Disperse Systems - Solutions PDF

Disperse systems Solutions Asoc. Prof. Simona Sutkuvienė Solution Disperse system, consisting of two or more components. The term "Disperse System" refers to a system in which one substance (the dispersed phase) is distributed, in discrete units, throughout a second substance (the dispersed me...

Disperse systems Solutions Asoc. Prof. Simona Sutkuvienė Solution Disperse system, consisting of two or more components. The term "Disperse System" refers to a system in which one substance (the dispersed phase) is distributed, in discrete units, throughout a second substance (the dispersed medium, continuous phase ). The term "Disperse System“ refers to a system in which one substance (the dispersed phase) is distributed, in discrete units, throughout a second substance (the dispersed medium, continuous phase ). Each phase can exist in solid, liquid or gaseous state. TYPES OF DISPERSE SYSTEM Dispersed phase Dispersed Gas Liquid Solid Medium None Liquid aerosol Solid aerosol (All gases are (fog, hair sprays) (smoke cloud, Gas mutually miscible dust) (homogenous systems) Foam Emulsion Sol (soap, beer foams; (milk, (blood, Liquid whipped cream, mayonnaise) pigmented ink) shaving cream) Solid foam Gel Solid sol Solid (pumice foam) (agar, gelatine (jewel, gemstone) jelly, opal) TYPES OF DISPERSE SYSTEM Particles size a, Class Characteristics of system nm Molecular dispersion 500 Don't diffuse, visible under Heterogenous systems microscope (Suspension & emulsion) Coarse dispersions/ Heterogenous dispersion heterogenous systems a > 500 nm ( blood, milk); Colloid dispersion/ Heterogenous dispersion microheterogenous system a = 1 – 500 nm (plasma, macromolecular solutions); Molecular dispersion/ True solutions homogenous system a 100 nm Can be filtered. Must stir to stay suspended. Example: Blood is a susspension. Its red and white cells can be easily separated from the plasma by a centrifuge. COARSE DISPERSION SUSPENSION AND EMULSION Emulsion - suspension of liquid droplets (dispersed phase) of certain size within a second immiscible liquid (continuous phase). Classification of emulsions Based on dispersed phase Oil in Water (O/W): Oil droplets dispersed in water (milk) Water in Oil (W/O): Water droplets dispersed in oil (butter) Stable suspensions of liquids constituting the dispersed phase, in an immiscible liquid constituting the continuous phase is brought about using emulsifying agents such as surfactants. Stability of emulsions may be engineered to vary from seconds to years depending on application. COLLOIDAL DISPERSION a system in which particles of colloidal size (a=1-100 nm) of any nature (e.g. solid, liquid or gas) are dispersed in a continuous phase of a different composition (or state). HISTORY OF COLLOIDS The word "Colloid" was derived from the Greek, "kolla" for glue, as some of the original organic colloidal solutions were glues. This term was first coined in 1862 by the father of Physical chemistry, Thomas Graham to distinguish colloids from crystalloids such as sugar and salt. “colloid” - substances that do not diffuse through a semi-permeable membrane; Thomas Graham 1805–1869 “ crystalloid” - those which do diffuse and which are therefore in true solution. TYPES OF COLLOIDAL SYSTEMS Classification based on the nature of interaction between dispersed phase and dispersion medium Colloidal systems, depending on the nature of attraction between the dispersed phase and the dispersion medium are classified into: Lyophilic Colloids Lyophobic Colloids Associated Colloids  “Solvent-loving” colloids (Lyo means solvent and philic means loving). When these colloids are mixed with the suitable liquid, high force of attraction exists between colloidal particles and liquid. This result in formation of very stable LYOPHILIC solution called lyophilic sol. These sols are formed by COLLOIDS substances like gums, starch and proteins. Lyophilic sol can be easily prepared by directly mixing colloid with the liquid. Water loving colloids are called Hydrophilic colloids and colloidal dispersion formed is called Hydrophilic sol. STABILITY & REVERSIBILITY of lyophilic colloids In Lyophilic sol, forces of interaction between colloidal particles and liquid are quite strong. Hence, Lyophilic Sols are very stable and do not precipitate/coagulate easily. However addition of very large quantities of electrolytes can cause particles to precipitate. If large quantity of liquid is added to precipitations or the colloidal solution is stirred properly lyophilic sols can regain their original state. This shows that lyophilic sols are also reversible in nature.  “solvent-hating” colloids (Lyo means solvent and phobic means hating). When these colloids are mixed with the suitable liquid, very weak force of attraction exists between colloidal particles and LYOPHOBIC liquid and system does not pass into colloidal state readily. COLLOIDS Therefore, lyophobic sols are difficult to prepare. Special techniques are employed to prepare these sols. Water hating colloids are called hydrophobic colloids and colloidal dispersion formed is called hydrophobic sol. STABILITY & REVERSIBILITY of lyophobic colloids In Lyophobic sol, weak forces of interaction exist between colloidal particles and liquid. Hence, lyophobic sols are less stable. Addition of even small quantities of electrolytes can cause particles to precipitate. Unlike Lyophilic colloids, the precipitations of lyophobic colloids do not regain their original state as coagulated mass cannot be dispersed into colloidal form. This shows that lyophobic sols are also irreversible in nature. FACTORS AFFECTING STABILITY OF COLLOIDAL SOLUTIONS Instability of colloidal solution is known as coagulation. Coagulation is the destabilization of colloids by neutralizing the electric charge of the dispersed phase particles, which results in aggregation* of the colloidal particles. It is induced by: Heating colloidal solutions. Increasing concentration of colloidal solutions. Adding electrolytes - decreases electrostatic repulsion. Adding another colloid containing oppositely charged particles. *Aggregation is a formation of groups of particles (aggregates) bonded to each other by van der Waals or other intermolecular forces. PREPARATION OF COLLOIDS Lyophilic and lyophobic colloidal solutions (or sols) are generally prepared by different types of methods. Some of the common methods are as follows. PREPARATION OF LYOPHILIC COLLOIDS 1)The lyophilic colloids have strong affinity between particles of dispersed phase and dispersion medium. 2) Simply mixing the dispersed phase and dispersion medium under ordinary conditions readily forms these colloidal solutions. 3) For example, the substance like gelatin, gum, starch, egg, albumin etc. pass readily into water to give colloidal solution. 4) They are reversible in nature become these can be precipitated and directly converted into colloidal state.  PREPARATION OF LYOPHOBIC COLLOIDS Lyophobic sols - unstable and irreversible, 1) Condensation method difficult to 2) Dispersion method prepare. CONDENSATION METHOD In these method, smaller particles of dispersed phase are condensed suitably to be of colloidal size. This is done by the following methods. Chemical methods: a) by oxidation A colloidal solution of sulphur can be obtained by bubbling oxygen (or any other oxidising agent like HNO3, Br2 etc.) through a solution of hydrogen sulphide in water. 2H2S + O2 (or any other agent) → 2H2O + 2S 2) by reduction. Metal sols are generally prepared by this method. A number of metals such as silver, gold and platinum, have been obtained in colloidal state by treating the aqueous solution of their salts, with a suitable reducing agent such as formaldehyde, phenyl hydrazine, hydrogen peroxide, stannous chloride etc. 2AuCl3 + 3SnCl2 → 3SnCl + 2Au Gold sol The gold sol, thus prepared, has a purple colour. 3) by hydrolysis. To obtain sols of oxides or hydroxides of weakly electropositive metals. Many salt solutions are rapidly hydrolysed by boiling dilute solutions of their salts. For example, ferric hydroxide and aluminium hydroxide sols are obtained by boiling solutions of the corresponding chlorides. FeCl3 + 3H2O → Fe(OH)3 + 3HCl Colloidal sol  CONDENSATION METHOD Physical methods: 1) by exchange of solvent Colloidal solution of certain substances such as sulphur, phosphorus, which are soluble in alcohol but insoluble in water can be prepared by pouring their alcoholic solution in excess of water. For example, alcoholic solution of sulphur on pouring into water gives milky colloidal solution of sulphur. 2) by excessive cooling A colloidal solution of ice in an organic solvent like ether or chloroform can be prepared by freezing a solution of water in the solvent. The molecules of water which can no longer be held in solution, separately combine to form particles of colloidal size. DISPERSION METHOD In these methods, larger particles of a substance (suspensions) are broken into smaller particles. The following methods are employed. a) mechanical dispersion 1) In this method, the substance is first ground to coarse particles. 2) It is then mixed with the dispersion medium to get a suspension. 3) The suspension is then grinded in colloidal mill. 4) It consists of two metallic discs nearly touching each other and rotating in opposite directions at a very high speed about 7000 revolution per minute. 5) The space between the discs of the mill is so adjusted that coarse suspension is subjected to great shearing force giving rise to particles of colloidal size. 6) Colloidal solutions of black ink, paints, varnishes, dyes etc. are obtained by this method. b) by peptisation The process of converting a freshly prepared precipitate into colloidal form by the addition of suitable electrolyte is called peptisation. The electrolyte is used for this purpose is called peptizing agent or stabilizing agent. Cause of peptisation is the adsorption of the ions of the electrolyte by the particles of the precipitate. Important peptizing agents are sugar, gum, gelatin and electrolytes. Freshly prepared ferric hydroxide can be converted into colloidal state by shaking it with water containing Fe3+ or OH– ions, viz. FeCl3 or NH4OH respectively. PROPERTIES OF COLLOIDS The main properties of Colloidal Solutions are as follows: 1. Physical properties Heterogeneous nature: Colloidal sols are heterogeneousin nature. They consists of two phases; the dispersed phase and the dispersion medium. Stable nature: The colloidal solutions are quite stable. Their particles are in a state of motion and do not settle down at the bottom of the container. Filterability: Colloidal particles are readily passed through the ordinary filter papers. However they can be retained by special filters known as ultrafilters (parchment paper). 2. Mechanical properties 1) Brownian movement a) Robert Brown, a botanist discovered in 1827 that the pollen grains suspended in water do not remain at rest but move about continuously and randomly in all directions. b) Later on, it was observed that the colloidal particles are moving at random in a zig – zag motion. This type of motion is called Brownian movement. c) colloidal particles are subjected to random collision with molecules of the dispersion medium (solvent) so each particle move in irregular and complicated zigzag pathway. d) The Brownian movement explains the force of gravity acting on colloidal particles. This helps in providing stability to colloidal sols by not allowing them to settle down. 2) Diffusion As a result of Brownian motion the sol particles diffuse from higher concentration to lower concentration region. However, due to bigger size, they diffuse at a lesser speed. 3) Sedimentation a) The colloidal particles settle down under the influence of gravity at a very slow rate. This phenomenon is used for determining the molecular mass of the macromolecules. b) At small particle size (less than 5 nm) Brownian motion is significant and tend to prevent sedimentation due to gravity and promote mixing in stead. c) Use an ultracentrifuge which provide stronger force so promote sedimentation in a measurable manner. 3. Optical properties 1. Tyndall effect True solutions do not scatter light and appear clear but colloidal dispersions contain opaque particles that do scatter light and thus appear turbid. a) When light passes through a sol, its path becomes visible because of scattering of light by particles. It is called Tyndall effect. This phenomenon was studied for the first time by Tyndall. The illuminated path of the beam is called Tyndall cone. b) The intensity of the scattered light depends on the difference between the refractive indices of the dispersed phase and the dispersion medium. c) In lyophobic colloids, the difference is appreciable and, therefore, the Tyndall effect is well - defined. But in lyophilic sols, the difference is very small and the Tyndall effect is very weak. d) The Tyndall effect confirms the heterogeneous nature of the colloidal solution. e) The Tyndall effect has also been observed by an instrument called ultra – microscope.  Tyndall effect True solution Colloidal sol No scatering of light Scatering of light Some example of Tyndall effect are as follows: ✓Tail of comets is seen as a Tyndall cone due to the scattering of light by the tiny solid particles left by the comet in its path. ✓Due to scattering the sky looks blue. ✓The blue color of water in the sea is due to scattering of blue light by water molecules. ✓Visibility of projector path and circus light. ✓Visibility of sharp ray of sunlight passing through a slit in dark room. ASSOCIATED COLLOIDS (Micelles) A stearate micelle in water solution Micelle is an electrically neutral particle, which contains a colloidal particle as the nucleus (core) and two layers of ions as the shell. The particles of a colloid selectively absorb ions and acquire an electric charge. All of the particles of a given colloid take on the same charge (either positive or negative) and thus are repelled by one another.  The layer of ions closer to the nucleus - adsorption layer. The outer layer - diffusion layer. Granule has an electrokinetic potential - zeta () potential. Zeta-potential is responsible for the stability of colloidal solutions. Granules of coloidal solution have the Adsorption layer same – positive or negative - charge. That the main reason Diffusion layer why the particles of solution are stabile and do not precipitate. Sol of Fe(OH)3 is the product of hydrolysis reaction, formed by condensation method: FeCl3 + 3H2O ↔ Fe(OH)3 + 3HCl The nucleu (core) of the micelle consist of m number‘s Fe3+ and OH– ions, which are connected to insoluble crystal of (Fe(OH)3)m. NUCLEAR  STRUCTURE OF Fe(OH)3 MICELLE The molecules on the nucleu surface react with the HCl from the solution: Fe(OH)3 + HCl → FeOCl+2H2O During reaction formed iron oxy-chloride dissociates in to ions: FeOCl ↔ FeO+ + Cl- The n number of FeO+ ions are adsorbed on the surface of core and they attract the (n-x) number of oppositively charged Cl– ions. These ions form adsorption layer. Remained ions of xCl– spread out distantly from the core and form diffusion layer. The charge of colloidal particle is the same as the first layer’s ions of Diffusion layer the adsorption layer. The structure of Fe(OH)3 sol micelle could be expressed by following formula: m[Fe(OH)3]nFeO+(n-x)Cl-}  xCl-} X+ Diffusion Nucleu Adsorption layer layer GRANULE M I C E LLE True Solution True Solution is a homogeneous mixture of two or more substances in which substance dissolved (solute) in solvent has the particle size of less than 10-9 m or 1 nm. Particles of true solution cannot be filtered through filter paper and are not visible to naked eye. Simple solution of sugar in water is an example of true solution. True solution The solute is the component that is dissolved or is the least abundant component of the solution. The solvent is the dissolving agent or the most abundant component in the solution. Solutions are classified according to: PHYSICAL STATE 1) solid solutions – alloys of metal - steel is a solid solution in which iron is the solvent and carbon and manganese are the solutes; 2) liquids – physiological solutions, etc.; 3) gaseous – the air you breathe is consisting primarily of nitrogen and oxygen. In each solution, the solvent is the substance which determines the state of the finished solution. For example, when you add salt (a solid) to water (a liquid), you can tell that the water is the solvent since the resulting solution is a liquid. If both the solute and the solvent have the same state, the solvent is typically the part of the solution which is present in the highest concentration. Solutions are classified according to: ELECTRICAL CONDUCTIVITY – 1) Electrolytes are substances that solutions conduct electricity. If an ionic compound is dissolved in water, it dissociates into ions and the resulting solution will conduct electricity. E.g. Table salt is an electrolyte. 2) Non-Electrolytes are substances that don't conduct electricity. A non electrolyte does not dissociate at all in solution and therefore does not produce any ions. E.g. Sugar is a non-electrolyte.  Properties of true solutions A mixture of two or more components–solute and solvent is homogeneous and The dissolved solute is It is either colored or colorless has a variable composition. molecular or ionic in size. and is usually transparent. This means that the ratio of solvent to solute can be varied. The solute particles of a true The solute remains uniformly The solute can generally be solution are molecular or ionic distributed throughout the separated from the solvent by in size. solution and will not settle out purely physical means such as The particle size range is 0.1 nm to with time. evaporation. 1 nm (10-8 cm to10-7cm). The particles are invisible. Solubility Solubility describes the amount of solute that will dissolve in a specified amount of solvent. Terms that describe the extent of solubility of a solute in a solvent: very soluble; soluble; moderately soluble; slightly soluble; insoluble. When two liquids are completely soluble in each other (forming homogeneous solutions) they are said to be miscible (methyl alcohol and water). Liquids that are insoluble in each other they are said immiscible (oil and water). Factors related to solubility Factors that affect solubility are: nature of solute and solvent temperature pressure Nature of solute One useful classification of materials is and solvent polarity. Molecular polarity is dependent on the difference in electronegativity between atoms in a compound and the asymetry of the compound's structure. Polarity Electrons are not always shared equally between two bonding atoms: one atom might exert more of a force on the electron cloud than the other. This "pull" is termed electronegativity and measures the attraction for electrons a particular atom has. The more different electronegativity of two atoms, the more polar compound is formed. ELECTRONEGATIVITY OF ELEMENTS 4-2,1=1,9 HF 2,8-2,1=0,7 HBr HF molecule is more polar, HBr molecule is more non-polar, Substances such as H2, O2, N2, CH4, CCl4, CO2 etc. are called non-polar compounds, whereas H2O, NH3, CH3OH, NO, CO, HCl, NaCl, H2S etc. are called polar compounds. Nature of solute and solvent Solute and solvent are both polar The opposite charges help attract the particles to each other and solvation occurs. One polar, one nonpolar – There is no attraction between the particles and solvation does not occur. These substances are described as insoluble or immiscible. Polar compounds Polar compounds tend to be more soluble in polar solvents than nonpolar solvents. NaCl (sodium chloride) is ✓soluble in water Polarity Solvent ✓slightly soluble in ethyl alcohol ✓insoluble in ether and benzene Non-polar compounds Nonpolar compounds tend to be more soluble in nonpolar solvents than in polar solvents. Benzene (C6H6) is: H Polarity Solvent H C H ✓ insoluble in water C C C C H C H ✓ soluble in ether H The effect of temperature on solubility Most solutes have a limited solubility in a specific solvent at a fixed temperature. For most solids dissolved in a liquid, an increase in temperature results in increased solubility. Some solids increase in solubility only slightly with increasing temperature. Some solids decrease in solubility with increasing in temperature. Any point on any solubility curve represents a saturated solution of that solute The effect of temperature on solubility large increase in solubility with temperature slight increase in decrease in solubility with solubility with increasing temperature temperature The Effect of Pressure on Solubility Small pressure changes have little effect on the solubility of: solids in liquids liquids in liquids Small pressure changes have a great effect on the solubility of gases in liquids. The solubility of a gas in a liquid is directly proportional to the pressure of that gas above the liquid.  Saturated, Unsaturated and Supersaturated Solutions Saturated solutions A saturated solution contains the maximum quantity of solute that dissolves at that temperature. A saturated solution represents equilibrium: rate of dissolving equals to rate of crystallization. In a saturated solution two processes are occurring simultaneously: The solute is crystallizing out of solution. The solid is dissolving into the solution. Unsaturated and Supersaturated solution An unsaturated solution contains less than the maximum amount of solute that can dissolve at a particular temperature. A supersaturated solutions contain more solute than is possible to be dissolved. Supersaturated solutions are unstable. The supersaturation is only temporary, and usually accomplished in one of two ways: Warm the solvent so that it will dissolve more, then cool the solution Evaporate some of the solvent carefully so that the solute does not solidify and come out of solution. Factors affecting the “rate” of solubility The rate of solubility is how quickly something dissolves. The factors that affect the rate of solution are: ✓ Particle size ✓ Temperature ✓ Concentration of the Solution ✓ Agitation (Which means movement or stirring)  Factors affecting the “rate” of solubility Particle size – an increase in surface area to the solvent will increase rate of solution. So the particle size should be reduced by comminution before it is dissolved. A solid can dissolve only at the surface that is in contact with the solvent. Smaller crystals have a larger surface to volume ratio than large crystals. Smaller crystals dissolve faster than larger crystals. Example: A sugar cube takes longer to dissolve in a cup of tea than an equal amount of granulated sugar.  Factors affecting the “rate” of solubility Temperature In most cases, the rate of dissolving of a solid increases with temperature. At high temperatures, solvent particles move faster and solvation occurs more quickly. Example: When making sweetened iced tea, it is much easier to add the sugar while the tea is still hot. The hot temperature helps the sugar dissolve more quickly. Factors affecting the “rate” of solubility Concentration As solution concentration Δc increases, the rate of Δt dissolving decreases. The rate of dissolving Δc is at a maximum when solute and solvent are Δt first mixed.  Factors affecting the “rate” of solubility Agitation or stirring Solutes dissolve faster when the solution is agitated by stirring or shaking. When a solid is first put into water, it comes in contact only with water. The rate of dissolving is then a maximum. As the solid dissolves, the amount of dissolved solute around the solid increases and the rate of dissolving decreases. Stirring distributes the dissolved solute throughout the water; more water is in contact with the solid causing it to dissolve more rapidly. When adding sugar to coffee, stirring helps the sugar dissolve faster. Concentration of Solutions The concentration of a solution expresses the amount of solute dissolved in a given quantity of solvent or solution. Generally amounts or measures are masses, moles or liters. Dilute and Concentrated Solutions The terms dilute and concentrated are qualitative expressions of the amount of solute present in a solution. A dilute solution contains a relatively small amount of dissolved solute. A concentrated solution contains a relatively large amount of solute. Concentration of Solutions Mass Percent Solution Mass percent expresses the concentration of solution as the percent of solute in a given mass of solution. mass of component in solution mass % of component =  100 total mass of solution g solute g solute mass percent = x 100 = x 100 g solute + g solvent g solution  What is the mass percent of sodium hydroxide in a solution that is made by dissolving 12.00 g NaOH in 50.0 g H2O? grams of solute (NaOH) = 12.0 g grams of solvent (H2O) = 50.0 g  g solute  mass percent =   x 100  g solute + g solvent  12 g NaOH mass percentage =  100% = 19.3% NaOH solution 12g NaOH + 50g H 2 O Concentration of Solutions Molarity Molarity of a solution is the number of moles of solute per liter of solution. number of moles of solute moles molarity = M = = liter of solution liter ns Ms = V Concentration of Solutions Molality Is the number of moles of solute per kilogram of solvent. The molality of a solution is calculated by taking the moles of solute and dividing by the kilograms of solvent. mol of solute m= kg of solvent  Concentration of Solutions Mole fraction The molar fraction of any one component in a solution (or any other mixture) is the ratio of its number of moles to the total number of moles of all components present. Expressed mathematically: nA XA = n A + n B + n C.... + n X Where XA is the mole fraction of component A, and nA, nB, nC,.....nX are the numbers of moles of each component A, B, C,........X respectively. The sum of all mole fractions for a mixture must equal 1. Where are the mole fractions of each component is a solution consisting of 1.00 mol ethyl alcohol, 0.500 mol methyl alcohol and 6.00 mol water? We have tu apply equation nA XA = to each component, in turn. n A + n B + n C.... + n X Ethyl alcohol: 1.00 X C2 H5OH = = 0.133 Methyl alcohol: 1.00 + 0.500 + 6.00 0.500 Water: X CH3OH = = 0.067 1.00 + 0.500 + 6.00 6.00 Total =1.000 X H 2O = = 0.800 1.00 + 0.500 + 6.00 Homework What is the molarity of a solution containing 1.4 mol of acetic acid (CH3COOH) in 250 ml of solution? What is the molality of a 6.0 M nitric acid solution for which 298 g of solution contains 95 g, or 1,5 mol of HNO3? What masses of potassium chloride and water are needed to form 250 g of 5.00% solution?