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Chemistry of Solutions. Chapter 7. Types of Solutions.

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types of solutions
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.
water
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.
like dissolves like
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).
electrolytes and nonelectrolytes
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.
equivalents
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
example on equivalents
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?
solubility
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.
solubility example
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?
concentrations
Concentrations
  • Defined in the form
percent concentrations
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
example
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?
molarity
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.
example 1
Example 1
  • Calculate the molarity of 5.85 g of sodium chloride in 400. ml of solution.
example 2
Example 2
  • Calculate the number of grams of solute needed to make 175 ml of 3.00 M sodium nitrate?
example 3
Example 3
  • How many milliliters of 0.800 M calcium nitrate contain 0.0500 moles of this solute?
example 4
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?
example 5
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?
example 6
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?