrose in the solution; convert the mass of solvent to units of kilograms; from the number of moles and the mass of solvent, determine the molality; determine the difference between the boiling point of water and the boiling point of the solution; determine the new boiling point.
• 100.5 °C

What is the boiling point of a solution of 9.04 g of I 2 in 75.5 g of benzene, assuming the I 2 is nonvolatile?

What is the freezing temperature of a solution of 115.0 g of sucrose, C 12 H 22 O 11 , in 350.0 g of water, which freezes at 0.0 °C when pure?

(a) Determine the molar mass of sucrose; determine the number of moles of sucrose in the solution; convert the mass of solvent to units of kilograms; from the number of moles and the mass of solvent, determine the molality; determine the difference between the freezing temperature of water and the freezing temperature of the solution; determine the new freezing temperature. (b) −1.8 °C

What is the freezing point of a solution of 9.04 g of I 2 in 75.5 g of benzene?

What is the osmotic pressure of an aqueous solution of 1.64 g of Ca(NO 3 ) 2 in water at 25 °C? The volume of the solution is 275 mL.

(a) Determine the molar mass of Ca(NO 3 ) 2 ; determine the number of moles of Ca(NO 3 ) 2 in the solution; determine the number of moles of ions in the solution; determine the molarity of ions, then the osmotic pressure. (b) 2.67 atm

What is osmotic pressure of a solution of bovine insulin (molar mass, 5700 g mol −1 ) at 18 °C if 100.0 mL of the solution contains 0.103 g of the insulin?

What is the molar mass of a solution of 5.00 g of a compound in 25.00 g of carbon tetrachloride (bp 76.8 °C; K b = 5.02 °C/ m ) that boils at 81.5 °C at 1 atm?

(a) Determine the molal concentration from the change in boiling point and K b ; determine the moles of solute in the solution from the molal concentration and mass of solvent; determine the molar mass from the number of moles and the mass of solute. (b) 2.1 × 10 2 g mol −1

A sample of an organic compound (a nonelectrolyte) weighing 1.35 g lowered the freezing point of 10.0 g of benzene by 3.66 °C. Calculate the molar mass of the compound.

A 1.0 m solution of HCl in benzene has a freezing point of 0.4 °C. Is HCl an electrolyte in benzene? Explain.

No. Pure benzene freezes at 5.5 °C, and so the observed freezing point of this solution is depressed by Δ T f = 5.5 − 0.4 = 5.1 °C. The value computed, assuming no ionization of HCl, is Δ T f = (1.0 m)(5.14 °C/ m ) = 5.1 °C. Agreement of these values supports the assumption that HCl is not ionized.

A solution contains 5.00 g of urea, CO(NH 2 ) 2 , a nonvolatile compound, dissolved in 0.100 kg of water. If the vapor pressure of pure water at 25 °C is 23.7 torr, what is the vapor pressure of the solution?

A 12.0-g sample of a nonelectrolyte is dissolved in 80.0 g of water. The solution freezes at −1.94 °C. Calculate the molar mass of the substance.

144 g mol −1

Arrange the following solutions in order by their decreasing freezing points: 0.1 m Na 3 PO 4 , 0.1 m C 2 H 5 OH, 0.01 m CO 2 , 0.15 m NaCl, and 0.2 m CaCl 2 .

Calculate the boiling point elevation of 0.100 kg of water containing 0.010 mol of NaCl, 0.020 mol of Na 2 SO 4 , and 0.030 mol of MgCl 2 , assuming complete dissociation of these electrolytes.

0.870 °C

How could you prepare a 3.08 m aqueous solution of glycerin, C 3 H 8 O 3 ? What is the freezing point of this solution?

A sample of sulfur weighing 0.210 g was dissolved in 17.8 g of carbon disulfide, CS 2 ( K b = 2.43 °C/ m ). If the boiling point elevation was 0.107 °C, what is the formula of a sulfur molecule in carbon disulfide?

In a significant experiment performed many years ago, 5.6977 g of cadmium iodide in 44.69 g of water raised the boiling point 0.181 °C. What does this suggest about the nature of a solution of CdI 2 ?

Lysozyme is an enzyme that cleaves cell walls. A 0.100-L sample of a solution of lysozyme that contains 0.0750 g of the enzyme exhibits an osmotic pressure of 1.32 × 10 −3 atm at 25 °C. What is the molar mass of lysozyme?

1.39 × 10 4 g mol −1

The osmotic pressure of a solution containing 7.0 g of insulin per liter is 23 torr at 25 °C. What is the molar mass of insulin?

The osmotic pressure of human blood is 7.6 atm at 37 °C. What mass of glucose, C 6 H 12 O 6 , is required to make 1.00 L of aqueous solution for intravenous feeding if the solution must have the same osmotic pressure as blood at body temperature, 37 °C?

What is the freezing point of a solution of dibromobenzene, C 6 H 4 Br 2 , in 0.250 kg of benzene, if the solution boils at 83.5 °C?

What is the boiling point of a solution of NaCl in water if the solution freezes at −0.93 °C?

100.26 °C

The sugar fructose contains 40.0% C, 6.7% H, and 53.3% O by mass. A solution of 11.7 g of fructose in 325 g of ethanol has a boiling point of 78.59 °C. The boiling point of ethanol is 78.35 °C, and K b for ethanol is 1.20 °C/ m . What is the molecular formula of fructose?

The vapor pressure of methanol, CH 3 OH, is 94 torr at 20 °C. The vapor pressure of ethanol, C 2 H 5 OH, is 44 torr at the same temperature.

• Calculate the mole fraction of methanol and of ethanol in a solution of 50.0 g of methanol and 50.0 g of ethanol.
• Ethanol and methanol form a solution that behaves like an ideal solution. Calculate the vapor pressure of methanol and of ethanol above the solution at 20 °C.
• Calculate the mole fraction of methanol and of ethanol in the vapor above the solution.

(a) \(X_\mathrm{CH_3OH}=0.590\); \(X_\mathrm{C_2H_5OH}=0.410\); (b) Vapor pressures are: CH 3 OH: 55 torr; C 2 H 5 OH: 18 torr; (c) CH 3 OH: 0.75; C 2 H 5 OH: 0.25

The triple point of air-free water is defined as 273.15 K. Why is it important that the water be free of air?

Meat can be classified as fresh (not frozen) even though it is stored at −1 °C. Why wouldn’t meat freeze at this temperature?

The ions and compounds present in the water in the beef lower the freezing point of the beef below −1 °C.

An organic compound has a composition of 93.46% C and 6.54% H by mass. A solution of 0.090 g of this compound in 1.10 g of camphor melts at 158.4 °C. The melting point of pure camphor is 178.4 °C. K f for camphor is 37.7 °C/ m . What is the molecular formula of the solute? Show your calculations.

A sample of HgCl 2 weighing 9.41 g is dissolved in 32.75 g of ethanol, C 2 H 5 OH ( K b = 1.20 °C/ m ). The boiling point elevation of the solution is 1.27 °C. Is HgCl 2 an electrolyte in ethanol? Show your calculations.

\(\mathrm{Δbp}=K_\ce{b}m=(1.20\:°\ce C/m)\mathrm{\left(\dfrac{9.41\:g×\dfrac{1\:mol\: HgCl_2}{271.496\:g}}{0.03275\:kg}\right)=1.27\:°\ce C}\)

The observed change equals the theoretical change; therefore, no dissociation occurs.

A salt is known to be an alkali metal fluoride. A quick approximate determination of freezing point indicates that 4 g of the salt dissolved in 100 g of water produces a solution that freezes at about −1.4 °C. What is the formula of the salt? Show your calculations.

11.6: Colloids

Identify the dispersed phase and the dispersion medium in each of the following colloidal systems: starch dispersion, smoke, fog, pearl, whipped cream, floating soap, jelly, milk, and ruby.

Distinguish between dispersion methods and condensation methods for preparing colloidal systems.

Dispersion methods use a grinding device or some other means to bring about the subdivision of larger particles. Condensation methods bring smaller units together to form a larger unit. For example, water molecules in the vapor state come together to form very small droplets that we see as clouds.

How do colloids differ from solutions with regard to dispersed particle size and homogeneity?

Colloidal dispersions consist of particles that are much bigger than the solutes of typical solutions. Colloidal particles are either very large molecules or aggregates of smaller species that usually are big enough to scatter light. Colloids are homogeneous on a macroscopic (visual) scale, while solutions are homogeneous on a microscopic (molecular) scale.

Explain the cleansing action of soap.

Soap molecules have both a hydrophobic and a hydrophilic end. The charged (hydrophilic) end, which is usually associated with an alkali metal ion, ensures water solubility The hydrophobic end permits attraction to oil, grease, and other similar nonpolar substances that normally do not dissolve in water but are pulled into the solution by the soap molecules.

How can it be demonstrated that colloidal particles are electrically charged?

If they are placed in an electrolytic cell, dispersed particles will move toward the electrode that carries a charge opposite to their own charge. At this electrode, the charged particles will be neutralized and will coagulate as a precipitate.

Contributors and Attributions

Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors.  Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/[email protected] ).