Difference between Colloid and Solution

Back to elementary, Homogeneous and Heterogeneous are mentioned... both of these are Solution and Colloid. If you can't figure what which is Heterogeneous or Homogeneous for Solution, Homogeneous is Solution the remaining is for Colloid.

Let's define again what is Colloid and Solution.

Colloid — mixture of solid particles that are suspended within the liquid.
Solution — mixture of solid particles that are completely dissolved within the liquid.

The only difference between Colloid and Solution is the properties of solid particles within the liquid.

Definition of Solutions

Definition of Solutions

• In chemistry, a solution is a homogeneous mixture composed of only one phase. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent. Thesolvent does the dissolving. The solution more or less takes on the characteristics of the solvent including its phase, and the solvent is commonly the major fraction of the mixture. The concentration of a solute in a solution is a measure of how much of that solute is dissolved in the solvent.

• Characteristics

• A solution is a homogeneous mixture.
• A solution is a single phase system.
• The particles of solute in solution cannot be seen by naked eye.
• The solution does not allow beam of light to scatter.
• A solution is stable.
• The solute from the solution cannot be separated by filtration (or mechanically).

Types

• Homogeneous means that the components of the mixture form a single phase. The properties of the mixture (such as concentration, temperature, and density) can be uniformly distributed through the volume but only in absence of diffusion phenomena or after their completion. Usually, the substance present in the greatest amount is considered the solvent. Solvents can be gases, liquids or solids. One or more components present in the solution other than the solvent are called solutes. The solution has the same physical state as the solvent.

Gas

• If the solvent is a gas, only gases are dissolved under a given set of conditions. An example of a gaseous solution is air (oxygen and other gases dissolved in nitrogen). Since interactions between molecules play almost no role, dilute gases form rather trivial solutions. In part of the literature, they are not even classified as solutions, but addressed as mixtures.
  
  

  Liquid

  If the solvent is a liquid, then gases, liquids, and solids can be dissolved. Here are some examples:
  • Gas in liquid:
    • Oxygen in water
    • Carbon dioxide in water – a less simple example, because the solution is accompanied by a chemical reaction (formation of ions). Note also that the visible bubbles in carbonated water are not the dissolved gas, but only an effervescence of carbon dioxide that has come out of solution; the dissolved gas itself is not visible since it is dissolved on a molecular level.
  • Liquid in liquid:
    • The mixing of two or more substances of the same chemistry but different concentrations to form a constant. (Homogenization of solutions)
    • Alcoholic beverages are basically solutions of ethanol in water.
  • Solid in liquid:
    • Sucrose (table sugar) in water
    • Sodium chloride (table salt) or any other salt in water, which forms an electrolyte: When dissolving, salt dissociates into ions.
  Counterexamples are provided by liquid mixtures that are not homogeneous: colloids, suspensions, emulsions are not considered solutions.
  Body fluids are examples for complex liquid solutions, containing many solutes. Many of these are electrolytes, since they contain solute ions, such as potassium. Furthermore, they contain solute molecules like sugar and urea. Oxygen and carbon dioxide are also essential components of blood chemistry, where significant changes in their concentrations may be a sign of severe illness or injury.
  
  

  Solid

  If the solvent is a solid, then gases, liquids and solids can be dissolved.
• Gas in solids:
  • Hydrogen dissolves rather well in metals, especially in palladium; this is studied as a means of hydrogen storage.
• Liquid in solid:
  • Mercury in gold, forming an amalgam
  • Hexane in paraffin wax
• Solid in solid:
• Steel, basically a solution of carbon atoms in a crystalline matrix of iron atoms.
• Alloys like bronze and many others.
• Polymers containing plasticizers.

Solubility

Main article: Solubility

Main article: Solvation

• The ability of one compound to dissolve in another compound is called solubility. When a liquid can completely dissolve in another liquid the two liquids are miscible. Two substances that can never mix to form a solution are called immiscible.
• All solutions have a positive entropy of mixing. The interactions between different molecules or ions may be energetically favored or not. If interactions are unfavorable, then the free energy decreases with increasing solute concentration. At some point the energy loss outweighs the entropy gain, and no more solute particles can be dissolved; the solution is said to be saturated. However, the point at which a solution can become saturated can change significantly with different environmental factors, such as temperature, pressure, and contamination. For some solute-solvent combinations a supersaturated solution can be prepared by raising the solubility (for example by increasing the temperature) to dissolve more solute, and then lowering it (for example by cooling).

Usually, the greater the temperature of the solvent, the more of a given solid solute it can dissolve. However, most gases and some compounds exhibit solubilities that decrease with increased temperature. Such behavior is a result of an exothermic enthalpy of solution. Some surfactants exhibit this behaviour. The solubility of liquids in liquids is generally less temperature-sensitive than that of solids or gases.

Properties

• The physical properties of compounds such as melting point and boiling point change when other compounds are added. Together they are called colligative properties. There are several ways to quantify the amount of one compound dissolved in the other compounds collectively called concentration. Examples include molarity, mole fraction, and parts per million (PPM).
• The properties of ideal solutions can be calculated by the linear combination of the properties of its components. If both solute and solvent exist in equal quantities (such as in a 50% ethanol, 50% water solution), the concepts of "solute" and "solvent" become less relevant, but the substance that is more often used as a solvent is normally designated as the solvent (in this example, water).

Liquid

• In principle, all types of liquids can behave as solvents: liquid noble gases, molten metals, molten salts, molten covalent networks, and molecular liquids. In the practice of chemistry and biochemistry, most solvents are molecular liquids. They can be classified into polar and non-polar, according to whether their molecules possess a permanent electric dipole moment. Another distinction is whether their molecules can form hydrogen bonds (protic and aprotic solvents). Water, the most commonly used solvent, is both polar and sustains hydrogen bonds.

Water is a good solvent because the molecules are polar and capable of forming hydrogen bonds(1).

• Salts dissolve in polar solvents, forming positive and negative ions that are attracted to the negative and positive ends of the solvent molecule, respectively. If the solvent is water, hydration occurs when the charged solute ions become surrounded by water molecules. A standard example is aqueous saltwater. Such solutions are called electrolytes.
• For non-ionic solutes, the general rule is: like dissolves like.
• Polar solutes dissolve in polar solvents, forming polar bonds or hydrogen bonds. As an example, all alcoholic beverages are aqueous solutions of ethanol. On the other hand, non-polar solutes dissolve better in non-polar solvents. Examples are hydrocarbons such as oil and grease that easily mix with each other, while being incompatible with water.
• An example for the immiscibility of oil and water is a leak of petroleum from a damaged tanker, that does not dissolve in the ocean water but rather floats on the surface.

IMAGE OF SOLUTION

Making a saline water solution by dissolving table salt (NaCl) in water. The salt is the solute and the water the solvent.

Questions

1. What is a Colloid?

Answer

• Colloids are solutions that, though homogenous, have particles of one substance suspended within another substance . Some common colloids include milk, fog, and jello.

Definition of Colloid

Definition of Colloids

• Before we start to explore various examples of colloids, let us do a quick recap of basic Definition of Colloids. A colloid is a heterogeneous system in which one substance is dispersed (called dispersed phase) as very fine particles in another substance called dispersion medium. The size of the dispersed molecule is larger than a simple molecule (having diameter between 1 to 1000 nm) but small enough to remain suspended. So colloid is an intermediate state between suspensions and solutions.
• A colloid is a substance microscopically dispersed throughout another substance.^[1]
  The dispersed-phase particles have a diameter of between approximately 1 and 1000nanometers.^[2] Such particles are normally invisible in an optical microscope, though their presence can be confirmed with the use of an ultramicroscope or an electron microscope. Homogeneous mixtures with a dispersed phase in this size range may be called colloidal aerosols, colloidal emulsions, colloidal foams, colloidal dispersions, orhydrosols. The dispersed-phase particles or droplets are affected largely by the surface chemistry present in the colloid.
  Some colloids are translucent because of the Tyndall effect, which is the scattering of light by particles in the colloid. Other colloids may be opaque or have a slight color.
  Colloidal solutions (also called colloidal suspensions) are the subject of interface and colloid science. This field of study was introduced in 1861 by Scottish scientistThomas Graham.
  
  

Classification[edit]

• Because the size of the dispersed phase may be difficult to measure, and because colloids have the appearance of solutions, colloids are sometimes identified and characterized by their physico-chemical and transport properties. For example, if a colloid consists of a solid phase dispersed in a liquid, the solid particles will not diffuse through a membrane, whereas with a true solution the dissolved ions or molecules will diffuse through a membrane. Because of the size exclusion, the colloidal particles are unable to pass through the pores of an ultrafiltration membrane with a size smaller than their own dimension. The smaller the size of the pore of the ultrafiltration membrane, the lower the concentration of the dispersed colloidal particules remaining in the ultrafiltered liquid. The exact value of the concentration of a truly dissolved species will thus depend on the experimental conditions applied to separate it from the colloidal particles also dispersed in the liquid. This is particularly important for solubility studies of readily hydrolysed species such as Al, Eu, Am, Cm, or organic matter complexing these species. Colloids can be classified as follows:

  Medium / Phases		Dispersed phase
  		Gas	Liquid	Solid
  Continuous medium	Gas	NONE (All gases are mutually miscible)	Liquid aerosol Examples: fog, mist, hair sprays	Solid aerosol Examples: smoke, cloud, air particulates
  	Liquid	Foam Example: whipped cream, Shaving cream	Emulsion Examples: milk, mayonnaise, hand cream	Sol Examples: pigmented ink, blood
  	Solid	Solid foam Examples: aerogel, styrofoam,pumice	Gel Examples: agar, gelatin, jelly	Solid sol Example: cranberry glass

  Based on the nature of interaction between the dispersed phase and the dispersion medium, colloids can be classified as: Hydrophilic colloids: These are water-loving colloids.The colloid particles are attracted toward water. They are also called reversible sols. Hydrophobic colloids: These are opposite in nature to hydrophilic colloids. The colloid particles are repelled by water. They are also called irreversible sols.

  
  

  In some cases, a colloid can be considered a homogeneous mixture. This is because the distinction between "dissolved" and "particulate" matter can be sometimes a matter of approach, which affects whether or not it is homogeneous or heterogeneous.

  
  

Hydrocolloids[edit]

• A hydrocolloid is defined as a colloid system wherein the colloid particles are hydrophilic polymers dispersed in water. A hydrocolloid has colloid particles spread throughout water, and depending on the quantity of water available that can take place in different states, e.g., gel or sol (liquid). Hydrocolloids can be either irreversible (single-state) or reversible. For example, agar, a reversible hydrocolloid ofseaweed extract, can exist in a gel and solid state, and alternate between states with the addition or elimination of heat.

  Many hydrocolloids are derived from natural sources. For example, agar-agar and carrageenan are extracted from seaweed, gelatin is produced by hydrolysis of proteins of bovine and fish origins, and pectin is extracted from citrus peel and apple pomace.

  Gelatin desserts like jelly or Jell-O are made from gelatin powder, another effective hydrocolloid. Hydrocolloids are employed in food mainly to influence texture or viscosity (e.g., a sauce). Hydrocolloid-based medical dressings are used for skin and wound treatment.

  Other main hydrocolloids are xanthan gum, gum arabic, guar gum, locust bean gum, cellulose derivatives as carboxymethyl cellulose,alginate and starch.

  Based on the nature of interaction between the dispersed phase and the dispersion medium, colloids can be classified as: Hydrophilic colloids: These are water-loving colloids.The colloid particles are attracted toward water. They are also called reversible sols. Hydrophobic colloids: These are opposite in nature to hydrophilic colloids. The colloid particles are repelled by water. They are also called irreversible sols.

  
  

nteraction between particles[edit]

• The following forces play an important role in the interaction of colloid particles:

  • Excluded volume repulsion: This refers to the impossibility of any overlap between hard particles.
  • Electrostatic interaction: Colloidal particles often carry an electrical charge and therefore attract or repel each other. The charge of both the continuous and the dispersed phase, as well as the mobility of the phases are factors affecting this interaction.
  • van der Waals forces: This is due to interaction between two dipoles that are either permanent or induced. Even if the particles do not have a permanent dipole, fluctuations of the electron density gives rise to a temporary dipole in a particle. This temporary dipole induces a dipole in particles nearby. The temporary dipole and the induced dipoles are then attracted to each other. This is known as van der Waals force, and is always present (unless the refractive indexes of the dispersed and continuous phases are matched), is short-range, and is attractive.
  • Entropic forces: According to the second law of thermodynamics, a system progresses to a state in which entropy is maximized. This can result in effective forces even between hard spheres.
  • Steric forces between polymer-covered surfaces or in solutions containing non-adsorbing polymer can modulate interparticle forces, producing an additional steric repulsive force (which is predominantly entropic in origin) or an attractive depletion force between them. Such an effect is specifically searched for with tailor-made superplasticizers developed to increase the workability of concrete and to reduce its water content.

Preparation[edit]

• 
  

  There are two principal ways of preparation of colloids:^[3]

  • Dispersion of large particles or droplets to the colloidal dimensions by milling, spraying, or application of shear (e.g., shaking, mixing, or high shear mixing).
  • Condensation of small dissolved molecules into larger colloidal particles by precipitation, condensation, or redox reactions. Such processes are used in the preparation of colloidal silica or gold.

  
  

Stabilization (peptization)[edit]

• 
  

  
  

  The stability of a colloidal system is the capability of the system to remain as it is.

  Stability is hindered by aggregation and sedimentation phenomena, which are driven by the colloids tendency to reduce surface energy. Reducing the interfacial tension will stabilize the colloidal system by reducing this driving force.

  
  

  

  Examples of a stable and of an unstable colloidal dispersion.

  Aggregation is due to the sum of the interaction forces between particles.^[4][5] If attractive forces (such as van der Waals forces) prevail over the repulsive ones (such as the electrostatic ones) particles aggregate in clusters.

  Electrostatic stabilization and steric stabilization are the two main mechanisms for stabilization against aggregation.

  • Electrostatic stabilization is based on the mutual repulsion of like electrical charges. In general, different phases have different charge affinities, so that an electrical double layer forms at any interface. Small particle sizes lead to enormous surface areas, and this effect is greatly amplified in colloids. In a stable colloid, mass of a dispersed phase is so low that its buoyancy or kinetic energy is too weak to overcome the electrostatic repulsion between charged layers of the dispersing phase.
  • Steric stabilization consists in covering the particles in polymers which prevents the particle to get close in the range of attractive forces.

  A combination of the two mechanisms is also possible (electrosteric stabilization). All the above mentioned mechanisms for minimizingparticle aggregation rely on the enhancement of the repulsive interaction forces.

  Electrostatic and steric stabilization do not directly address the sedimentation/floating problem.

  Particle sedimentation (and also floating, although this phenomenon is less common) arises from a difference in the density of the dispersed and of the continuous phase. The higher the difference in densities, the faster the particle settling.

  • The gel network stabilization represents the principal way to produce colloids stable to both aggregation and sedimentation.^[6][7]

  The method consists in adding to the colloidal suspension a green biopolymer able to form a gel network and characterized by shear thinning properties. Examples of such substances are xanthan and guar gum.

  
  

  

  Steric and Gel network stabilization.

  Particle settling is hindered by the stiffness of the polymeric matrix where particles are trapped.^[6] In addition, the long polymeric chains can provide a steric or electrosteric stabilization to dispersed particles.

  The rheological shear thinning properties find beneficial in the preparation of the suspensions and in their use, as the reduced viscosity at high shear rates facilitates deagglomeration, mixing and in general the flow of the suspensions.

Destabilization (flocculation)[edit]

• 
  

  
  

  Unstable colloidal dispersions can form flocs as the particles aggregate due to interparticle attractions. In this way photonic glasses can be grown. This can be accomplished by a number of different methods:

  • Removal of the electrostatic barrier that prevents aggregation of the particles. This can be accomplished by the addition of salt to a suspension or changing the pH of a suspension to effectively neutralize or "screen" the surface charge of the particles in suspension. This removes the repulsive forces that keep colloidal particles separate and allows for coagulation due to van der Waals forces.
  • Addition of a charged polymer flocculant. Polymer flocculants can bridge individual colloidal particles by attractive electrostatic interactions. For example, negatively-charged colloidal silica or clay particles can be flocculated by the addition of a positively-charged polymer.
  • Addition of non-adsorbed polymers called depletants that cause aggregation due to entropic effects.
  • Physical deformation of the particle (e.g., stretching) may increase the van der Waals forces more than stabilization forces (such as electrostatic), resulting coagulation of colloids at certain orientations.

  Unstable colloidal suspensions of low-volume fraction form clustered liquid suspensions, wherein individual clusters of particles fall to the bottom of the suspension (or float to the top if the particles are less dense than the suspending medium) once the clusters are of sufficient size for the Brownian forces that work to keep the particles in suspension to be overcome by gravitational forces. However, colloidal suspensions of higher-volume fraction form colloidal gels with viscoelastic properties. Viscoelastic colloidal gels, such asbentonite and toothpaste, flow like liquids under shear, but maintain their shape when shear is removed. It is for this reason that toothpaste can be squeezed from a toothpaste tube, but stays on the toothbrush after it is applied.

Image of Colloid