Solubility of Ionic Compounds and Other Substances: A Laboratory Investigation, Study notes of Chemistry

Download Solubility of Ionic Compounds and Other Substances: A Laboratory Investigation and more Study notes Chemistry in PDF only on Docsity! 5.1 Solubility and the Property of Liquids - Pre-Lab Questions Name: Date: Instructor: Section/Group: 1. Describe the groups of elements that form ionic compounds. 2. Calculate maximum amount of sodium chloride that can be dissolved in 2.0 g of water at 20 °C? (Refer to Figure 5 in the Background section.) 3. How many grams of sodium nitrate, NaNO3, are soluble in 100 g of water at 10 °C? 4. Seltzer is a beverage made by dissolving carbon dioxide in water. Carbon dioxide (CO2) solubility is about 1.7 grams per kilogram in water at room temperature. What compound is the solute? What compound is the solvent? 5. Is sugar more soluble in hot tea or cold tea? What does this tell us about the solubility of sugar as a function of temperature? Downloads 5.2 Solubility and the Property of Liquids - Introduction A solution is a homogeneous mixture of substances containing one or more solutes dissolved in a solvent. The solvent is the component of a solution that is present in the greatest amount (majority component). The solute is the substance that is dissolved in a solution and is present in the least amount (minority component). Both the solvent and solutes may be solids, liquids, or gases. In the General Chemistry lab, the solvent is most often a liquid. For this experiment, you will be working with solid and liquid solutes placed into liquid solvents. Background Solubility Some compounds dissolve easily in a given solvent. Others will not dissolve at all. An easy way to determine whether a substance will dissolve is the “like dissolves like” expression where polar substances dissolve in polar substances and non-polar substances dissolve in non-polar substances. Additionally, polar substances will not dissolve in non-polar substances (and vice versa). To better understand this concept, compounds can be separated into four categories based on the types of chemical bonds they contain: ionic, hydrogen, polar covalent, and non- polar covalent. Figure 2 displays the structures of these different compounds. Compound Bonding Categories Ionic compounds form between metal and non-metal elements from opposite sides of the periodic table. Their metal and non-metal elements vary greatly in electronegativity (the power of an atom to attract electrons to itself). This results in the formation of ions when electrons are transferred from the less electronegative element (metal), to the more electronegative element (non-metal). Many ionic compounds will dissolve in polar solvents because charged ions interact with partially charged solvent molecules. Polar covalent molecules are composed of nonmetal elements whose electronegativity differences are between 0.4 and 2.0. The electronegativity difference creates polar bonds within the molecule. The overall polarity of the molecule is then dependent on its molecular shape. Symmetry in a molecule, even one with very polar bonds, can create an overall non-polar molecule. Figure 1. Polar and non-polar compounds. Figure 5. Solubility of sodium chloride (NaCl) in water to form unsaturated and saturated solutions. Temperature Effects Solubility depends on temperature. Some substances dissolve better at high temperatures, and some dissolve better at low temperatures. Table sugar (sucrose) is an example of a substance with a higher solubility in water at high temperatures than at low temperatures. A large amount of sugar can be dissolved in hot water. If a stick is placed in the hot water solution, and the solution is slowly cooled down, sugar comes out of solution and attaches to the stick. This is how rock candy is made (Figure 6). Figure 6. Rock candy on a stick. Intermolecular Forces All liquid molecules attract each other, but some attract more strongly than others. Molecules interact due to intermolecular forces called London dispersion forces. The dispersion force is the only way non-polar molecules can interact. The dispersion force is based on surface area, so the larger molecules have stronger dispersion forces. Polar molecules attract each other by dipole- dipole attraction in addition to the London dispersion force. A polar molecule, even one that cannot hydrogen bond, will generally have stronger intermolecular forces than a non-polar molecule unless the non-polar molecule is very large. Molecules that can hydrogen bond attract most strongly; hydrogen bonding can be considered an extreme form of dipole-dipole attraction. Solution Properties Affected by Intermolecular Forces A molecule inside a liquid is attracted in all directions and feels no net force. However, a molecule on the surface of a liquid experiences a net force pulling it inside. Therefore, all liquids have surface tension, which acts like a “liquid skin.” The strength of the skin is directly related to the strength of the intermolecular attractions of the molecules. The term viscosity means a liquid’s resistance to flow. If a liquid has a high viscosity, it does not flow easily. This is why honey pours so slowly. A low viscosity means it flows easily. The stronger the intermolecular forces, the higher the viscosity. The vapor pressure of a liquid determines how quickly a liquid evaporates. A liquid with strong intermolecular forces will have a low vapor pressure and will evaporate slowly. A liquid with a high vapor pressure will evaporate quickly and is described as volatile. 5.3 Solubility and the Property of Liquids - The Experiment This experiment consists of five parts where you will: 1. dissolve different substances and observe the “like dissolves like” rule; 2. determine the solubility of potassium chlorate (KClO3) at a given temperature; 3. observe and measure the surface tension of various liquids; 4. determine the melting point of a solid compound; and 5. determine the boiling point of a liquid. Part IV: Determining the Melting Point of Molecular and Ionic Compounds 1. Obtain ~0.500 g each of sucrose, succinic acid, sodium chloride, and potassium iodide. 2. Place the sucrose in a clean dry watch glass and obtain a capillary tube that’s open on one end and closed on the other. Figure 7. Packing a capillary tube. Note: Your instructor may wish to demonstrate how to pack the capillary tube. 3. To pack the tube, the open end of the capillary tube is pressed gently into a small amount of the sample on a watch glass. The tube should contain a few millimeters of the solid sample. 4. Using a long glass or stiff plastic tube, drop the capillary tube (closed side down) so that compound packs down to the closed side.Your long tube can be as short as about 18 inches, but longer is fine as well. 5. Once the capillary tube is packed, place it in the melting point device and start increasing the temperature. 6. If you don’t know the theoretical melting point, try for about a 5 °C/minute rate of temperature increase. Carefully monitor the capillary tube and thermometer. Note the temperature at which the compound starts to melt (it appears wet or glossy), and then the temperature at which the melting is complete. 7. If the compound melts and you miss recording the precise temperature, repeat at a slower rate of temperature increase. 8. Repeat Steps 2-6 for succinic acid, sodium chloride, and potassium iodide and record all your melting temperature ranges. If you get to the top limit of your device (either digitally or as read on the thermometer) and the compound still has not melted, record “Did not melt.” 9. State whether your compound is a molecular compound (covalently bonded) or an ionic compound and describe your reasoning. Part V: Determining the Boiling Point of a Liquid 1. Construct a boiling point apparatus as shown in Figure 8. The thermometer bulb should be as close to the bottom of the test tube as possible in order to measure the temperature at the bottom of the test tube as accurately as possible. 2. Use a 600-mL beaker filled with ~400 mL of tap water and a small test tube to hold your liquid. CAUTION: The liquids you are trying to determine the boiling point of in this procedure could be flammable. For this reason, NEVER use an open flame as a heat source. ALWAYS use a hot plate. Remember that as you increase a liquid’s temperature its vapor pressure also rises, increasing the amount of potentially flammable vapor in the immediate vicinity of the heat source. You can lessen this danger by performing this determination in a ventilation hood or in a very well- ventilated area. Figure 8. Set-up for determining the boiling point of a liquid. 3. Obtain ~1 mL of isopropanol and place it in the test tube. Add a couple of boiling chips to the test tube. Place the test tube in the test tube clamp. 4. Turn on the heat source and increase the temperature of the water surrounding the test tube. CAUTION: The liquid inside the test tube may be hard to see, but DO NOT hold your face, and especially your eyes, over the mouth of the test tube in order to see better. 5. Record the temperature when the liquid boils and vaporizes. If the liquid boils too quickly and you do not note the temperature, redo the determination. 6. After determining the boiling point of the isopropanol, remove the test tube from the apparatus. Leave the thermometer in the water and allow the water to continue to heat up. When it boils, record its temperature. Note: A liquid is not boiling when you first see bubbles start to form on the bottom of the container. The first bubbles to appear do indicate, however, that the liquid is approaching its boiling point. To be sure of your reading, wait until the temperature stops increasing – THEN record the temperature. Boiling Point Temperature ____________________°C Part III: Properties of Liquids Trials Number of Drops Water Acetone Vegetable Oil Trial 1 Trial 2 Trial 3 Average Calculations 1. Using the data collected in Part II, determine the total mass of the solution. 2. Determine the mass of KClO3 dissolved in solution. 3. Determine the mass of water in the solution. 4. Dividing the mass of KClO3 dissolved in solution by the mass of the water gives the solubility of KClO3 in one gram of water. What is the solubility of KClO3 in one gram of water at your temperature? Class Data Create a graph of solubility vs. temperature with the class data and attach it to your report. Temperature, °C Solubility (g/100g) Plot of Class Data