Guest Lecture by Dr. Chris Harmon covered material, but not as presented below, which are my notes—rap.
Water is a very unusual, even incredible substance whose amazing properties are often unappreciated because of its ubiquitousness. Water's special properties include extremely high mp and bp (0 °C & 100 °C, compare to methane, -183 °C & -161 °C, with a MW of 16 vs. water's 18); a high heat capacity (18 cal/°C mol vs. 8 cal/°C mol for methane); it has a high viscosity; its solid form is less dense than the liquid form at the same temperature (ice floats on water - very rare), it has a high dielectric constant (78.5 vs. 1.9 for hexane; 'blocks electric fields') and it has a high surface tension.
The high mp, bp, heat capacity and surface tension all predict relatively strong bonding between water molecules, H-bonding. Note environmental consequences - Earth's weather is much more pleasant because it is moderated by water, especially along coasts. Ice floating prevents "solid" seas, definitely a downer in environmental terms, many animals use surface tension to "float."
Water of course is a covalent structure: H-O-H. But what gives it its special properties is the polarity of its O-H bonds and the resultant dipole moments of the bonds and the molecule itself.
The water molecule itself is bent, with an angle of 104.5° between the hydrogens (compare to 109.5° for sp3 tetrahedron) as seen in Figure 4.1 on p 128 of your text.
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Because of the very strong dipole moments of these bonds and the very small size of the hydrogen substituents on water, a slight degree of orbital overlap occurs between adjacent water oxygens and hydrogens to give partial covalent bonds known as H-bonds (effectively, can only form with O, N, & F). Note that the partial covalent character means that they are directional!
Within solid bulk water (ice) every water molecule is bonded to 4 others.
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In liquid water the molecules are still bonded to a large degree (the heat of fusion for ice is only 13% of the heat of vaporization for ice, thus most of the H-bonds must survive melting). Of course in liquid water the bonds are very unstable (average lifetime about 10 psec = 10-11 sec), exchanging constantly to give a "flickering cluster" structure. The various properties of water arise from this structure. (Note hi bp & mp, heat cap., viscosity, and, less obviously, that ice floats. This is because the molecules are in an open lattice rather than close-packed. G&G note that close-packed molecules would only occupy about 57% of volume. This would lead one to expect that ice would float "high." It doesn't because most of the structure remains in the liquid phase at 0° C.)
Water is also an excellent solvent for polar substances since its dipolar structure enables it to insulate them from each other and it can make good dipole-dipole and dipole-charge bonds. Figure 4.2 on pg. 128 shows the hexavalent liganding of water to sodium and chloride ions to form hydration shells (For sodium ions, the waters in the inner hydration-shell exchange every 2-4 nsec.).
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We say that substances dissolved in water are hydrated. (More generally, solvents solvate, with hydration being the special case for aqueous solvation.) Anything which can H-bond (such as alcohol or acetic acid) will also of course be quite soluble.
As a general rule for solubility we can say that "Like dissolves like.
A major family of solutes are the electrolytes. An electrolyte is a substance whose solution conducts electricity. Electrolytes are ionic substances which dissociate in solution to conduct charge. In contrast, a nonelectrolyte gives a solution which does not conduct electricity. Nonelectrolytes are generally non-ionic substances.
Electrolytes are in turn classified in two broad categories:
Strong electrolytes: solutions are generally good conductors; substances are completely ionized in solution.
Weak electrolytes: solutions are not very good conductors; substances only partially ionized in solution.
Notice that electrolytes can be ionic compounds, like sodium chloride, or they can be covalent, molecular compounds like hydrogen chloride or hydrogen fluoride, both of which exist as perfectly good polar covalent molecules in the gas phase, but which react in water to dissociate into ions.
Strong Electrolytes are completely ionized in aqueous solution in concentrations of 1 M or less. Three common classes:
Virtually all inorganic salts are strong electrolytes in the sense that only ions exist in solution. For example sodium chloride dissociates into hydrated sodium ions (Na+(aq)) and hydrated chloride ions (Cl-(aq)). Note that each ion is surrounded by a "shell" of water molecules, countering and diffusing their charges, insulating them from other ions, and keeping them in solution.
Acids which completely dissociate at concentrations of 1 M or less. Thus hydrochloric acid (HCl) dissociates to give H+ and Cl-, nitric acid (HNO3) gives H+ and NO3-, and sulfuric acid (H2SO4) gives H+ and HSO4-.
Strong bases dissociate or react completely with water to give hydroxide ions. Thus sodium hydroxide (NaOH) gives Na+ and OH-, and potassium hydroxide gives K+ and OH-.
Weak electrolytes only partially dissociate in aqueous solution in concentrations of 1 M or less. The most common weak electrolytes are weak acids. For example acetic acid (CH3OOH) and hydrofluoric acid (HF) both dissociate only slightly (a few percent or less).
The most commonly used concentration term in chemistry is molarity = 1 mole of solute dissolved in enough solvent to give 1 L = M = mol/L. This is the most popular concentration unit at least in part because of the convenience of making molar solutions with volumetric flasks (show flask).
Note that molarity has various meanings in describing solutions. Commonly molarity refers to the amount of substance dissolved in one liter of liquid to give a solution. Thus if 158.52 g (1 mole) of strontium chloride is dissolved in enough water to give one liter we will have a 1 L solution of strontium chloride. Note however, that while the solution is 1M in Sr2+, it is 2 M in Cl- and 3 M in terms of particles!
There are a number of different kinds of problems.
Make up a 1.000 L solution of 0.25 M NaCl (note that water is the "default" solvent).
What is the concentration of a solution made by dissolving 10.00 g of KI in enough water to make 1.000 L
What is the concentration of a solution resulting from adding 25.0 mL of 0.60 M CaCl2 to 475 mL of water. (Note - both aqueous, so the volumes are additive.)First, I like to keep in mind that Moles = Moles, that is we have conservation of mass.
(Note the units of volume will cancel, so we don't HAVE to convert to L, though it doesn't hurt.)
First need to determine volumes:
volume1 = 0.025 L
volume2 = 0.025 L + 0.475 L = 0.500 L
What is the concentration of chloride ion in this solution?
How much 1.000 M MgSO4 is needed to make 500.0 mL of a 0.25 M solution.
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© R A Paselk
Last modified 24 February 2013