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Temperature
Basically, solubility increases with temperature. It is the case for most of the solvents. The situation is though different for gases. With increase of the temperature they became less soluble in each other and in water, but more soluble in organic solvents. [|**SOLVENT**] [|**(noun)**] The noun **SOLVENT** **:** a liquid substance capable of dissolving other substances.

dis·solve
 / dɪˈzɒlv  /   [ dih- **zolv**  ]   **//verb//** -solved, -solv·ing,  **//noun//** **//–verb (used with object) //**  1. to make a solution of, as by mixing with a liquid; pass into solution: //to dissolve salt in water.//

SOLUBILITY For the best view of solubility, we will use the examples of a solid solute dissolved into a liquid solvent. This does not mean that other materials do not work in the same fashion. The solubility of a solution is a measure of how much of the solute can be dissolved into the solvent. The solution reaches a point called the //saturation point// when no more solute will be accepted by the solvent. Any further addition of solute will result in solid solute mixed in with the saturated solution. Each solvent and solute pair has a characteristic solubility at a given temperature. Usually as you increase the temperature, an increased amount of solute will be able to dissolve. Take a Pyrex measuring cup and put in exactly a cup of table sugar. Heat water to boiling and pour in a small amount. Notice what happens. The volume of material in the cup appears to shrink! Continue adding boiling water until the level is back up to the 'one cup' mark. Notice the temperature of the solution. It takes heat to dissolve sugar. Stir. You should be able to almost dissolve all the sugar. The solution should be very close to the saturation point at that temperature. The solution should end up at about room temperature. Now add a few heaping tablespoons of sugar. Stir and attempt to dissolve all the sugar. If you succeed, add another few tablespoonsful of sugar. Put the saturated solution with a lot of undissolved sugar into the microwave, and heat until all the sugar is dissolved. If you have a meat thermometer, find the temperature of the boiling mixture. (Be careful. The solution is VERY hot. Handle with something to insulate you from the heat.) Observe the solution after you take it out of the microwave and put it on the counter. Notice the temperature at which the sugar crystals begin to form again. If you have done the experiment just right, you may see the crystals appearing at a temperature far below what you might think. If you boil the solution enough in the microwave, you will dissolve all traces of a seed crystal for the saturated solution to deposit sugar onto. At one time your solution will be //supersaturated//, or beyond the normal amount of solute in the solution. Supersaturation is an unstable condition. If any crystal is presented to a supersaturated solution, the crystallization of the solute onto it will occur fairly rapidly. At home if you have done this demonstration with only sugar and water in a clean cup, don't waste the sugar solution. A little bit of maple flavoring will make it a fine syrup for pancakes, or you can use it in the [|frosting] of the chocolate cake I have published here on the site. Do not eat any material made at school. Lab materials may contain traces of contaminants. If you eat anything in the school laboratory, the school lawyers will turn green and purple, have a conniption fit, and likely take their discomfort out upon you. Solubility of salts depends upon the type of ions in the salt. There is a very great range of solubility of salts in water. Even the most insoluble, such as silver chloride, have a very small but detectable solubility. Some salts, called //deliquescent salts//, are so soluble that they grab water molecules out of the air and can dissolve themselves in this way. Using the simplification of classifying materials as either soluble or not in water at room temperature, there are some nice easy general rules for predicting whether or not a salt will dissolve in water. These rules are useful not just for predicting how to make solutions, but ion reactions, such as a double displacement reaction, depend upon the insolubility of a salt as a possible product for the reaction to happen. Depending upon what your instructor suggests, it may be a good idea for you to know the following rules: (a) Almost all simple ionic compounds with Group I elements or ammonium ion, (NH**4**)**+**, are soluble. (b) All nitrates (NO**4**)**-**, most sulfates, (SO**4**)**2-**, and most chlorides, Cl**-**, are soluble. **Notable exceptions to this rule are: barium sulfate, BaSO**4**)**2-**, lead II sulfate, PbSO**4**)**2-**, and silver chloride, AgCl. (c) Most hydroxides, (OH)**-**, carbonates, (CO**3**)**2- **sulfides, S**2-**, and phosphates, (PO**4**)**3-**, are insoluble except for the compounds of rule (a). Barium hydroxide, Ba(OH)**2**