9. Solubilities of Ionic and Molecular Substances

Transcription

1 9. Solubilities of Ionic and Molecular Substances What you will accomplish in this experiment You ll investigate the like dissolves like rule for predicting the ability of a solute to dissolve in a given solvent. You ll confirm your predictions by assessing the solubility of a variety of ionic and molecular substances in solvents which have varying degrees of polarity and hydrogen-bonding capability. oncepts you need to know to be prepared Most of the substances you encounter in everyday life are mixtures. You should recall from one of your earliest lecture discussions that a mixture that s uniform throughout and doesn t separate into phases is called a homogeneous mixture. If a sample is taken from any part of a homogeneous mixture and chemically analyzed, the same components in the same concentration would be found in each and every sample. The most familiar type of homogeneous mixture is a liquid solution: a solute (which may be solid, liquid, or gaseous) is thoroughly and uniformly dispersed throughout a solvent (a liquid). If a closed container of a solution is allowed to remain standing, the components will not separate out, no matter how much time passes. There s usually a limit as to the amount of a particular solute that can be dispersed or dissolved in a given amount of a particular solvent. This limit is the solubility, and it s defined as the maximum mass of solute that can be dissolved in 100 grams (or 100 ml) of a solvent at a stated temperature. For example, you can see in the table of solubilities to the left that sucrose (table sugar) has a solubility of grams per 100 grams of water at 0 o. This means that you can dissolve up to grams of sugar, but no more, in 100 grams of water at 0 o. If more sugar is added, the excess will fall to the bottom of the container and remain undissolved. A solution in this state is said to be saturated. A solution with less than the maximum amount of dissolved solute is unsaturated.. Graham Brittain Page 1 of 10 11/4/2010

2 Solubility varies with temperature: solids and liquids will generally have a much higher solubility as the temperature of the solvent is increased. The figure to the right shows that the solubility of sucrose increases from grams (per 100 grams of water) at 0 o to approximately 300 grams at 60 o (140 o F). You ve observed this variability yourself, if you ve closely observed sugar dissolving in a glass of iced tea as compared to a cup of hot tea. Gaseous solutes tend to become less soluble as the temperature of the solvent increases. n the figure to the right, note the decrease in the solubility of sulfur dioxide (S 2 ) gas in water as the temperature is increased. And consider your own experience in the observed the loss of carbonation (the solute 2 ) in a glass of soda when the temperature is raised from ice temperature to room temperature. There are some liquid substances that can be homogeneously mixed in ALL proportions; these liquids are said to be completely miscible. Ethyl alcohol (ethanol) and water are two such substances. (You can see that the table on the previous page indicates the infinity symbol as the solubility for ethanol and water.) The solution of commercial antifreeze (ethylene glycol and water) that used in automobile radiators is another example of a homogenous mixture of miscible liquids. ow can you predict the solubility of a particular solute in a given solvent? The rule of thumb for assessing solubility is like dissolves like. This means that the more similar the polarities of the solute and solvent, the more likely it is that the two substances will form a homogeneous mixture. Polar solutes, with partial positive ( + ) and partial negative charges ( - ), will tend to dissolve in polar solvents such as water. Nonpolar solvents (such as hydrocarbons) will dissolve in other nonpolar solvents, such as oil or gasoline. Put another way, a solute is likely to dissolve in a solvent if the two substances are capable of SIMILAR types of intermolecular attractive forces. As you learned from the ne Page Lesson posted with Lab 8, there are three types of intermolecular attractive forces: 1) Dipole-dipole interactions (dipolar forces), 2) ydrogen bonds, and 3) Instantaneous dipole-dipole interactions (abbreviated as IDDI, and also called dispersion forces ). A dipole-dipole interaction is the attractive force that exists between the partial positive charge of a polar group on one molecule and the partial negative charge of a polar group on a neighboring molecule. Another way to say this is that a region of slight excess positive charge on a molecule is attracted to a complementary region (of slight excess negative charge) on a neighboring molecule.. Graham Brittain Page 2 of 10 11/4/2010

3 nly molecules that contain polar groups of atoms will have this charge imbalance (a dipole). Polar groups in organic molecules typically contain either an or a N atom. rganic molecules that are purely hydrocarbon (contain only and atoms) are nonpolar, so they are not capable of adhering by dipole-dipole interaction (electronegativities: = 2.1, = 2.5 so close that the difference is negligible). You should be able to recognize that the molecule on the top right, with a carbon-oxygen double bond, has a dipole (with a slight excess of negative charge on the more electronegative oxygen atom). Two of these polar molecules would be attracted to one another by a dipoledipole interaction. The molecule on the bottom right, made purely of carbon and hydrogen atoms, is nonpolar (has no dipole), so it is incapable of dipole-dipole interaction. ydrogen bonds are a very special, rather elite type of dipole-dipole interaction. Although called hydrogen bonds, it s essential to remember that these are still simply a type of attraction between partially-charged groups of atoms on neighboring molecules. owever, the hydrogen bond, always indicated by a dotted line (... N ), is an especially strong attractive force (about one-tenth the strength of a covalent bond). This is because it consists of two extremely electronegative atoms (two, or two N, or one and one N) with a atom between them N N... N... N Using the example (... N ), the dotted hydrogen bond is formed because the atom in the polar group is so powerfully electronegative that it draws the shared electrons in the covalent bond toward itself, leaving the positive charge of s nucleus almost fully exposed (electronegativities: = 2.1, = 3.5). The partial-positive charge of the proton is extremely attractive to electrons on nearby molecules, particularly to a lone pair of electrons on an electronegative or N atom. Thus a hydrogen bond always occurs when a + that is covalently attached to an or N atom on one molecule is interacting with a - or N atom on a neighboring molecule. In examining the pair of molecules on the right, you should now be able to recognize that both molecules contain a polar group of atoms. But because the molecule on the top has a atom covalently attached to its atom, its polar group is capable of hydrogen bonding. As the molecule on the bottom has no atom attached to its atom, it is only capable of the weaker dipole-dipole attractive force. All molecules are capable of interacting with one another via dispersion forces (more descriptively called instantaneous dipole-dipole interactions, abbreviated IDDI). To understand the origin of this attraction, you ll need to recall the clouds of electrons (orbitals) that surround the nucleus of every atom. You should never think of that electron cloud as being static, or frozen in time; rather, you should think of it as a dynamic, swirling mist. Because the electrons are constantly moving, at any given place, the electron cloud may be thin in one instant, and thick in the next. Wherever the cloud briefly thickens, the negative charge of the electrons overshadows the positive charge of the nucleus. And wherever the cloud is briefly thin, the positive charge on the nucleus is able to shine through. Thus whenever two neighboring molecules closely approach one another, the temporary charges caused by their swirling electron clouds can interact: the partially-revealed positive charge of one nucleus will be momentarily attracted to the partially-accumulated negative charge of a denser electron cloud surrounding a nucleus on the neighboring molecule. So for a brief instant, there s a very slight and very fleeting dipole, and the two molecules are weakly attracted to one another.. Graham Brittain Page 3 of 10 11/4/2010

4 All organic molecules are capable of this extremely weak instantaneous dipole-dipole interaction; however, this is the only type of intermolecular attractive force available to molecules that are purely hydrocarbon, and thus nonpolar (have no permanent dipole). After this review of intermolecular attractive forces, you should have a better appreciation for our rule of thumb for predicting solubility: like dissolves like. Again, a solute is likely to dissolve in a solvent IF the two substances are capable of SIMILAR types of intermolecular attractive forces. For example, molecules that can interact with one another by forming hydrogen bonds (such as ammonia, N 3 ) are likely to be soluble in a solvent if it can also form hydrogen bonds (such as water, 2 ). Molecules that contain polar groups of atoms (called hydrophilic groups meaning water-loving ) tend to be soluble in water and other polar solvents. Molecules that are primarily nonpolar ( hydrophobic meaning water-fearing ) are unable to dissolve in strongly polar solvents, but they are quite likely to dissolve in nonpolar hydrocarbon solvents. The figure to the right shows a polar solute (indicated by the partial charges of the permanent dipole) interacting with the partial charges of polar water molecules. Keeping the like dissolves like rule in mind, you should now be able to study the structures of the molecules below and begin making conclusions about the water solubility of each The molecule to the left is purely hydrocarbon, so it is nonpolar (hydrophobic) and would not be water-soluble. The molecule at the center has a polar group that is capable of hydrogen-bonding with water. The nonpolar hydrocarbon region of this molecule is relatively small, so the hydrogen-bonding properties of the polar group should be able to act as a tugboat, hauling the rest of the molecule into a polar solvent and causing it to be at least partially water-soluble. The molecule at the right has the same hydrogen-bonding group, but the group is attached to a nonpolar hydrocarbon chain containing seven carbons! A carbon chain this long is too hydrophobic to allow the hydrogenbonding group to coax the molecule into solution in a solvent that is as highly polar as water. This molecule is far more likely to be soluble in a nonpolar solvent (or in a solvent that is less polar than water) Graham Brittain Page 4 of 10 11/4/2010

5 An ionic compound (with FULL positively and FULL negatively-charged ions) will generally dissolve in a polar solvent such as water (that has partial positive and partial negative charges). The figure to the right shows that when ionic compounds dissolve in water, the individual cations and anions are dislodged from the crystal lattice and hydrated by the polar water molecules. The cations will be surrounded by water molecules (oriented so the partial negative oxygen atom is interacting with the positive charge of the solute cation). And the anions will be surrounded by water molecules (oriented so the partial positive hydrogen atoms are interacting with the negative charge of the solute anion). Since the cations and anions are no longer locked into their crystalline lattice, they are mobile and can move freely from one part of a solution to the other. For this reason, solutions of ionic compounds can conduct electricity. Electrolytes are solutes which can form ions when dissolved in water and conduct an electric current. Solutes that cannot ionize and conduct an electric current are called nonelectrolytes. For some ionic compounds, the attraction of the ions for one another is greater than the attraction for the partial charges of the water molecules; thus the ions are unable to break free from the structure of the crystal lattice. These ionic compounds are insoluble in water. Procedure that you will follow 1. btain a test tube rack and four clean 6 test tubes. Label each test tube as follows: Water Ethanol Acetone exane 2. Place approximately 0.1 gram of the first solute (listed in the observation table on the Report Sheet) into each test tube. (Your lab instructor will weigh 0.10 grams as a demonstration of the appropriate amount of solute; use your spatula to estimate the 0.1 gram-sized sample.) Identify the solute as either an ionic or a molecular substance. 3. Add 4.0 ml of the appropriate solvent to the labeled test tube. (You can use your graduated cylinder to add 4.0 ml of water to the first tube, then fill the other test tubes to approximately the same level.) Mix each tube thoroughly. Sharp tapping of the test tube with your fingers or stirring with a glass rod will agitate the contents.. Graham Brittain Page 5 of 10 11/4/2010

6 4. Record your observations of the solubility of the solid in each solvent: completely soluble, slightly soluble, or insoluble. Note any temperature changes as the solute dissolves. Also note any color changes, remembering that the intensity of the solution s color is a function of the amount of solute dissolved. For insoluble solutes, note whether the solute is more dense (sinks to the bottom) or less dense (floats on top) than the solvent. 5. After completing all observations, empty the tubes into a large waste beaker at your lab bench. Rinse each tube with a small amount of the specified solvent to prepare it for use with the next solute listed in the observation table. 6. When all solutes have been tested, rinse each tube with solvent, and then with soap and water, always emptying the contents into the large waste beaker at your lab station. When the waste beaker fills, take it to the hood and empty it into the appropriate liquid waste container.. Graham Brittain Page 6 of 10 11/4/2010

7 Report Sheet 9: Solubilities of Ionic and Molecular Substances Student Lab Partner Date Lab Performed Section # Lab Instructor Date Report Received Lab Notebook: Data and bservations Solute Solvent Name & hemical Formula Ionic or ovalent? Water 2 Ethanol 3 2 Acetone 3 3 exane 3 ( 2 ) 4 3 opper (II) Sulfate us 4 Sucrose Aspirin Naphthalene Graham Brittain Page 7 of 10 11/4/2010

8 Formal Report: Results and onclusions The Lewis structures for the solvents ethanol, acetone, and hexane are shown below. 1. For each structure, draw circles around the polar and nonpolar regions of the molecule, and label EA circled region with the type of intermolecular attractive force that it is capable of using. Ethanol Acetone exane 2. Draw the Lewis structure of the fourth solvent you used (a water molecule), including all covalent bonds and all lone pairs. Label the + and - regions of the molecule. Which type of intermolecular attractive force is the water molecule capable of using? 3. Rank the four solvents (ethanol, acetone, hexane, and water) according to their degree of polarity, with #1 being most polar, and #4 being least polar). #1: #2: #3: #4: 4. Does the like dissolves like rule account for your observations of opper (II) Sulfate solubility in each solvent? ompare and EXPLAIN the solubility of opper (II) Sulfate in EA solvent. Water: Ethanol: Acetone: exane:. Graham Brittain Page 8 of 10 11/4/2010

9 5. The structure of the solute Sucrose is shown to the right. In this condensed ring structure, it s understood that there s a carbon atom at each corner of the ring (except where an oxygen atom is shown). Draw circles around the various polar and nonpolar regions of the molecule, and LABEL each region with the TYPE of attractive force it can use to interact with other molecules. 2 2 Sucrose 2 Does the like dissolves like rule account for your observations of Sucrose solubility in each solvent? ompare and EXPLAIN the solubility of Sucrose in EA solvent. Water: Ethanol: Acetone: exane: 6. The structure of the solute Aspirin is shown to the right. Draw circles around the various polar and nonpolar regions of the molecule, and LABEL each region with the TYPE of attractive force it can use to interact with other molecules. Does the like dissolves like rule account for your observations of Aspirin solubility in each solvent? ompare and EXPLAIN the solubility of Aspirin in EA solvent. Water: Ethanol: Acetone: exane:. Graham Brittain Page 9 of 10 11/4/2010

10 7. The structure of the solute Napthalene is shown to the right. Draw circles around the various polar and nonpolar regions of the molecule, and LABEL each region with the TYPE of attractive force it can use to interact with other molecules. Does the like dissolves like rule account for your observations of Napthalene solubility in each solvent? ompare and EXPLAIN the solubility of Napthalene in EA solvent. Water: Ethanol: Acetone: exane: 8. The concentration of a solution (amount of solute/amount of solution) is often expressed with molarity units: 1.00 gram of Vitamin (ascorbic acid, 6 8 6, structure shown to the right) can dissolve in water to make 3.00 ml of solution. a) Illustrate how ascorbic acid molecules dissolve in water by drawing the structures of at least FUR water molecules, labeling the + and - regions of each, and sketching the way in which they interact with the ascorbic acid molecule to the right. b) Use the definition of molarity, the molar mass of ascorbic acid, and the solubility information above to calculate the concentration of this solution in molarity units (moles ascorbic acid/ liter solution). 2. Graham Brittain Page 10 of 10 11/4/2010