1. Chapter 6: Sports Drink

2. Introductory Activity What do you think are the benefits of drinking a sports drink while exercising rather than plain water? How are your ideas influenced by the marketing strategies of the companies that sell these drinks?

3. Sports Drinks This chapter will introduce the chemistry needed to understand how Sports Drinks work – Section 6.1: Solutions & electrolytes – Section 6.2: Concentrations of solutions – Section 6.3: Acidity & pH – Section 6.4: Solubility & precipitates – Section 6.5: Stoichiometry – Section 6.7: Limiting Reactants – Section 6.6: Properties of solutions

4. Sports Drinks Solution Is a Electrolytes Solubility With that need to all dissolve when mixed together Concentrations How much solute is in it? pH Some affect Titrations Can be determined by Differ from pure liquids in Properties

5. Section 6.1—Solutions & Electrolytes What are those “electrolytes” they say you’re replacing by drinking sports drinks?

6. Dissolving substances Substances are dissolved by a process called hydration – The solvent and solute need to break intermolecular forces within themselves – New intermolecular forces are formed between the solvent and solute – The solvent “carries off” the solute particles

7. Dissolving Ionic Compounds -+ + + + - -- - ++ -- O H H - + water Ionic compound Water molecules are polar and they are attracted to the charges of the ions in an ionic compound. When the intermolecular forces are stronger between the water and the ion than the intramolecular between the ions, the water carries away the ion.

8. -+ + - -- - + + Dissolving Ionic Compounds ++ -- O H H - + water Ionic compound As more ions are “exposed” to the water after the outer ions were “carried off”, more ions can be “carried off” as well.

9. Dissolving Ionic Compounds ++ -- O H H - + water Ionic compound -+ + + + - -- - These free-floating ions in the solution allow electricity to be conducted

10. Electrolytes When there are free-floating charges in a solution then it can conduct electricity. Things that produce free-floating charges when dissolved in water are called electrolytes.

11. Dissolving Covalent Compounds -+-+ -+-+ -+-+ -+-+ -+-+ Solvent, water (polar) ++ -- -+-+ Solute, sugar (polar) Polar covalent molecules are formed in the same way—water forms intermolecular forces with the solute and “carries” the solute particles away.

12. Dissolving Covalent Compounds Solvent, water (polar) ++ -- -+-+ Solute, sugar (polar) -+-+ -+-+ -+-+ -+-+ -+-+ However, the polar covalent molecules themselves do not split into charged ions—the solute molecule stays together and just separates from other solute molecules.

13. Non-electrolytes When molecules separate from other molecules (breaking intermolecular forces), but free-floating charges are not produced from breaking intramolecular forces, the solution cannot conduct electricity. These are called non-electrolytes

14. Types of Electrolytes Strong Electrolytes Ionic compounds Weak Electrolytes Ionic Compounds Non-Electrolytes Covalent Compounds Almost all ions are separated when dissolved in water. Only a few ions are separated when dissolved in water No molecules separate—ions are not formed Easily conducts electricity when dissolved in water Conducts electricity slightly when dissolved in water Does not conduct electricity at all when dissolved in water

15. Breaking up Electrolytes Leave polyatomic ions in-tact (including the subscript within the polyatomic ion) All subscripts not within a polyatomic ion become coefficients Be sure to include charges on the dissociated ions! Example: Break up the following ionic compounds into their ions KNO 3 Ca(NO 3 ) 2 Na 2 CO 3

16. Misconceptions about dissolving People often describe something that dissolves as having “disappeared” Before the solute dissolves, it’s in such a large group of particles that we can see it. After dissolving, the solute particles are still there—they’re just spread out throughout the solution and are in groupings so small that our eyes can’t see them

17. Types of Solutions Unsaturated More solute can be dissolved Saturated No more solute can be dissolved—it’s “full” Super-Saturated Has more solute than would make a saturated solution dissolved In general, the higher the temperature of a solution, more solid can be dissolved.

18. Section 6.2—Concentration How do we indicate how much of the electrolytes are in the drink?

19. Concentrated versus Dilute solute solvent Lower concentration Not as many solute (what’s being dissolved) particles Higher concentration More solute (what’s being dissolved) particles

20. Concentration Concentration gives the ratio of amount dissolved to total amount There are several ways to show concentration

21. Percent Weight/Volume This is a method of showing concentration that is not used as often in chemistry However, it’s used often in the food and drink industry – For example, your diet drink can might say you have less than 0.035 g of salt in 240 mL. – That would give you a concentration of 0.035 g / 240 mL, which is 0.015% solution

22. %(W/V) Example Example: If you dissolve 12 g of sugar in 150 mL of water, what percent weight/volume is the solution?

23. Concentration using # of molecules When working with chemistry and molecules, it’s more convenient to have a concentration that represents the number of molecules of solute rather than the mass (since they all have different masses) Remember, we use moles as a way of counting molecules in large numbers

24. Quick Mole Review 1 mole = 6.02 × 10 23 molecules The molecular mass of a molecule is found by adding up all the atomic masses in the atom Molecular mass in grams = 1 mole of that molecule

25. Quick Mole Example Example: How many moles are in 25.5 g NaCl?

26. Molarity Molarity (M) is a concentration unit that uses moles of the solute instead of the mass of the solute

27. Molarity Example Example: If you dissolve 12 g of NaCl in 150 mL of water, what is the molarity?

28. Molarity Example Example: If you dissolve 12 g of NaCl in 150 mL of water, what is the molarity? 1.4 M NaCl 12 g NaCl = _______ mole NaCl g NaCl mole NaCl 1 58.44 0.21 1 mole NaCl molecules = 58.44 g Na Cl 1 1 22.99 g/mole 35.45 g/mole  = 22.99 g/mole = 35.45 g/mole + 58.44 g/mole  Remember to change mL to L! 150 mL of water = 0.150 L

29. Converting between the two If you know the %(W/V), you know the mass of the solute You can convert that mass into moles using molecular mass You can then use the moles solute to find molarity

30. Converting from % to M Example Example: What molarity is a 250 mL sample of 7.0 %(W/V) NaCl?

31. Concentration of Electrolytes An electrolyte breaks up into ions when dissolved in water You have to take into account how the compound breaks up to determine the concentration of the ions CaCl 2  Ca +2 + 2 Cl -1 For every 1 CaCl 2 unit that dissolves, you will produce 1 Ca +2 ion and 2 Cl -1 ions If the concentration of CaCl 2 is 0.25 M, the concentration of Ca +2 is 0.25 M and Cl -1 is 0.50 M

32. Let’s Practice #1 Example: You want to make 200 mL of a 15% (W/V) solution of sugar. What mass of sugar do you need to add to the water?

33. Let’s Practice #2 Example: What is the %(W/V) of a 500. mL sample of a 0.25 M CaCl 2 solution?

34. Let’s Practice #3 Example: What are the molarities of the ions made in a 0.75 M solution of Ca(NO 3 ) 2

35. Section 6.3—Acidity, pH How does concentration of acid affect the pH of a sports drink?

36. A Review of Acids & Bases

37. Acids – Arrhenius Definition Produce Hydronium ion (H 3 O +1 ) in water Hydronium ion is water + a hydrogen cation H O H water H +1 H O H H By this definition, if an acid is to give a H +1 to water, then all acids will have hydrogen as the cation (first element written).

38. How do Acids produce Hydronium? H O H H - water acid Hydrogen cation with some anion

39. How do Acids produce Hydronium? H O H H - +1

40. How do Acids produce Hydronium? H O H H +1 - Hydronium ion Anion

41. Bases – Arrhenius Definition Bases produce the hydroxide ion in water H O Hydroxide Ion

42. Characteristics of Acids & Bases BasesAcids Produce H 3 O +1 (hydronium ion) in water Produce OH -1 (hydroxide ion) in water Tastes sourTastes Bitter React with active metals to form hydrogen gas Feels slippery

43. Strength versus Concentration

44. Strong versus Weak Acids + + + --- Strong acid Most of the acid molecules have donated the H +1 to water How many hydronium ion – anion pairs can you find? How many intact acid molecules can you find? 3 1

45. Strong versus Weak Acids + - Weak acid Only a few of the acid molecules have donated the H +1 to water How many hydronium ion – anion pairs can you find? How many intact acid molecules can you find? 1 3

46. Concentrated versus Dilute solute solvent Lower concentration Not as many solute (what’s being dissolved) particles Higher concentration More solute (what’s being dissolved) particles

47. Combinations of Concentration & Strength DiluteConcentrated A lot of acid added & most dissociates Not much acid added, but most of what’s there dissociates A lot of acid added, but most stays together Not much acid added and most of what is there stays together Strong Weak

48. Acids and Bases as Electrolytes Acids and bases dissociate into ions in water Free-floating ions in water conduct electricity Acids & Bases are electrolytes Strong acids and bases are strong electrolytes Weak acids and bases are weak electrolytes

49. pH

50. Is a scale to measure the acidity of a sample pH Scale 114 Highly acidicVery basic (not acidic) neutral 7 Chapter 6 will give more detail about how pH is calculated!

51. pH is a Logarithmic Scale Logarithm –The number of times a base must be multiplied by itself to reach a given number # of multiples Base # you’re trying to reach

52. Calculating pH pH scale – Logarithmic scale of the acidity of a solution The pH scale uses base “10” pH has not units [ ] = concentration in Molarity

53. The “-” in the pH equation Because pH is the negative log of concentration of hydronium, as concentration increases, the pH goes down. The lowest pH is the highest concentration of hydronium

54. What does a “log” scale really mean? pH 4 3 2 1 10x more acidic 100x more acidic 1000x more acidic Level of acidity increases Every change of 1 in pH shows a change of 10x in concentration of hydronium

55. An example of calculating pH Example: Find the pH if the concentration of [H 3 O +1 ] is 0.25 M

56. An example of calculating hydronium Example: Find the [H 3 O +1 ] if the pH is 2.7

57. Auto-ionization of Water Water will split into ions – 2 H 2 O  H 3 O +1 + OH -1 Water will do this to make sure that at 25°C the following is true: – [H 3 O +1 ] × [OH -1 ] = 1 × 10 -14 So if you know the hydronium concentration at 25°C (which can be found from pH), then you can also find the hydroxide concentration

58. An example of calculating hydroxide Example: Find the [OH -1 ] if the pH is 10.7

59. Let’s Practice #1 Example: Find the pH if the concentration of [H 3 O +1 ] is 2.5 × 10 -5 M

60. Let’s Practice #2 Example: Find the [OH -1 ] if the pH is 3.6

61. Let’s Practice #3 Example: Find the [H 3 O +1 ] if the pH is 11.2

62. Section 6.4—Solubility & Precipitation How can we make sure everything that’s added to the sports drink will dissolve?

63. A Review of Double-Replacement Reactions

64. NaCl + AgNO 3  AgCl + NaNO 3 Double Replacement Reactions The cations from two compounds replace each other. Cl Na Ag O O O O N N O O Cl Ag Na O O O O N N O O Two ionic compounds switch ions

65. Double Replacement Reactions General format of a double replacement reaction:

66. Combine the cation of the first reactant with the anion of the second reactant CaCl 2 +AgNO 3 1 Products of a Double Replacement

67. Combine the cation of the second reactant with the anion of the first reactant CaCl 2 +AgNO 3 2 Products of a Double Replacement

68. & balance charges with subscripts when writing formulas Remember to write cations first … AgCl CaCl 2 +AgNO 3 3 Ca(NO 3 ) 2 + CaCl 2 AgNO 3 + Products of a Double Replacement Only leave subscripts that are in the original compound there if they are a part of a polyatomic ion!

69. Precipitation Reactions

70. A precipitation reaction is when 2 soluble substances are mixed together and they form an insoluble substance 2 soluble chemicals: NaOH and Cu(NO 3 ) 2

71. Precipitation Reactions A precipitation reaction is when 2 soluble substances are mixed together and they form an insoluble substance 2 soluble chemicals: NaOH and Cu(NO 3 ) 2 2 soluble chemical: NaNO 3 1 insoluble chemical (the precipitate): Cu(OH) 2

72. Why do some things dissolve and others don’t?

73. Remember the dissolving process? Substances are dissolved by a process called hydration – The solvent and solute need to break intermolecular forces within themselves – New intermolecular forces are formed between the solvent and solute – The solvent “carries off” the solute particles

74. Review--Dissolving Ionic Compounds -+ + + + - -- - ++ -- O H H - + water Ionic compound Water molecules are polar and they are attracted to the charges of the ions in an ionic compound. When the intermolecular forces are stronger between the water and the ion than the intramolecular between the ions, the water carries away the ion.

75. -+ + - -- - + + Review--Dissolving Ionic Compounds ++ -- O H H - + water Ionic compound As more ions are “exposed” to the water after the outer ions were “carried off”, more ions can be “carried off” as well.

76. Review--Dissolving Ionic Compounds ++ -- O H H - + water Ionic compound -+ + + + - -- - These free-floating ions in the solution allow electricity to be conducted

77. What about with stronger ionic bonds? 2-2+ 2- ++ -- O H H 2+ water Ionic compound Ion charge can affect strength of ionic bond—the higher the charges, the stronger the bond. (How closely the two ions can pack together also affects ionic bond strength)

78. ++ -- O H H 2- 2+ water Ionic compound If the connection between the water and the ions is not similar in strength or stronger than the ion-ion and water-water connections that are being broken… What about with stronger ionic bonds? 2-2+ 2-

79. What about with stronger ionic bonds? 2-2+ 2- If the connection between the water and the ions is not similar in strength or stronger than the ion-ion and water-water connections that are being broken… The water won’t be able to carry the ions away…it won’t dissolve the solid. ++ -- O H H 2- 2+ water Ionic compound

80. Solubility Rules

81. Solubility Rules Table THESE ANIONS FORMS SOLUBLE COMPOUNDS WITH THESE CATIONS FORM INSOLUBLE COMPOUNDS WITH THESE CATIONS NO 3 - nitrateMost cationsNo common cations CH 3 COO – acetate Most cationsAg + Cl - chlorideMost cationsAg +, Pb 2+, Hg 2 2+, Tl + Br - bromideMost cationsAg +, Pb 2+, Hg 2 2+, Tl + I - iodideMost cationsAg +, Pb 2+, Hg 2 2+, Tl + SO 4 2- sulfateMost cationsBa 2+, Sr 2+, Pb 2+, Ag +, Ca 2+ CrO 4 2- chromate Most cationsBa 2+, Sr 2+, Pb 2+, Ag + S 2- sulfideNH 4 +, cations of column 1, cations of column 2 Most other cations OH - hydroxide NH 4 +, cations of column 1, and Ba 2+ and Sr 2+ Most other cations CO 3 2- carbonate NH 4 +, cations of column 1 except Li + Most other cations PO 4 3- phosphate NH 4 +, cations of column 1 except Li + Most other cations This table, found at the end of Chpt 6 and in the Appendix, can help you figure out which compounds dissolve (those that are soluble) and which form precipitate (insoluble)

82. Let’s Practice #1 Example: Decide whether each is soluble or not NaNO 3 AgCH 3 COO CaBr 2 Ba(OH) 2 Cu(OH) 2

83. Let’s Practice #2 Example: Write the products for this reaction Na 2 CrO 4 (aq) + BaCl 2 (aq)  Remember to indicate compounds that dissolve with “aq” for “aqueous” and compounds that don’t dissolve with “s” for “solid”

84. Let’s Practice #3 Example: Write the products for this reaction NaCH 3 COO (aq) + KCl (aq)  Remember to indicate compounds that dissolve with “aq” for “aqueous” and compounds that don’t dissolve with “s” for “solid”

85. Section 6.5—Stoichiometry How can we determine in a lab the concentration of electrolytes?

86. 2 H 2 + O 2  2 H 2 O 2 No coefficient = 1 2 For every 2 moles of H 2 … 1 mole of O 2 is need to react… and 2 moles of H 2 O are produced What do those coefficients really mean?

87. What is stoichiometry? Stoichiometry – Using the mole ratio from the balanced equation and information about one compound in the reaction to determine information about another compound in the equation.

88. Stoichiometry with Moles Example: If 4.2 mole of H 2 reacts completely with O 2, how many moles of O 2 are needed? 2 H 2 + O 2  2 H 2 O

89. Stoichiometry with Moles 4.2 mole H 2 mole H 2 mole O 2 = ________ mole O 2 2 1 2.1 From balanced equation: 2 mole H 2  1 mole O 2 Example: If 4.2 mole of H 2 reacts completely with O 2, how many moles of O 2 are needed? 2 H 2 + O 2  2 H 2 O

90. But we can’t measure moles in lab! We can’t go to the lab and count or measure moles…so we need a way to work in measurable units, such as grams and liters! Molecular mass gives the grams = 1 mole of a compound!

91. Stoichiometry with Moles & Mass Example: How many grams of AgCl will be precipitated if 0.45 mole AgNO 3 is reacted as follows: 2 AgNO 3 + CaCl 2  2 AgCl + Ca(NO 3 ) 2

92. Stoichiometry with Mass Example: How many grams Ba(OH) 2 are precipitated from 14.5 g of NaOH in the following reaction: 2 NaOH + BaCl 2  Ba(OH) 2 + 2 NaCl

93. But what about for solutions? Molarity gives the number of moles of the solute that are in 1 L of a solution

94. Stoichiometry with Solutions Example: If you need 15.7 g Ba(OH) 2 to precipitate, how many liters of 2.5 M NaOH solution is needed? 2 NaOH + BaCl 2  Ba(OH) 2 + 2 NaCl

95. From balanced equation: 2 mole NaOH  1 mole Ba(OH) 2 Stoichiometry with Solutions 15.7 g Ba(OH) 2 g Ba(OH) 2 mole Ba(OH) 2 = ________ L NaOH 171.35 1 0.0733 Concentration of NaOH: 2.5 mole NaOH = 1 L mole Ba(OH) 2 mole NaOH 1 2 L NaOH 2.5 1 Molar Mass of Ba(OH) 2 : 1 mole Ba(OH) 2 = 171.35 g Example: If you need 15.7 g Ba(OH) 2 to precipitate, how many liters of 2.5 M NaOH solution is needed? 2 NaOH + BaCl 2  Ba(OH) 2 + 2 NaCl

96. What about gases? Standard Temperature and Pressure (STP) – 1 atm (or 101.3 kPa) and 273 K (0°C) Molar Volume of a Gas – at STP, 1 mole of any gas = 22.4 liters

97. Stoichiometry with Gases Example: If you need react 1.5 g of zinc completely, what volume of gas will be produced at STP? 2 HCl (aq) + Zn (s)  ZnCl 2 (aq) + H 2 (g)

98. Keeping all these equalities straight! TO GO BETWEENUSE THE EQUALITY Grams & molesMolecular Mass in grams = 1 mole moles & liters of a solutionMolarity in moles = 1 L Moles & liters of a gas at STP 1 mole = 22.4 L at STP 2 different chemicals in a reaction Coefficient ratio from balanced equation

99. Titrations—Using Stoichiometry Titration – Addition of a known volume of a known concentration solution to a known volume of unknown concentration solution to determine the concentration.

100. End Point End Point (or Stoichiometric Point) – When there is no reactant left over—they have all be reacted and the solution contains only products Indicators – Paper or liquid that change color based on pH level. The end point must be reached in order to use stoichiometry to calculate the unknown solution concentration If the pH of the products is known, the indicator can be chosen to indicate the end point

101. Gravemetrics—Using Stoichiometry Gravemetric Analysis – Using a reaction to precipitate out an insoluble compound. The solid is dried and massed. Stoichiometry can then be used to determine the original substance’s concentration from the mass of the precipitate

102. Let’s Practice #1 Example: If you are making 0.57 moles H 2 O, how many moles of O 2 are needed? 2 H 2 + O 2  2 H 2 O

103. Let’s Practice #2 Example: If you need to precipitate 10.7 g of Ba(OH) 2, how many grams NaOH are needed? 2 NaOH + BaCl 2  Ba(OH) 2 + 2 NaCl

104. Let’s Practice #3 Example: How many moles AgNO 3 are needed to react with 10.7 g CaCl 2 ? 2 AgNO 3 + CaCl 2  2 AgCl + 2 Ca(NO 3 ) 2

105. Let’s Practice #4 Example: How many liters of 0.10 M NaOH is needed to react with 0.125 L of 0.25 M BaCl 2 ? 2 NaOH + BaCl 2  Ba(OH) 2 + 2 NaCl

106. Let’s Practice #5 Example: If you produce 15.4 L of H 2 at STP, how many grams of ZnCl 2 is also produced? 2 HCl (aq) + Zn (s)  ZnCl 2 (aq) + H 2 (g)

107. Section 6.6—Limiting Reactants What happens if you don’t add reactants in a molar ratio?

108. Planning a Meal You go to the grocery store and you buy 1 package of Brats (5 Brats), 1 package of cheese (16 slices) and 1 package of hot dog buns (8 buns). If you use all of these…You can make this many… 5 Brats 5 meals 16 slices of cheese 8 hot dog buns 16 meals 8 meals So you have the possibility of making 5, 16 or 8 meals…which is it? You’ll never get the chance to make 8 or 16 meals…you’ll run out of Brats after 5. Once you run out of one component, you have to stop making meals.

109. What’s a limiting reactant? Limiting Reactant – The reactant that runs out and causes the reaction to stop. In the previous example, the Brats were the limiting reactant—once they were gone, you had to stop! Once even one of the reactants runs out, the reaction stops…it can’t make any more product.

110. Limiting Reactant Example Example: How many moles of H 2 O is produced when 2.3 moles O 2 and 2.3 moles H 2 react? 2 H 2 + O 2  2 H 2 O

111. Let’s Practice Example: If you react 10.5 g of NaOH and 7.5 g of BaCl 2, how many grams NaCl is produced? 2 NaOH + BaCl 2  Ba(OH) 2 + 2 NaCl

112. Section 6.7—Properties of Solutions How do all those dissolved things affect the properties of the drink?

113. What’s Vapor Pressure? Vapor Pressure – Pressure created above a sample by particles evaporating from the sample and becoming gas particles.

114. To evaporate, molecules must break intermolecular forces—this requires a minimum amount of energy As temperature increases, the average energy of the molecules increase More molecules will have the minimum needed to evaporate from the liquid As temperature increases, the vapor pressure increases. Vapor Pressure & Temperature Temperature is proportional to the average kinetic energy of the molecules. Average means some will have more and some will have less!

115. Vapor Pressure of Solutions Only solvent particles on the very top layer of the sample can evaporate Looking down on the top of beaker: Solvent particles Beaker with solvent only If a solvent particle on the top layer has enough energy, it can break the IMF’s and evaporate Once evaporated, they cause vapor pressure

116. Vapor Pressure of Solutions Only solvent particles on the very top layer of the sample can evaporate Looking down on the top of a solution in a beaker: Beaker with solvent only Solvent particles Solute particles The solvent and solute form intermolecular forces (connections) with each other. The solvent must now break those connections in order to evaporate. The connections are holding the solvent particles back.

117. Vapor Pressure of Solutions Only solvent particles on the very top layer of the sample can evaporate Looking down on the top of a solution in a beaker: Beaker with solvent only The fewer particles that evaporate, the lower the vapor pressure. The vapor pressure of a solution is always less than the pure solvent.

118. When does something boil? Heat source (usually underneath) heats the molecules closest to it the fastest Molecules are gaining the energy to break intermolecular forces and become a gas Atmospheric pressure pushes down on the top of the liquid

119. When does something boil? When enough water molecules turn to gas and create as much pressure as the atmosphere is pushing down with, a bubble can form (counter-act the atmospheric pressure)

120. Boiling and Atmospheric Pressure Boiling occurs when vapor pressure of liquid = atmospheric pressure Higher altitude means lower atmospheric pressure The vapor pressure of the liquid doesn’t need to be as high to boil with lower atmospheric pressure The bubbles can form at a lower temperature The boiling point of a liquid is always lower at higher altitude

121. Boiling Points of Solutions Boiling occurs when vapor pressure of liquid = atmospheric pressure Solutions have lower vapor pressure than the pure solvent. The solution does not have high enough vapor pressure to boil at the solvent’s boiling point The temperature needs to be raised until the vapor pressure of the solution = atmospheric pressure The boiling point of a solution is always higher than the pure solvent

122. When do things freeze? When you’re above the freezing point, solid will melt to liquid When you’re below the freezing point, liquid will freeze to solid Freezing point is when there is equilibrium between solid & liquid—the amount of solid and liquid stay the same This occurs when the rate of evaporation from the solid is the same as the rate of evaporation from the liquid Every time a molecule evaporates from the solid, one also evaporates from the liquid. Every time a molecule re-forms into the solid, one also reforms into the liquid. Neither one can “get ahead”—it’s at equilibrium

123. In order for a liquid to freeze, the solid’s vapor pressure and the liquid’s vapor pressure must be equal The solid is the pure solvent. The liquid is the solution. The vapor pressure of the liquid (solution) is lower than the solid’s (solvent) The temperature is lowered until the solid’s vapor pressure = the liquid solution’s vapor pressure The freezing point of a solution is always lower than the pure solvent Freezing Points of Solutions This is the point where the speed of molecules joining to form a solid equals the speed molecules leave the solid to be liquid

124. When solutes are electrolytes, the impact is greater For every 1 mole of ___ added __ moles of particles are in solution Sugar (non-electrolyte)1 (C 6 H 12 O 6 stays together) NaCl (electrolyte) CaCl 2 2 (Na + + Cl - ) 3 (Ca 2+ + 2 Cl - ) Electrolytes break up into more than one particle when added to water. Therefore, there are even more particles when considering colligative properties. What effects do electrolytes cause?

125. What did you learn about sports drinks?

126. Sports Drink Solution Is a Electrolytes Solubility With that need to all dissolve when mixed together Concentrations How much solute is in it? pH Some affect Titrations Can be determined by Differ from pure liquids in Properties