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New microsoft office word document (2)

  1. 1. Thermodynamic of Solutions .phasecomposed of only onemixturehomogeneousis asolution, achemistryIn ubstance, knownin another sdissolvedis a substancesoluteIn such a mixture, a . The solvent does the dissolving. The solution more or less takes onsolventas a characteristics of the solvent including its phase, and the solvent is commonlythe of a solute in a solution is aconcentrationthe major fraction of the mixture. The much of that solute is dissolved in the solventmeasure of how Characteristics  A solution is a homogeneous mixture  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).
  2. 2. 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 aneffervescence of carbon dioxide that has come out of solution; the dissolved gas itself is not visible since it is dissolved on a molecular level.
  3. 3.  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.
  4. 4. 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 besaturated. 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 anexothermic 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.
  5. 5. 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, volume fraction, and mole fraction. 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 See also: Solvent § Solvent classifications 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.
  6. 6. 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 andgrease 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. Preparation from constituent ingredients It is common practice in laboratories to make a solution directly from its constituent ingredients. There are three cases in practical calculation:  Case 1: amount of solvent volume is given.  Case 2: amount of solute mass is given.  Case 3: amount of final solution volume is given. In the following equations, A is solvent, B is solute, and C is concentration. Solute volume contribution is considered through ideal solution model.  Case 1: amount (mL) of solvent volume VA is given. Solute mass mB = C VA dA /(100-C/dB)  Case 2: amount of solute mass mB is given. Solvent volume VA = mB (100/C- 1/ dB )  Case 3: amount (mL) of final solution volume Vt is given. Solute mass mB = C Vt /100; Solvent volume VA=(100/C-1/ dB) mB  Case 2: solute mass is known, VA = mB 100/C  Case 3: total solution volume is know, same equation as case 1. VA=Vt; mB = C VA /100 Example: Make 2 g/100mL of NaCl solution with 1 L water Water (properties). The density of resulting solution is considered to be equal to that of water, statement holding especially for dilute solutions, so the density information is not required.
  7. 7. Freezing and Boiling Points For a solution with a liquid as solvent, the temperature at which it freezes to a solid is slightly lower than the freezing point of the pure solvent. This phenomenon is known as freezing point depression and is related in a simple manner to the concentration of the solute. The lowering of the freezing point is given by ΔT 1 = K fm where K f is a constant that depends on the specific solvent and m is the molality of the molecules or ions solute. Table 1 gives data for several common solvents.
  8. 8. Henry's law In physics, Henry's law is one of the gas laws formulated by William Henry in 1803. It states: "At a constant temperature, the amount of a given gas that dissolves in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid." An equivalent way of stating the law is that the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid. An everyday example of Henry's law is given by carbonated soft drinks. Before the bottle or can of carbonated drink is opened, the gas above the drink is almost pure carbon dioxide at a pressure slightly higher than atmospheric pressure. The drink itself contains dissolved carbon dioxide. When the bottle or can is opened, some of this gas escapes, giving the characteristic hiss. Because the partial pressure of carbon dioxide above the liquid is now lower, some of the dissolved carbon dioxide comes out of solution as bubbles. If a glass of the drink is left in the open, the concentration of carbon dioxide in solution will come into equilibrium with the carbon dioxide in the air, and the drink will go "flat". Temperature dependence of the Henry constant When the temperature of a system changes, the Henry constant will also change.[1] This is why some people prefer to name it Henry coefficient. Multiple equations assess the effect of temperature on the constant. These forms of the van 't Hoff equation are examples: where kH for a given temperature is Henry's constant (as defined in this article's first section). Note that the sign of C depends on whether kH,pc or kH,cp is used. T is any given temperature, in K
  9. 9. Thermodynamic of Solutions / ‫االسم‬ / ‫السكشن‬ / ‫شعبة‬ ‫رزق‬ ‫د/محمد‬ ‫أشراف‬ ‫تحت‬