Unit 10 Solutions

1. Unit 10 Solutions

2. Mixtures • Mixture – is a combination of two or more pure substances • Pure substances include elements and molecules • Examples of Mixtures: air, water from the tap, hot chocolate, paint, 14k gold, stainless steel (metal mixtures are called alloys) • Heterogeneous mixture means the composition is different throughout • Homogenous mixture means the composition is the same throughout • Solution = homogenous mixture Homogenized milk Non-homogenized milk

3. Solutions • Solutions are homogeneous mixtures • The solute is the substance to be dissolved (sugar). The solvent is the one doing the dissolving (water). • So, if I mix cream into my coffee, which is the solute and which is the solvent? • Once the solute is evenly dissolved, what do you call the new coffee? • A homogenous mixture, or solution

4. Solutions • Characteristics of Solutions: • Homogeneous • Solute will not settle out of solution • Cannot be separated by filtering • The solution has different properties than the solute and the solvent • i.e., salt and water have different individual properties than salt water • The solution is considered to be in one physical state of matter • Even though the solute and the solvent can be in different states of matter • Not all substances can combine to form a solution

5. Separating solutions • Evaporation:great for separating a solute from an aqueous solution. • The process involves heating the solution until the solvent evaporates leaving behind the solid solute • Magnetism: ideal for separating mixtures of two solids with one part having magnetic properties. • Some metals like iron, nickel and cobalt have magnetic properties whiles gold, silver and aluminum do not. • Distillation: Separates a liquid from an aqueous solution or a solution made of two liquids • Chromatography: This method is often used in the food industry. • It is used to identify chemicals (coloring agents) in foods or inks. For example, if a scientist wants to know how many substances are in a particular food dye, chromatography can be used.

6. Measuring Solutions (1) (4) • Measure solutions from eye level at the bottom of the meniscus • A graduatedcylinder(1) is used to measure the amount of a liquid and it is used to measure volume. • The traditional use of an Erlenmeyer flask (2)and a beaker (3) are to hold liquids and mix chemicals • Volumetric flasks (4) are used to make solutions and perform dilutions (2) (3)

7. States of Matter of Solutions Solutions are considered to be in one physical state of matter but can be made up of any combination of states of matter for their solutes and solvents Gas in gas: Nitrogen, oxygen, argon, carbon dioxide, water vapor, a small amount of other gases combine to make air Solid in liquid: Salt in water (= sea water) Solid in solid: Copper and nickel to make a nickel coin All metal solutions are called ____ Alloys • Can you think of an example for each? • Liquid in liquid: • Cream and coffee to make a solution • Gas in liquid: • CO2 in water to make soda

8. Solubility • Solubility refers to a solute’s ability to be dissolved by the solvent. • A substance is soluble if it is able to be dissolved in a given solvent. • A substance is said to be insoluble if it stays in its original state in the solvent (aka it does not dissolve) • Solubility is a range (very soluble, mostly soluble, partially soluble, slightly soluble, insoluble) • How do we know if a substance is soluble in water? • Solubility rules • Liquids that dissolve in one another are called miscible and liquids that do not dissolve in one another are called immiscible

9. Solubility and Polarity • Polar means having opposite ends (a positive end and a negative end) • Polar substances include: ionic compounds, acids, and polar covalent compounds • Non-polar substances include: non-polar covalent molecules and diatomic molecules • Water is a very polar substance • “Like dissolves like” • Polar substances dissolve other polar substances • Non-polar substances dissolve non-polar substances • Polar will not dissolve non-polar and vice versa • This is a generalization • Water is the universal solvent (meaning it can dissolve most substances) because it is very polar • Why do you think that oil and water won’t mix? • Oil and other fats are non-polar and don’t dissolve in water. • Soap has a polar end, so it is able to dissolve in water and a non-polar end, so it is able to dissolve grease.

10. Dissolving • Remember: polar substances, like ionic compounds and water, are composed of a positive end and a negative end • There are 2 phases of dissolving: • Dissociation- Separation of solute molecules and spreading of solute through solution • Solvation- the process of the solvent attracting and surrounding molecules or ions of a solute • For example when sodium chloride (NaCl) dissolves in water, sodium chloride dissociates (or separates into sodium ions and chlorine ions) in the solution and water solvates (or surrounds) these ions • Model this and label solvation and dissociation: Salt dissolving in water

11. Rate of Dissolution • Rate of dissolution is how quickly something dissolves • If you are making some sweet tea and you are trying to dissolve a lot of sugar, how can you make the process of the sugar dissolving faster? • 4 factors affect the rate of dissolving: surface area, agitation, temperature, and pressure • Surface area- the dissolving process occurs at the surface of the solid being dissolved. The more surface exposed the quicker the dissolving. • Which would dissolve faster: a cube of sugar or crushed sugar? • Crushed because of more surface area

12. Rate of Dissolution • Agitation- agitation such as stirring, shaking, or mixing removes newly dissolved particles form the solid surface and continuously exposes the surface to fresh solvent. • Agitation increases the rate at which solids and liquids dissolve, but decreases the rate at which a gas dissolves. • Agitating a solution increases the energy of a gas and actually causes a gas solute to escape solution. • Temperature- higher temperature causes the solvent to move more rapidly, thus increasing the rate of dissolving for solids, but decreasing the dissolving rate of gases • Solubility for gases decrease when temperature is increased because the molecules move too fast and escape rather than dissolve • This is shown on a solubility curve by a downward sloping curve

13. Rate of Dissolution • Pressure- Increasing the pressure above a solution increases the rate at which a gas dissolves. • Pressure only effects gases, NOT solids or liquids. • Increasing pressure pushes more gas molecules into solution.

14. Describing Solutions • concentration = amount of solute in a solution. • A concentrated solution has a relatively high amount of solute. • A dilute solution has a relatively low amount of solute.

15. Saturation • Do you think you could dissolve a whole bag of sugar in a cup of water? Why or why not? • There is a limit to the amount of solute that can be dissolve into a solution- called the Saturation Limit • Saturation limits change when temperature and pressure change • A solution is saturated if the solution is at its maximum concentration for the given temperature and pressure. • Equilibrium solubility is the concentration of compound in a saturated solution when excess solid is present, and solution and solid are at equilibrium • Equilibrium is when the solid and the compounds in the solution dissolve and precipitate at the same rate, so the amount of solid and dissolved ions remains constant

16. Saturation • Unsaturated solutions are solutions in which the solute concentration is lower than the saturation limit • Supersaturated solutionsare solutions in which the solute concentration is higher than the saturation limit • Process of making a supersaturated solution: • Add solute to a solution until it is saturated • Heat solution to raise the saturation limit • Add more solute to saturate the solution at the new saturation limit • SLOWLY cool the solution • At the cooler temperature, any slight disturbance, such as agitation or adding solute, will trigger the precipitation of the solid back out of the solution.

17. Saturation • To test for saturation, add solute • If the solute dissolves, the solution is unsaturated • If the solute falls to the bottom, the solution is saturated • If the solute falls to the bottom along with additional solute from the solution, the solution is supersaturated

18. Solubility Curves • A solubility curve is a graphical representation of solubility of multiple different substances • Solubility curves shows the solubility (in grams) of different solutes that can be dissolved in a specified amount of water (usually 100g or 100mL) at a given temperature • What happens to solubility of a substance as temperature increases? • Do any substances break this rule? • Because they are gases

19. Solubility Curves • On a solubility curve, the curves indicate the concentration of a saturated solution • Values on the graph below a curve represent unsaturated solutions • Values on the graph above a curve represent supersaturated solutions

20. Solubility Curves Which solute is most soluble at 10◦C? Which solute is least soluble at 40◦C? What mass of ammonia is dissolved in a saturated solution at 25◦C? If KNO3 has 70g of solute dissolved at 60◦C is the solution saturated, unsaturated, or supersaturated? If a solution of NaCl has 20g of solute dissolved at 90◦C is the solution saturated, unsaturated, or supersaturated?

21. Molarity • A standard solution or a stock solution is a solution whose concentration is accurately known. • There are multiple ways to measure concentration mathematically • Represented by M and units are M (molar) • Most commonly used expression of concentration

22. Molarity Examples • What is the molarity of a 23.0 mL solution that was made using 0.760 moles of magnesium sulfide? M = (0.760 moles) / (0.0230 L) = 33.0 M • How many moles of hydrochloric acid are dissolved in 5.55 L of a 0.250 M solution? (0.250 M) = moles / (5.55L) moles= 0.138 mol • What is the volume of a 0.88 M solution that contains 0.35 moles of potassium fluoride? (0.88 M) = (0.35 mole) / V V = 0.40 L

23. Molarity Examples • What mass of lithium hydroxide is dissolved in 700.0 mL of a 0.250 M solution? (0.250 M) = moles / (0.7000 L) moles= 0.175 moles = 4.19 g LiOH • What is the molarity of a 0.450 L solution that contains 100. g of sodium chloride? = 1.71 mole NaCl M = (1.71 mole) / (0.450 L) M =3.80 M

24. Measuring Concentration Experimentally • Use a reaction to create a precipitate and use mass of precipitate and stoichiometry to determine original concentration of solution • Use a spectrophotometer that measures concentration by shining a light through a solution and measuring how much of the light absorbed • Perform a titration experiment that uses a solution of known concentration to determine a solution of unknown concentration

25. Process of Making a Solution • If you are asked to make a solution with a certain molarity and a certain volume: For example 250.0 mL of a 0.35 M solution of sodium chloride • You will need to first determine the mass of solute you will need: • Why do we use mass instead of moles? • Procedure: • Measure amount of solute in grams (be specific) using a scale and add to a volumetric flask • Add water to the volumetric flask until it reaches the desired volume of solution (give a specific volume with units) • Is this exactly 250mL of water? • NO • Remember the volume is of SOLUTION not the volume of solvent • Mix solution • Procedure for the example above: • Measure 5.114 g of NaCl using a scale and add to a volumetric flask. • Add water to the volumetric flask until it reaches 250.0 mL • Mix solution

26. Dilution • If you made a solution of sweet tea (solute= sugar and solvent=tea), if it was not sweet enough, how could you increase the concentration? • If it was too sweet, how could you decrease the concentration? • Dilution is the process of adding solvent to a solution to lower the concentration of solution • The actual number of moles of the solute never changes • Only the amount of solvent changes to reduce the molarity of the solution.

27. Dilutions • Which solution is the most concentrated? • Which is the most dilute? • What happens to the volume as we dilute the solution? • What happens to the amount of solvent throughout the dilution? • What happens to the amount of solute throughout the dilution? • What happens to the concentration of the solution throughout the dilution?

28. Dilution Calculations • When you need to dilute a solution, you know 3 of the following 4 things: • the molarity of what you started with (M1), • the new molarity of the solution you want to make (M2) • The volume of the solution you start with (V1) • The volume of the new solution (V2) • Use the dilution equation M1V1= M2V2to solve for the last variable. • Volume can be in units of mL or L. • Watch out for the wording of the questions!

29. Dilution Calculations • Calculate the molarity that results when 250 mL of water is added to 125 mL of .251 M HCl. • First assign variables: • V1 = 125 mL • V2 = 125 mL + 250 mL = 375 mL • M1 = .251 mol/liter • M2 = ? • Then plug into formula and solve for unknown • M2 = M1V1/V2 = (.251 M)(125 mL)/375 mL • M2 = .0837 M

30. Dilution Practice • Suppose you wished to make a 0.879 L of 0.250 M aqueous silver nitrate by dilution a stock solution of a 0.675M aqueous silver nitrate. What will the volume of the stock solution would you need to use? • Calculate the new concentration when 50.0mL of water is added to 735mL of 1.25M NaCl. • How much water would need to be added to 750mL of a 2.8 M HCl solution to make a 1.0M solution?

31. Example #5 An aqueous solution is prepared by diluting 3.30 mL acetone solution (d = 0.789 g/mL) with water to a final volume of 75.0 mL. The density of the solution is 0.993 g/mL. What is the molarity, molality and mole fraction of acetone in this solution?

32. Procedure for Diluting • Performing a dilution is similar to making a solution • Measure specific volume of stock solution (be specific) using a graduated cylinder and place into a volumetric flask (V1) • Fill volumetric flask with water until it reaches final volume (V2) (be specific) • Mix solution

33. Dilutions • Describe the step-by-step process of diluting 25.0 mL of a 1.25 M KCl solution to a 0.750 M KCl solution. • (25.0 mL)(1.25 M) = V2(0.750M) • V2 = 41.7 mL • Measure 25.0 mL of 1.25 M KCl solution using a graduated cylinder and place into a volumetric flask • Fill volumetric flask with water until it reaches final volume of 41.7 mL • Mix solution

34. Stoichiometry with Molarity • Molarity (mol/L) can be used as a conversion factor in stoichiometry • Remember 22.4L/mol can ONLY be used for gases at STP • Molarity can be used for aqueous solutes • Use Molarity to convert from volume to moles • Mole ratio can be used to convert from one substance to another • Use conversion factor to convert to desired units

35. Example • Calcium hydroxide is sometimes used in water treatment plants to clarify water for residential use. Calculate the volume of 0.0250 M calcium hydroxide solution that can be completely reacted with 25.0 mL of 0.125 M aluminum sulfate solution. Al2(SO4)3(aq)+3Ca(OH)2(aq)2Al(OH)3(s)+3CaSO4(s) 0.375 L

36. Example • Ammonium sulfate is manufactured by reacting sulfuric acid with ammonia. What concentration of sulfuric acid is needed to react with 24.4 mL of a 2.20 Mammonia solution if 50.0 mL of sulfuric acid is used? H2SO4(aq) + 2NH3(aq)  (NH4)2SO4(aq)

37. Practice • H2SO4 reacts with NaOH, producing water and sodium sulfate. What volume of 2.0 M H2SO4 will be required to react completely with 75 mL of 0.50 MNaOH? • How many moles of Fe(OH)3 are produced when 85.0 L of iron(III) sulfate at a concentration of 0.600 Mreacts with excess NaOH? • What mass of precipitate will be produced from the reaction of 50.0 mL of 2.50 Msodium hydroxide with an excess of zinc chloride solution.

38. 0.50 mol NaOH 0.600 mol Fe2(SO4)3 1 L H2SO4 1 mol H2SO4 2 mol Fe(OH)3 x x x x x 1 L NaOH 1 L Fe2(SO4)3 2.0 mol H2SO4 1 mol Fe2(SO4)3 2 mol NaOH Answers # L H2SO4= 3. H2SO4(aq) + 2NaOH(aq)  2H2O + Na2SO4(aq) 0.075 LNaOH = 0.009375 L = 9.4 mL H2SO4 4. Fe2(SO4)3(aq)+6NaOH(aq)2Fe(OH)3(s)+3Na2SO4(aq) # mol Fe(OH)3= 85 L Fe2(SO4)3 = 102 mol Fe(OH)3

39. 2.50 mol NaOH 1 mol Zn(OH)2 99.40 g Zn(OH)2 x x x L NaOH 1 mol Zn(OH)2 2 mol NaOH Answers 5. 2NaOH(aq)+ZnCl2(aq)  Zn(OH)2(s) + 2NaCl(aq) # g Zn(OH)2= 0.0500 LNaOH = 6.21 g Zn(OH)2

40. Other ways of Measuring Concentration • There are multiple ways to measure concentration mathematically, below are some examples: • Percent by Mass: very similar to percent composition; nutrition labels • Percent by Volume: used for products such as hydrogen peroxide and alcohols • Mole Fraction: important for gases • Molality: important for colligative properties • Molarity: the most common way to express concentration

41. Measures of Concentration • Units of % by mass are % • Represented by Greek lower case chi (χ) and has NO UNITS • Represented by m and units are m (molal) • Water has a density of 1 g/mL • Represented by M and units are M (molar) • Most commonly used expression of concentration

42. Example #1 • M= • If 5.7 g KNO3 was dissolved in 233 mL of solution, what is the molarity of the solution? • Convert grams to mole of solute and mL to L 5.7 g KNO3 1 mole = .056 mol KNO3 233mL = 0.233L 101.10 g 0.056 mol KNO3 = 0.24 M KNO3 0.233 L

43. Example #2 62.1 g (1.00 mol) of ethylene glycol was dissolved in 250. g of H2O. Calculate molality, mole fraction, and % by mass of ethylene glycol.

44. Examples • How many grams of NaCl are needed to prepare 250 g of a 10.0% (by mass) NaCl solution? What is the mole fraction of NaCl in this solution? • What is the molality of a solution that contains 42.0 g of KCl in a 798 mL of water? • How many grams of NaOH are required to prepare 200mL of a 0.45M solution?

45. Colligative Properties • Van’t Hoff factor, (i):is the number of particles a solute breaks into when it dissolves. • Colligative property is a property that is dependent only on the number of solute particles present in solution (aka depends on the molality and van’t Hoff factor factor) • Examples of colligative properties are Freezing Point and Boiling Point • An increase or decrease in concentration will effect colligative properties • When concentration is increased • The freezing point is always lowered this is called freezing point depression • And the boiling point is always raised this is called boiling point elevation

46. Freezing Point Depression and Boiling Point Elevation • Why do we salt the roads when is snows? • Why do we add salt to the pot of boiling water when we cook noodles?

47. Boiling-Point Elevation and Freezing-Point Depression • van’t Hoff Factor, i:This factor equals the number of ions produced from each molecule of a compound upon dissolving. • i= 1 for covalent compounds • i= # of ions for ionic compounds • i = 1 for CH3OH i = 3 for CaCl2 • i = 2 for NaCli = 5 for Ca3(PO4)2

48. Colligative Properties • Which would have a lower freezing point: a 0.50M solution of NaCl water or a 1.0M solution of NaCl?Which would have a lower boiling point? Why? • Which would have a higher freezing point: a 0.50M solution of NaCl water or a 0.50M solution of CaCl2? Which would have a higher boiling point? Why?

49. Boiling point elevation • Boiling-Point Elevation (∆Tb):∆Tb= i .Kb. m • ∆Tb= change in boiling point • i = van’t Hoff factor • Kb= molal boiling-point elevation constant • m = molality Boiling Point of solution = normal boiling point of solvent + ΔTb What is the normal boiling point of water?

50. Freezing point depression • Freezing-Point Depression (∆Tf):∆Tf = i . Kf. m • ∆Tf= change in freezing point • i = van’t Hoff factor • Kf= molalfreezing-point depression constant • m = molality Freezing Point of solution = normal freezing point of solvent - ΔTf What is the normal freezing point of water?