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Chemistry of Solutions

Chemistry of Solutions. Chapter 7. Types of Solutions.

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Chemistry of Solutions

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  1. Chemistry of Solutions Chapter 7

  2. Types of Solutions • Although there are many examples of solutions in different phases – gases in gases; gases, liquids, or solids in liquids; and liquids or solids in solids – the most frequent situation in chemistry is working with something dissolved in a liquid. • A solution is a homogeneous mixture – i.e., no separation of solute and solvent, concentration the same everywhere.

  3. Water • Water is the most common solvent in a chemistry laboratory. Dissolves many materials because of its ability to form hydrogen bonds or because of its polarity. • However, water has trouble dissolving many non-polar substances, particularly organic compounds.

  4. Like Dissolves Like • Polar solvents like to dissolve polar or ionic solutes – salt in water, acetic acid in water, methanol in water, acetic acid in methanol • Nonpolar solvents like to dissolve nonpolar solutes – toluene in hexane, hexane in carbon tetrachloride • Note that surfactants work by having a nonpolar end that is attracted to nonpolar grease and an opposite polar end attracted to water to carry the grease away. Also a model for cell walls (lipid chemistry).

  5. Electrolytes and Nonelectrolytes • An electrolyte is a solute that separates into ions in water. • Differentiated by labels “strong” and “weak”. • Strong – dissociate 100% into ions. (NaCl) • Weak – stays mostly as intact molecules. Only a small portion dissociates into ions – (acetic acid, phenol, ammonia) • A nonelectrolyte does not dissociate at all (sugar, ethanol). Stays as intact molecules. • Conductivity of a solution is a good measure of strength of an electrolyte.

  6. Equivalents • Used to describe electrolyte concentrations – examples in book are taken from medical applications. • Def: an equivalent is the number of moles of an ion providing one mole of positive or negative charge. • # equivalents of ion = # moles of ion * Absolute value of charge of the ion • e.g., 0.4 moles Ca+2 = 0.8 equivalents Ca+2

  7. Example on Equivalents • A solution contains 40 mEq/L Cl- and 15 mEq/L of HPO42- . If Na+ is the only cation in the solution, what is the sodium ion concentration in milliequivalents per liter? • What are the molar concentrations of each component of the solution?

  8. Solubility • Not every solution system is completely miscible. It is possible to saturate a solution. A saturated solution has the maximum amount of solute dissolved in a solvent at a given temperature. We see this all the time with the solubility of, for example, sugar in water. • Solubility usually increases with temperature. Hence, more sugar dissolves in hot tea than in iced tea. This is because most solution processes are endothermic – they absorb heat to make them go.

  9. Solubility Example • The solubility of KCl in water: • At 20 deg C, 34 g KCl will dissolve in 100 g water • At 50 deg C, 43 g KCl will dissolve in 100 g water • A solution containing 80. g of KCl in 200. g of water at 50 deg C is cooled to 20 deg C. How many grams of KCl remain in solution at 20 deg C? How many grams of KCl crystallized from solution after cooling?

  10. Concentrations • Defined in the form

  11. Percent concentrations • Mass Percent – most common, except in medical applications • Volume Percent (volumes not strictly additive) • Mass / Volume Percent (using grams of solute, ml of solution) – seems to be commonly used in medical applications

  12. Example • A patient needs 100. g of glucose in the next 12 hours. How many liters of a 5% (m/v) glucose solution must be given?

  13. Molarity • Most common in the chemistry laboratory • Gives the number of moles of solute present in a given volume. Easy to relate back to chemical equations which operate based on moles.

  14. Example 1 • Calculate the molarity of 5.85 g of sodium chloride in 400. ml of solution.

  15. Example 2 • Calculate the number of grams of solute needed to make 175 ml of 3.00 M sodium nitrate?

  16. Example 3 • How many milliliters of 0.800 M calcium nitrate contain 0.0500 moles of this solute?

  17. Example 4 • What is the final concentration in molarity of a solution in which water is added to 25 ml of a 25% (m/v) solution of sulfuric acid until the final volume is 100.0 ml?

  18. Example 5 • How many liters of 0.50 M phosphoric acid can be made from 0.500 liter of a 6.0 M phosphoric acid stock solution?

  19. Example 6 • Lead(II) nitrate reacts with potassium chloride to produce lead(II) chloride and potassium nitrate. The lead(II) chloride precipitates as a solid and is removed from the reaction as it is formed. • Write a balanced equation for this reaction. • How many grams of lead(II) chloride will be formed from 50.0 ml of 1.50 M potassium chloride and excess lead(II) nitrate? • How many milliliters of 2.00 M lead(II) nitrate are needed to react completely with 50.0 ml of 1.50 M potassium chloride?

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