7.1a NOTES Types of Solutions

7.1a NOTES: Types of Solutions In this lesson you will... define a solution and define the terms solute and solvent. describe the composition of solutions. identify the solute and solvent in a solution. define colloids and suspensions. list and explain three factors that affect the rate of dissolving of a solid solute in a liquid solvent. explain solution equilibrium and distinguish between saturated, unsaturated, and supersaturated solutions. explain the effects of temperature on the solubility of solids and gases. use a solubility curve to determine the solubilities of substances at various temperatures. » OBJECTIVES covered: CH.4c Introduction In this module, we begin our study of solution chemistry. We all might think that we know what a solution is, listing a drink like tea or soda as an example of a solution. What you might not have realized, however, is that the air or alloys such as brass are all classified as solutions. Why are these classified as solutions? Why wouldn't milk be classified as a true solution? To answer these questions, we have to learn some specific properties of solutions. Let's begin with the definition of a solution and look at some of the different types of solutions. Why can you use certain soaps to wash your face and NOT clothes? Soaps are generally made from natural fats and oils. Soaps are excellent for cleaning our hands, face, and body, because the loosened dirt is rinsed away immediately. But soaps have definite drawbacks for cleaning laundry. For example, soaps often allow dirt lifted from clothes to redeposit on the clothes before the wash cycle is finished. And in hard water (water containing high levels of minerals), soaps react with minerals to form scum, called soap curd. Soap curd does not dissolve. It is difficult to remove from fabrics, and it makes the fabric feel stiff. Because of these drawbacks, laundry soaps have largely been replaced by detergents. Detergents are synthetic (artificial) mixtures of ingredients that not only clean clothes but also prevent redeposition of dirt, discourage scum formation, and possess other useful properties. The most important advantage of detergents is the ability to clean effectively in hard water. Determining the right cleaner largely depends on understanding Solutions. Solutions A solution is a homogeneous mixture of several substances, usually a solid dissolved in a liquid. Solutions are made up of two parts, the solvent and solute are the two basic parts of a solution. The solvent is the substance present in the greatest amount. The solute, then, is the substance present in the least amount. Page 1 of 9 What is happening? Solute particles (usually solids) move from the solid into the liquid. The process by which the positive and negative ions of an ionic solid become surrounded by solvent molecules is called Solvation. Types of Solutions • • • Saturated: The system will remain the same as long as the temperature remains constant. Such a solution is said to be saturated. A saturated solution contains the maximum amount of solute for a given quantity of solvent at a constant temperature and pressure. Unsaturated: A solution that contains less solute than a saturated solution at a given temperature and pressure is an unsaturated solution. Supersaturated: A supersaturated solution can temporarily hold more than it can. Semi- stable. Ex) honey, rock candy There are three states of matter: solid, liquid, and gas. If we think about solutions and the possibilities of combining these states together to form solutions, we have nine possibilities. Solid Solute Liquid Solute Gas solute Solid Solvent solid in solid liquid in solid gas in solid Liquid Solvent solid in liquid liquid in liquid gas in liquid Gaseous Solvent solid in gas liquid in gas gas in gas Although there are nine possibilities, there are 5 common types of solutions. EXAMPLES: 1. A solid in a solid solution would be an alloy like steel or brass. The other types of solutions are still solutions, but they are less common in everyday lives. 2. Solid in a liquid solution can be anything from a sugar solution to seawater. 3. Liquid in liquid solutions include vinegar and the antifreeze/coolant used in cars. 4. For a gas in a liquid solution, the most common example is in soft drinks: carbon dioxide is a gas dissolved in water. Another example is the ammonia solution used for cleaning. 5. Finally, to understand a gas in a gas solution, take a deep breath. Air is a solution made up of mostly oxygen gas and nitrogen gas. Heterogeneous vs Homogeneous Solutions Solutions are homogeneous mixtures, meaning that a solution is completely uniform throughout. There are two other types of mixtures that are related to a solution but are heterogeneous. Page 2 of 9 Suspensions Take a glass of water and throw in a handful of sand or dirt. Stir it and stir it and stir it. Have you made a solution? Sand and dirt do not dissolve in water, and, though it may look homogeneous for a few moments, the sand or dirt gradually sinks to the bottom of the glass A suspension is a heterogeneous mixture in which some of the particles settle out of the mixture upon standing. The particles in a suspension are far larger than those of a solution, so gravity is able to pull them down out of the dispersion medium (water). The diameter for the dispersed particles in a suspension, such as the sand in the suspension described above, is typically at least 1000 times greater than those in a solution. Unlike a solution, the dispersed particles can be separated from the dispersion medium by filtering. Suspensions are still considered heterogeneous because the different substances in the mixture will not remain uniformly dispersed if they are not actively being mixed. Colloids A colloid is a heterogeneous mixture in which the dispersed particles are intermediate in size between those of a solution and a suspension. The particles are spread evenly throughout the dispersion medium, which can be a solid, liquid, or gas. Because the dispersed particles of a colloid are not as large as those of a suspension, they do not settle out upon standing. The table below (Table below) summarizes the properties and distinctions between solutions, colloids, and suspensions. Properties of Solutions, Colloids, and Suspensions Solutions Colloids Suspensions Homogeneous Heterogeneous Heterogeneous Particle size: 1-1000 nm, Particle size: over 1000 nm, Particle size: 0.01-1 nm; dispersed; large molecules or suspended; large particles or atoms, ions, or molecules aggregates aggregates Do not separate on standing Do not separate on standing Do not separate on standing Cannot be separated by Cannot be separated by Can be separated by filtration filtration filtration May either scatter light or be Do not scatter light Scatter light (Tyndall effect) opaque • Colloids are unlike solutions because their dispersed particles are much larger than those of a solution. The dispersed particles of a colloid cannot be separated by filtration, but they scatter light, a phenomenon called the Tyndall effect. Tyndall Effect Colloids are often confused with true homogeneous solutions because the individual dispersed particles of a colloid cannot be seen. When light is passed through a true solution, the dissolved particles are too small to deflect the light. However, the dispersed particles of a colloid, being larger, do deflect light (Figure below). The Tyndall effect is the scattering of visible light by colloidal particles. You have undoubtedly "seen" a light beam as it passes through fog, smoke, or a scattering of dust particles suspended in air. All three are examples of colloids. Suspensions may scatter light, but if the number of suspended particles is sufficiently large, the suspension may simply be opaque, and the light scattering will not occur. Page 3 of 9 Learning Check #1: 1. Water is the solute in saltwater. a. True b. False 2. Which of the following are homogeneous? [mark all correct answers] a. gasoline b. chocolate c. air d. oil and water ANSWERS: 1. False, 2. a,b,c Compound Structure and Solution Formation Recall that a polar molecule is one that has a partially positive end and a partially negative end, while non-polar molecules have charges that are evenly distributed throughout the molecule. In fact, during the study of the Valence Shell Electron Pair Repulsion Theory (VSEPR) (from module IV), you learned that the chemical structures themselves have built in molecular polarity. In solution chemistry, we can predict when solutions will form using the saying, "like dissolves like." This expression means that solubility is based on the two parts of a solution having similar intermolecular forces. For example, suppose you are dissolving methanol in water. Both methanol and water are polar molecules that can form a solution because they both have permanent dipoles (partially positive and partially negative parts of the molecules) that allow the molecules of both substances to be attracted to the other. We say these substances are miscible, which means these two can dissolve in each other and make a solution. If they are immiscible, they are unable to be mixed together. The same is true for the case of a non-polar substance such as carbon tetrachloride CCl4 being dissolved in another nonpolar substance like pentane. London-dispersion forces are the intermolecular forces that hold the carbon tetrachloride together as a liquid and allow pentane to be a liquid at room temperature. Since both substances have the same intermolecular forces, when they are mixed together, a solution will be formed. In a polar solvent, the molecules of solvent are attracted to each other by the partial charges on the ends of the molecules. When a polar solute is added, the positive polar ends of the solute molecules attract the negative polar ends of the solvent molecules and vice versa. This attraction allows the two different types of molecules to form a solution. If a non-polar solute is added to a polar solvent, the non-polar solute molecules are unable to disrupt the solvent molecules. The polarity Page 4 of 9 of the solvent molecules makes the solvent molecules more attracted to each other rather than to the non-polar solute. As a result, the solute and solvent are immiscible. For example, if we were to add table salt, NaCl, to either carbon tetrachloride, we would find that the salt would not dissolve. Since carbon tetrachloride has no permanent dipoles in its molecules, there is no attraction among the charged ions in NaCl and the uncharged molecules of the solvent. Water is often referred to as the "universal solvent." The term "universal" is used to describe the fact that water can dissolve many types and kinds of substances. For instance, table salt, NaCl, is an ionic compound, but it easily makes a solution with water. This is true for many ionic compounds. Water also readily dissolves table sugar, a polar covalent compound, and other polar compounds, such as vinegar and corn syrup. Even some non-polar substances dissolve in water to a limited degree. Oxygen gas, a non-polar molecule, can dissolve in water and be taken up by fish through their gills. Carbon dioxide, another non-polar gas, is also soluble in water. Many drinks, such as sodas, beers, seltzer water, and champagne contain carbon dioxide. To keep as much gas in solution as possible, the sodas are kept under pressure. WATCH: The Dissociation of Salt (~1 min) https://youtu.be/EBfGcTAJF4o Learning Check #2: Will the following solutions form? 1. Polar substance in a polar substance. True False 2. Non-polar substance in a non-polar substance. True False 3. Polar substance in a non-polar substance. True False 4. Non-polar substance in a polar substance. True False 5. Ionic substance in a polar substance. True False 6. Ionic substance in a non-polar substance. True False Answers: 1.T, 2. T, 3. F, 4. F, 5. T, 6. F Factors Affecting Solubility Solubility is defined as the mass of a solute that dissolves in a given mass of a solvent at a specified temperature. 3 factors determine the Rate (how fast or slow) a substance dissolves: 1. Stirring(agitation)/shaking - The dissolving process occurs at the surface of the solute. Stirring speeds up the process because fresh solvent (most commonly water) is continually brought into contact with the surface of the solute (sugar). It's important to realize, however, that agitation (stirring or shaking) affects only the rate at which a solid solute dissolves. It does not influence the amount of solute that will dissolve. An insoluble substance remains undissolved regardless of how vigorously or for how long the solvent/solute system is agitated. 2. Temperature - At higher temperatures, the kinetic energy of water molecules is greater than at lower temperatures, so they move faster. The more rapid motion of the solvent molecules leads to an increase in the frequency and the force of the collisions between water molecules and the surfaces of the solute. 3. Surface area / particle size - The rate at which a solute dissolves also depends upon the size of the solute particles. For example, a spoonful of granulated sugar dissolves more quickly than a sugar cube because the smaller particles in granulated sugar expose a much greater surface area to the colliding water molecules. Page 5 of 9 Remember, the dissolving process is a surface phenomenon. The more surface area of the solute that is exposed, the faster the rate of dissolving. How much a solute will dissolve in solvent is known as Dissociation: • electrolyte - a compound that conducts an electric current when it is in an aqueous solution or in the molten state; all ionic compounds are electrolytes, but most covalent compounds are not. − strong electrolyte - a solution in which a large portion of the solute exists as ions, conducts electricity very well. − weak electrolyte - conducts electricity poorly because only a fraction of the solute in the solution exists as ions. • nonelectrolyte - a compound that does not conduct an electric current in aqueous solution or in the molten state, mostly due to not having mobile ions in solution. Solubility Solubility refers to the maximum amount of material that will dissolve in a given amount of solvent at a given temperature to produce a stable solution. By looking at the plot of solubilities below, you can see that most solids increase in solubility with an increase in temperature. (usually grams of solute per 100 grams solvent). 1. Saturated solution- Contains the maximum amount of solid dissolved. (Any point on the line) ex) (Red dot) At 20°C, 210 grams C12H22O11 2. Unsaturated solution- Can dissolve more solute. (any amount BELOW the line) ex) (Blue dot) At 20°C, 160 grams C12H22O11 3. Supersaturated- A solution that is temporarily holding more than it can. Semi- stable. Ex) honey (any amount over the line) ex) (Black dot) At 20°C, 220 grams C12H22O11 (Only if the solution is completely saturated, otherwise the extra solute would fall out of solution.) Factors that Affect solubility: • • temperature affects solubility of solids, liquids, and gases temperature AND pressure affect solubility of gases ONLY. Page 6 of 9 Solubility Curves Graphs that describe the behavior of solutes in solvents. • • • • • • Shows the relationship of temperature on solubility. a solution with an amount on the line indicates a saturated solution. a solution with an amount below the line indicates an unsaturated solution. a solution with an amount above the line indicates a supersaturated solution. Most solids and liquids increase solubility with increased temp. Most gases decrease solubility with increased temp. EXAMPLES: 1) How many grams of HCl are saturated at 40 C? ____? (66 g) 2) How many grams of HCl leave of solution when heated from 40C to 70C? _____? You must subtract the amount soluble at 40C and 70C. 66g (at 40C) - 52g (at 70C) = 14 grams 3) which substance has the lowest solubility at 30 C? _____? Find the lowest line at 30C, SO2 4) What is the amount of KCl that is saturated in 50 grams of water at 40 C? ____? you must set up a proportion. All values from the graph are in 100 grams of water. at 40 C, there are 40 grams of KCl in 100 grams of water. (40g KCl/100g H2O) so, 40g/100g = x / 50 g Learning Check #3: 1. Is a solution that contains 70g of NaNO3 at 30°C (in 100 mL H2O) saturated? True False 2. Is the solution that contains 20g of KClO3 at 50°C (in 50 mL H2O) saturated? True False 3. Find the mass of solute that will dissolve in 100 mL of water at the given temperature. KNO3 at 70°C 4. Find the mass of solute that will dissolve in 100 mL of water at the given temperature. NaCl at 100°C= ANSWER: 1. F, 2. T, 3. 135 grams, 4. 40 grams Colligative Properties • Colligative properties are properties of a solution – how a pure substance changes its physical properties when it is made into a mixture. • Depends only on the number of dissolved particles, Not what kind of particle. Ex) Pressure is increases with increasing number of particles, it does not matter what type. (oxygen, nitrogen, hydrogen, water vapor) Page 7 of 9 THREE important colligative properties: 1. Vapor-pressure lowering 2. Boiling-point elevation 3. Freezing-point depression 1.) Vapor-pressure lowering - the pressure created by vapor/gas molecules above a solution. The bonds between molecules keep molecules from escaping. In a solution, some of the solvent is busy keeping the solute dissolved. Lowers the vapor pressure. EXAMPLE: Electrolytes form ions when dissolved - more pieces. NaCl → Na+ + Cl- 2.) 2 particles: More particles, bigger effect. Boiling-point elevation - boiling point of a pure substance will increase if solute is added. The energy needed to overcome the forces between molecules increase, making it use more energy (increase temperature) to change solution from a liquid to a vapor. The vapor pressure determines the boiling point. Lower vapor pressure - higher boiling point. EXAMPLE: Saltwater boils above 100ºC, the amount of solute determines how much. 3.) Freezing-point depression - freezing point of a pure substance will decrease if solute is added. The energy needed to overcome the forces between molecules increase, requiring the temperature to drop further to change the solution from liquid to a solid. Solids form when molecules make an orderly pattern. The solute molecules break up the orderly pattern. Makes the freezing point lower. EXAMPLE: Saltwater freezes below 0ºC, how much depends on the solute. TERMS to KNOW: Suspensions - Less than 100nm particle size, a mixture in which the particles are large enough that they can settle out of solution due to the force of gravity ex) ketchup, cough syrup Colloids - (1-100 nm), are mixtures in which the particles are smaller than suspensions usually don't settle. Ex) marshmallows, cheese, paints True solution - a mixture with particle size <1nm with permanent suspension ex) household ammonia, vinegar, air Miscible - means that two liquids can dissolve in each other. Page 8 of 9 Immiscible - means they can't dissolve in each other. Electrolyte - is a substance that when dissolved in water yields a solution that conducts electric current (NaCl,...) Weak Electrolyte- is a substance that can partially dissolve in water and slightly (dimly) conduct electricity. (vinegar, ammonia) Nonelectrolyte - is a substance that when dissolved in water does NOT conduct electricity. (sugar) Solubility - The amount of substance that will dissolve at a specific temperature to produce a saturated solution. Saturated solution - Contains the maximum amount of solid dissolved. Unsaturated solution - Can dissolve more solute. Supersaturated - A solution that is temporarily holding more than it can. Semi- stable. Ex) honey Tyndall Effect - the scattering of visible light by colloidal particles, exhibited by suspensions and colloids Brownian motion - the chaotic movement of colloidal particles caused by collision with particles of the solvent in which they are dispersed Emulsion - colloidal dispersion of a liquid in a liquid Emulsifying agents - allow liquids that are not soluble to form colloidal dispersion: • • Soap is an emulsifying agent for oil and water. (cleaning purposes) Egg yolk is an emulsifying agent for vinegar and oil (makes mayonnaise) Resources: • • • • • • http://cps-www.bu.edu/Wasser/robert/work/node20.html http://www.science.uwaterloo.ca/~cchieh/cact/c120/solution.html https://www.ck12.org/book/CK-12-Chemistry-Intermediate/section/15.3/ http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/molvie1.swf http://sciencenotes.org/electrolytes-strong-weak-and-non-electrolytes/ http://www.cabarrus.k12.nc.us/site/handlers/filedownload.ashx?moduleinstanceid=37734&dataid=74059&FileName=solubility%20curve s%20worksheet.doc Page 9 of 9