Name Chemistry Pre-AP. Notes: Solutions

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1 Name Chemistry Pre-AP Notes: Solutions Period I. Intermolecular Forces (IMFs) A. Attractions Between Molecules Attractions between molecules are called and are very important in determining the properties of the compound, including boiling point, odor, and solubility. These forces are quite weak compared to the covalent bonds that hold molecules together. 1. London Dispersion Force (LDF) weakest of all IMFs found in all molecules and also in neutral atoms caused by temporary dipoles due to motion of more electrons = stronger LDF (so larger molecules have stronger LDFs) Ex: Halogens 2. Dipole-Dipole Force occur when molecules are attracted to one another (they have permanent dipoles) the slightly region of one molecule is attracted to the slightly region of another molecule Ex: 3. Hydrogen Bonding not a true bond, but rather a type of IMF special type of dipole-dipole force strongest of the IMFs What is needed for hydrogen bonding? At least one hydrogen atom in the molecule must be covalently-bonded to one of the most electronegative elements (N, O, or F), causing the hydrogen to acquire a significant amount of positive charge. Examples of compounds that exhibit H-bonding: Each of the elements to which the hydrogen is attached (N, O, or F) is not only significantly negative, but also has at least one lone pair. Lone pairs in the 2 nd energy level have the electrons contained in a relatively small volume of space, which therefore results in a high density of negative charge. (Lone pairs at higher energy levels are more diffuse and not so attractive to the hydrogen.) See next page for a diagram of hydrogen bonding between water molecules. 1

2 Q: Why is water the best at H-bonding? A: 2 hydrogens, 2 lone pairs on oxygen B. Intermolecular Forces and Molecular Properties the physical properties of a compound depend on the type (strength) of intermolecular forces in that compound Example: the greater the intermolecular forces, the the boiling point. Examine the graph below. Each set of data points is in order (left to right) from lowest to highest molar mass. Why do the compounds on the far left deviate from the normal pattern of boiling points? Note: The electrostatic (Coulombic) attractions of the ionic bond present in the lattice of an ionic crystal are stronger than any intermolecular force holding the molecules in a molecular compound together. That is why melting and boiling points of ionic compounds are so high, and why they are always solids at room temperature. If a liquid has weak intermolecular forces: If a liquid has strong intermolecular forces: - It will have a low boiling point - It will have a high boiling point - It will evaporate easily (notice odor) - It will evaporate slowly - It will have a low surface tension - It will have a high surface tension 2

3 Practice: Which would have a higher boiling point? And why? Think about IMFs! a) CCl 4 or CF 4 b) NaF or F 2 c) NH 3 or CH 4 II. Water A. Polarity of water Remember from our previous studies that water contains polar covalent bonds. Unequal sharing of the electrons between O and H sets up dipole moments in the water molecule, with a partial charge on hydrogen and a partial charge on oxygen. This uneven distribution of charge, along with water s bent shape, makes water a covalent molecule. B. Anomalies of Water due to Hydrogen Bonding Remember that water exhibits very strong bonding. These hydrogen bonds explain some very important and unusual properties of water: 1) Unusually high melting and points for its molar mass - Explained by It takes a lot of to overcome these IMFs. 2) High surface tension - Again, because of, water exhibits very high surface tension. Surface tension acts like a net across the surface of the water. - This high surface tension allows for the formation of water, allows for action in plants, and allows the water strider insect to walk on water. 3) High specific heat/high heat of vaporization - The specific heat of water is higher than many other liquids because of. This means that it takes more energy to raise the temperature of water compared to other liquids, and water requires more energy to change from liquid to gas as well. - Because of water s high specific heat, the water that covers most of earth s surface keeps temperature within limits that allow living organisms to survive. The high specific heat of water also moderates the of regions near large bodies of water. 3

4 4) Density of water - Water is very unusual in that the solid form can in the liquid form. (Ice is less dense than liquid water.) This anomaly is explained by. - As water cools and gets close to freezing, its molecules separate slightly to form the rigid structure of ice: Liquid: Solid: III. Solution Basics means capable of being dissolved (Ex: sugar is in ) A is a homogeneous mixture of two or more substances in a single. The dissolved particles are tiny: ions or small molecules. The dissolving medium in a solution is called the and the substance that dissolves is called the (Ex: when Kool Aid is dissolved in water, Kool Aid is the and water is the ) Solutions may exist as,, or Examples of Solutions Solute Solvent Example gas gas gas liquid liquid liquid liquid solid mercury in gold solid solid liquid solid An is a type of solid solution (Ex: copper and zinc mixed to create brass) 4

5 IV. The Solution Process A. Solvation forces exist between particles in a solute. They also exist between particles in a solvent. In addition, the solute and solvent particles have an attraction for each other. When a solute is placed in a, the solvent particles completely surround the surface of the (A). If the solute-solvent attraction is greater than the solute-solute and solvent-solvent attraction, then the solvent will pull particles away from the solid and surround them (B). This continues until the solute has reached (no more can dissolve). At this point, an condition exists. The rate of recrystallization equals the rate of dissolving (C). A B C B. Factors Affecting Rate of Dissolving There are several factors that affect the, or how fast, a will dissolve in a given. 1) Stirring: stirring or agitation will the rate of dissolving. 2) Temperature: temperature of the will the rate of dissolving. 3) Surface Area: the surface area of the solute will the rate of dissolving. (greater surface area = particle size). C. Solubility of a substance is the amount of that substance that will completely in a given quantity of at a certain. It is reported in g of solute/100 g of solvent at a given temperature. Note: solubility is not how a substance dissolves, but how solute will go into. 5

6 3 classifications exist regarding solutions and solubility: 1) A solution contains the maximum amount of for a given amount of at a certain. It is in. 2) An solution contains less than the maximum amount. (more solute can dissolve) 3) A solution contains more than it theoretically can hold at a given. It is an situation. It is typically made by adding solute to solvent and slowly and gently the solution. To determine if a solution is saturated, unsaturated, supersaturated, add more solute to the solution and observe. if it dissolves immediately: if it drops to bottom: if solution immediately crystallizes: D. Factors Affecting Solubility Solubility is affected by two factors: 1) nature of / combination 2) temperature 1. Nature of Solvent/Solute: Like dissolves like covalent solvents (like water) will dissolve solutes and covalent solutes. Can you explain WHY? covalent solvents (like hexane, C 6 H 14 ) will dissolve covalent solutes. Can you explain WHY? However, when a polar (or ionic) substance is mixed with a nonpolar substance, the two will not form a solution. Can you explain WHY? When two liquids are soluble in each other, they are called ; when two liquids are insoluble, they are called. Ex: oil and water are considered 6

7 Practice: Based on what you learned in the last unit about molecular polarity, predict whether or not the following solute/solvent combinations follow the rule like dissolves like : 1. carbon dioxide in water 2. magnesium sulfate in water 3. methane in carbon tetrachloride 2. Temperature 4. sodium chloride in carbon tetrachloride 5. ammonia in water 6. glucose (C 6 H 12 O 6 ) in water (glucose contains multiple O-H bonds) The of most solid substances when the of the solvent. For gases, the opposite is true (decrease in solubility with increasing temperature). This is readily seen on a curve. Note that a typical solubility curve shows the mass (in g) of a solute that will completely in 100 g of water at a given temperature. The solid line shows a solution. If the amount of solute dissolved at a given temperature is above that line, the solution is. If the amount of solute dissolved is below that line, the solution is. Solubility curves for various substances: 7

8 Using Solubility Curves Ex: I have a solution at 60 o C with 80 g of potassium nitrate dissolved in 100 g of water. Is it saturated, supersaturated, or unsaturated? Ex: How much more potassium nitrate could I add to the above to form a saturated solution? Ex: How many grams of potassium chloride will dissolve in 1.00 L of water at 50.0 o C? Ex: I have a saturated solution with 9.2 g of solute dissolved in 10 g of water at 25 o C. What is the solute? Practice 1) I have a solution with 15 g of sodium nitrate dissolved in 10.0 ml of water at 25 o C. Is it saturated, unsaturated, or supersaturated? 2) When I add 1 more seed crystal of sodium nitrate to a solution, all extra solute drops out of the solution. How much solute crystallized? 3) What volume of water do I need to make a saturated solution containing g of potassium chloride at 50 o C? 4) I have a saturated solution of lithium carbonate in 100. g of water at 20 C. How much will come out of solution when I heat it to 80 C? V. Concentration of Solutions A solution is one that contains a concentration of solute. A solution contains a concentration of solute. The of a solution is a quantitative measure of the amount of that is dissolved in a given quantity of. Many different units can be used to express the concentration of a solution. In this class, we will learn only Molarity. A. Molarity (M) is the number of of a solute dissolved per of solution. Ex. 6.0 M HCl solution contains 6.0 moles of HCl dissolved in 1.0 L of total solution. Read this as 6.0 molar HCl. Or it could be 3.0 moles dissolved in 0.50 L solution etc. Molarity (M) = 8

9 To make 1.00 liter of a 1.00 molar (1.0 M) solution: 1) Calculate the molar mass of the solute. 2) Add this mass (1 mole) of solute to a flask. 3) Add about ¼ flask of distilled water. Swirl the flask till the solute is dissolved. 4) Slowly add water while swirling until the volume reads 1.00 L. Fill to line, not top! Using Molarity Ex 1: A saline solution contains 0.90 g NaCl in exactly 100. ml of solution. What is the molarity of the solution? Ex 2: How many grams of solute are present in 1.5 L of 0.24 M Na 2 SO 4? Practice 1. A solution has a volume of 2.0 L and contains 36.0 g of glucose. If the molar mass of glucose is 180. g/mol, what is the molarity of the solution? 2. How many moles of ammonium nitrate are in 335 ml of M NH 4 NO 3? 3. How many grams of solute are in 250. ml of 2.0 M CaCl 2 solution? 4. Describe how you would prepare 250. ml of a 0.20 M NaOH solution. 9

10 B. Dilutions You can the concentration of a solution by diluting it with. The dilution reduces the of solute per unit, but the total of solute in solution do not change. Moles of solute before dilution = Moles of solute after dilution Moles of solute = Molarity x volume Dilution formula: Ex: How many ml of a stock solution of 2.00 M MgSO 4 would you need to prepare ml of M MgSO 4? Practice 1. How many ml of a stock solution of 4.00 M KI would you need to prepare ml of M KI? 2. What volume must you dilute to make 50.0 ml of 0.20 M KNO 3 from 4.0 M KNO 3? 3. What is the molarity of a solution formed when you add 200. ml of water to 50. ml of 5.0 M HCl? 10

11 VI. Precipitation Reactions, Solubility Rules, and Net Ionic Equations Two solutions of ionic compounds, when mixed together, may result in the formation of a precipitate, or insoluble solid substance, through a double replacement reaction. In order to determine if a precipitate forms, consult the solubility rules: Practice: Predict products for the following pairs of aqueous solutions (DR). Determine whether or not a precipitate will occur by consulting the solubility rules above. If a precipitate forms (and therefore a reaction occurs), write a balanced equation, INCLUDING ALL STATE SYMBOLS. (aq) means aqueous/soluble; (s) means solid/insoluble/precipitate. If no reaction occurs, write NR. 1. Sodium chloride and lead (II) nitrate 2. Magnesium nitrate and silver acetate 3. Ammonium sulfate and barium chloride 11

12 Net Ionic Equations: Precipitation Reactions Show formation of a precipitate (insoluble compound) when 2 aqueous solutions are reacted. Example 1: Given 2 aqueous solutions of compounds to be reacted: Na 2 CO 3 and Ca(NO 3 ) 2, write the net ionic equation for the reaction. Steps to writing net ionic equations: 1) Write the balanced equation for the double replacement reaction between the two compounds. This is known as a molecular equation. Na 2 CO 3 + Ca(NO 3 ) 2 2 NaNO 3 + CaCO 3 2) Use the solubility rules to determine whether each compound is soluble (aq) or insoluble (s) in water. Na 2 CO 3(aq) + Ca(NO 3 ) 2(aq) 2 NaNO 3(aq) + CaCO 3(s) 3) Break all soluble compounds down into their ions like this: (leave all insoluble compounds alone) 2Na + (aq) + CO 3 2- (aq) + Ca 2+ (aq) + 2NO 3 - (aq) 2Na + (aq) + 2NO 3 - (aq) + CaCO 3(s) This is known as a complete ionic equation. Use coefficients (not subscripts) to indicate the number of each ion. 4) Compare left to right sides. Any ions that are the same on both sides can be eliminated (cancelled out) to give: Ca 2+ (aq) + CO 3 2- (aq) CaCO 3(s) This is known as a net ionic equation. It shows only the important chemistry -- the formation of the insoluble compound or precipitate, which drives the chemical reaction. The ions that you eliminated are called spectator ions. In this example, the spectator ions are: Na +, NO 3 - Example 2: Write the net ionic equation for the reaction between AgNO 3 and NaCl 12

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