SAT Chemistry Liquids, Solids, and Phase Changes - Polarity and Hydrogen Bonding
SAT Chemistry Liquids, Solids, and Phase Changes - Polarity and Hydrogen BondingWater is different from most liquids in that it reaches its greatest density at 4°C and then its volume begins to expand. By the time water freezes at 0°C, its volume has expanded by about 9 percent. Most other liquids contract as they cool and change state to a solid because their
molecules have less energy, move more slowly, and are closer together. This abnormal behavior of water can be explained as follows. X-ray studies of ice crystals show that H2O molecules are bound into large molecules in which each oxygen atom is connected through hydrogen bonds to four other oxygen atoms as shown in Figure 30
This rather wide open structure accounts for the low density of ice. As heat is applied and melting begins, this structure begins to collapse but not all the hydrogen bonds are broken. The collapsing increases the density of the water, but the remaining bonds keep the structure from completely collapsing. As heat is absorbed, the kinetic energy of the molecules breaks more of these bonds as the temperature rises from 0° to 4°C. At the same time this added kinetic energy tends to distribute the molecules farther apart. At 4°C these opposing forces are in balance—thus the greatest density. Above 4°C the increasing molecular motion again causes a decrease in density since it is the dominate force and offsets the breaking of any more hydrogen bonds.
This behavior of water can be explained by studying the water molecule itself. The water molecule is composed of two hydrogen atoms bonded by a polar covalent bond to one oxygen atom.
Because of the polar nature of the bond, the molecule exhibits the charges shown in the above drawing. It is this polar charge that causes the polar bonding discussed in Chapter 3 as the hydrogen bond. This bonding is stronger than the usual molecular attraction called van der Waals forces or dipole-dipole attractions.
Water is often referred to as “the universal solvent” because of the number of common sub¬stances that dissolve in water. When substances are dissolved in water to the extent that no
more will dissolve at that temperature, the solution is said to be saturated. The substance dissolved is called the solute and the dissolving medium is called a solvent. To give an accu¬rate statement of a substance’s solubility, three conditions are mentioned: the amount of solute, the amount of solvent, and the temperature of the solution. Since the solubility varies for each substance and for different temperatures, a student must be acquainted with the use of solubility curves such as those shown in Figure 31.
These curves show the number of grams of solute that will dissolve in 100 grams (milliliters) of water over a temperature range of 0°C to 100°C. Take, for example, the very lowest curve at 0°C. This curve shows the number of grams of KClO3 that will dissolve in 100 grams of water over a temperature range of 0°C to 100°C. To find the solubility at any particular temperature, for example at 50°C, you follow the vertical line up from 50°C until it crosses the curve. At that point you place a ruler horizontally across the page and take the reading on the vertical axis. This point happens to be slightly below the 20 gram mark, or 18 grams. This means that 18 grams of KClO3 will dissolve in 100 grams (milliliters) of water at 50°C.
As a soluble solute is added to water at a given temperature, the solute will continue to go into solution until the water cannot quantitatively dissolve any more solute. This may take some time to be achieved depending on the rate at which the solute dissolves. Factors that influence the rate of making a solution are summarized after the General Rules of Solubility found below. At this point, solid solutes, like KClO3, will generally be found at the bottom of the container and appear not to further dissolve. In reality, portions of the undissolved solute continue to go into solution as previously dissolved solute particles re-crystallize. This type of condition—when two opposing processes [like dissolving and crystallization] equal each other in rate—is called an equilibrium state. At this point the solution is holding the maximum amount of solute that it can contain and is referred to as a saturated solution. If more water is added to a saturated solution, then more solute can be dissolved into it. A solution that contains less solute than a saturated solution is described as being unsaturated and the container will show no undissolved solute.
PROBLEM USING THE SOLUBILITY CURVE:
A solution contains 20 grams of KClO3 in 200 grams of H2O at 80°C. How many more grams of KClO3 can be dissolved to saturate the solution at 90°C?
Reading the graph at 90° and up to the graph line for KClO3, you find that 100 grams of H2O can dissolve 48 grams. Then 200 grams can hold (2 x 48) grams or 96 grams. Therefore, 96 g - 20 g = 76 g KClO3 can be added to the solution.
General Rules of Solubility
All nitrates, acetates, and chlorates are soluble.
All common compounds of sodium, potassium, and ammonium are soluble.
All chlorides are soluble except those of silver, mercury(I), and lead. (Lead chloride is noticeably soluble in hot water.)
All sulfates are soluble except those of lead, barium, strontium, and calcium. (Calcium sulfate is slightly soluble.)
The normal carbonates, phosphates, silicates, and sulfides are insoluble except those of sodium, potassium, and ammonium.
All hydroxides are insoluble except those of sodium, potassium, ammonium, calcium, barium, and strontium.
Some general trends of solubility are shown in the chart below.
Factors That Affect Rate of Solution Making (How Fast They Go Into Solution)
Summary of Types of Solutes and Relationships of Type to Solubility
Generally speaking, solutes are most likely to dissolve in solvents with similar characteristics; that is, ionic and polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents.
It should also be mentioned that polar molecules that do not ionize in aqueous solution (e.g., sugar, alcohol, glycerol) have molecules as solute particles; polar molecules that par¬tially ionize in aqueous solution (e.g., ammonia, acetic acid) have a mixture of molecules and ions as solute particles; and polar molecules that completely ionize in aqueous solution (e.g., hydrogen chloride, hydrogen iodide) have ions as solute particles.