Slide 2 / 142. Properties of Matter and Solutions

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1 Slide 1 / 142

2 Slide 2 / 142 Properties of Matter and Solutions

3 Slide 3 / 142 Properties of Matter and Solutions Pyrite, otherwise known as "fools gold" has fooled many a tourist over the years. Physical and chemical properties such as density or reactivity help us identify what substances are made of.

4 Slide 4 / 142 Matter We define matter as anything that has mass and takes up space. Atoms of an element Molecules of a compound molecules of a diatomic element Mixture of elements and a compound

5 Slide 5 / 142 What is Matter Made of? Elements and Compounds Substances that could not be broken down by any physical or chemical method were/are called elements Substances that could be broken down into different elements using physical or chemical methods were/are called compounds Element Ne(g) Ca(s) Au(s) Hg(l) Compound CO 2 (g) CaCO 3 (s) AuNO 3 (s) HgI(s)

6 Slide 6 / 142 Elements Elements are found on the periodic table. Mg-Magnesium Cu-Copper C-carbon diamond and graphite Al Aluminum foil I-Iodine vapor Na Sodium

7 Slide 7 / 142 Compounds Compounds are formed by combinations of different types of elements. CAFFEINE

8 Slide 8 / Which of the following would NOT be a compound? A HCl B CS 2 C H 2 O D CH 4 E I 2

9 Slide 9 / Which of the following is FALSE regarding compounds? A They consist of more than one element combined B C A compound has a set of properties distinct from the individual elements from which it is made When a compound is separated into its elements, the elements will have the same properties of the compound D Br 2 would not be considered a compound E NaCl would be considered a compound

10 Slide 10 / 142 Law of Definite Composition When electricity is passed through water (a compound), hydrogen and oxygen gas are produced. electricity liquid water > hydrogen gas + oxygen gas 100 grams 11.2 grams 88.8 grams When the amounts of gases produced are analyzed, no matter where the water came from or how large the sample, water always consists of exactly 11.2% hydrogen and 88.8% oxygen by mass.

11 Slide 11 / 142 Law of Definite Composition In fact, each compound had it's own definite composition by mass. Substance % carbon by mass % oxygen by mass carbon dioxide 27.3% 72.7% carbon monoxide 42.8% 57.1% This principle, that a certain substance will have it's own unique set composition of elements, is known as the Law of Definite Composition.

12 Slide 12 / 142 Pure Substances vs. Mixtures Some matter can be separated by heat, filtering, or boiling into other substances but did NOT obey the law of definite composition. These substances are known as mixtures and are NOT pure substances. More on mixtures later! Pure Substance Definitive Composition Examples: gold (Au) Mixture Non-definitive composition Examples: steel (Fe, C, Mn, Cr,...) pure water (H 2 O) salt water (H 2 O, Cl-, Na +,...)

13 Slide 13 / A sample of material A is collected in Nevada and found to consist of 94% oxygen and 6% hydrogen by mass. Another sample of material A is collected in Maine and found to contain 94% oxygen and 6% hydrogen. What kind of substance is this? A Element B Compound C Mixture D B and C E A, B, and C

14 Slide 14 / A sample of a material is found to contain 56% oxygen, 32% iron, and 12% sulfur. When another sample of the same material is collected, the composition was 44% oxygen, 30% iron, and 25% sulfur. What kind of substance is this? A element B compound C mixture D pure substance E B and D

15 Slide 15 / 142 Properties of Matter It was clear, even to the ancients, that not all matter shares the same characteristics/properties. gold Substance Property lustrous, soft metal, non-reactive, solid at room temperature salt water transparent, liquid at room temperature, could be separated by heat, no definite composition pure water calcium carbonate transparent, liquid at room temperature, definite composition, could be separated by electrolysis solid at room temperature, high melting point, non-lustrous, could be separated by heat

16 Slide 16 / 142 Physical Properties of Matter A physical property is a characteristic that can be observed WITHOUT altering the identity of the material. Physical Properties of water water melts at 0 Celsius at standard pressure water is transparent water has a density of roughly 1 g/ml at 25 C water is not soluble (does not dissolve) in gasoline water is colorless Notice all of these properties can be observed without changing the identity of the water - it is still water!

17 Slide 17 / 142 Physical Properties of Matter Who doesn't like brick oven pizza! A brick used in an oven is made of a mixture of aluminum oxide and silicon oxide. Think of as many physical properties of a brick that you can. Feel free to use terms like high and low if you don't know an exact number. high density high move melting for answer point reddish color brittle (break not bend)

18 Slide 18 / Which of the following IS NOT a physical property? A copper has a reddish gold color B iron reacts with oxygen to form rust C table salt dissolves easily in water D silver is an excellent conductor of electricity E all of theses are physical properties

19 Slide 19 / Which of the following IS a physical property? A acetone has a density of 0.87 g/ml B aluminum will burn in air to make aluminum oxide C water can undergo electrolysis and produce hydrogen and oxygen gas D Both A and C E Both B and C

20 Slide 20 / 142 Physical and Chemical Changes Chemical Changes Physical Changes Chemical changes result in new substances. Includes combustion, oxidation, decomposition, etc. Changes in matter that don't change the composition of a substance. Includes changes of state, temperature, volume, etc.

21 Slide 21 / 142 Chemical Properties These properties can only be observed when we attempt to change the identity of the material. There are a few tell tale signs that a chemical change has taken place: Color change Emission of Light Precipitate formation Production of gas

22 Slide 22 / 142 Chemical Properties Color change - marshmallow burning Emission of Light - wood burning

23 Slide 23 / 142 Chemical Properties Precipitate formation - solid forming from liquid mixtures + Production of gas - when limestone is heated heat +

24 Slide 24 / 142 Chemical Properties Class Discussion Compare the chemical properties of a pepperoni pizza with that of the brick oven. The pizza will react with the oxygen in the air and burn. The move for answer brick will not burn in the air.

25 Slide 25 / Which of the following is NOT a chemical property? A Silver tarnishing into silver oxide B gasoline burning in air C candle wax burning D candle wax melting E iron rusting

26 Slide 26 / All of the following are physical properties except.? A Gold's low reactivity with oxygen B Gasoline's inability to dissolve in water C Water melting at 0 C D Hot knife cutting through ice cream cake E evaporating water away from salt water

27 Slide 27 / In the following list, only is not an example of a chemical change. A B C D E dissolution of a penny in nitric acid the condensation of water vapor a burning candle the formation of polyethylene from ethylene the rusting of iron

28 Slide 28 / Which of the following are chemical changes? 1. rusting of a nail 2. freezing of water 3. decomposition of water into hydrogen and oxygen gases 4. compression of oxygen gas A 2, 3, 4 B 1, 3, 4 C 1, 3 D 1, 2 E 1, 4

29 Slide 29 / 142 Properties of Matter Application When you cook, cheese can be melted or it can be burned. One is a chemical change, the other a physical change. Explain which is which and how you knew! melted burned Melting is a physical change because the cheese has not changed - we know this because we see no evidence of a chemical change (no gas, light, precipitate, color change). move for answer However, burning cheese is a chemical change because we clearly see a color change, taste change, production of a gas when you set off the smoke detector!

30 Slide 30 / 142 Extensive Properties of Matter These are properties in which the value depends on how much of the material is present. Examples The mass of a glass of water is 30 grams. The stick has a length of 12.2 meters The helium balloon has a volume of 14.7 liters

31 Slide 31 / 142 Intensive Properties of Matter These are properties in which the value is independent of the amount of material. Examples The water is transparent and colorless The melting point of an iron chunk is 1538 Celsius The specific heat (amount of energy required to raise 1 gram by 1 degree celsius) of aluminum is 0.89 J/g*C

32 Slide 32 / 142 Intensive Physical Properties Density is an excellent example of an intensive property. No matter the size of the sample, the ratio of the mass to the volume for a given substance is the same. The higher the volume of the sample, the higher the mass will be. mass of water volume of water density of water grams ml g/ml grams ml g/ml grams ml g/ml note that the differences in density are the result of this being actual experimental data!

33 Slide 33 / 142 Application and Class Discussion Some meteorites found on the earth's surface are made of solid metal like iron. What kind of property - intensive or extensive - do you think would be most useful in identifying the metal in the meteorite? Explain. Intensive properties are unique to each substance so they are better for identifying. You can have 10 grams of just about anything or 5 ml of of just move about for answer anything, but only iron has a density of exactly 7.78 g/ml

34 Slide 34 / Which of the following would be an intensive physical property? A The color of the liquid bromine is reddish brown B The mass of the iron pipe is grams C The aluminum block engine has a density of 2.7 g/ ml D Both A and B E Both A and C

35 Slide 35 / Tungsten is a substance with an extremely high melting point and is used in light bulb filaments. Which of the following would be an extensive property of tungsten? A Tungsten melts at 3422 C B Tungsten has a silver color C Tungsten has a specific heat of J/gC D A tungsten filament is 10 cm long E A tungsten block will have a density of 15.6 g/ml

36 Slide 36 / Of the following, only is an extensive property. A B C D E density mass boiling point freezing point temperature

37 Slide 37 / Which one of the following is not an intensive property? A B C D E density mass boiling point freezing point temperature

38 Slide 38 / Which one of the following is an intensive property? A B C D E density mass boiling point freezing point temperature

39 Slide 39 / 142 Properties of Matter Summary Physical observed without changing identity of substance melting point, density, color, solubility, hardness, etc. Chemical observed by changing identity of substance reactivity with other substances Intensive independent of sample size color, melting point, density, etc. Extensive dependent on sample size mass, length, volume, etc.

40 Slide 40 / 142 Classification of Matter Earlier in the unit, we discussed that matter was either a pure substance or a mixture based on whether the composition was definite or variable. Matter Pure Substance Mixture definite composition variable composition

41 Slide 41 / 142 Mixtures Mixtures are a combination of two or more substances that vary can in composition. A classic example of a mixture would be salt water. Salt water can var in it's "saltiness" which makes it a mixture and not a pure substance. For example, the Mediterranean sea is roughly 5% more salty around Greece than it is off the coast of Spain. Mixtures can be separated into pure substances by physical means such as heating. Desalinization factories heat salt water to evaporate the water and leave the salt behind.

42 Slide 42 / 142 Types of Mixtures Heterogeneous mixtures are different throughout. For instance, a raisin muffin, a chocolate chip cookie are heterogeneous. But so is sand on the beach, since you can see differences in the sand due to grain size, etc. Homogeneous mixtures are the same throughout. These are also called solutions. Tap water and the air you breathe are excellent examples of solutions.

43 Slide 43 / 142 Solutions Solutions are defined as homogeneous mixtures of two or more pure substances. The solvent is the substance present in the greatest abundance. All other substances are solutes. Solvent dissolves the solute.

44 Slide 44 / 142 Credit totom Greebowe

45 Slide 45 / A combination of sand, salt, and water is an example of a. A B C D E homogeneous mixture heterogeneous mixture compound pure substance solid

46 Slide 46 / If matter is uniform throughout and cannot be separated into other substances by physical processes, but can be decomposed into other substances by chemical processes, it is called a (an). A B C D E heterogeneous mixture element homogeneous mixture compound mixture of elements

47 Slide 47 / Homogeneous mixtures are also known as. A B C D E solids compounds elements substances solutions

48 Slide 48 / 142 Dissociation _ _ When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them. _ 2+ _ 2+ _ 2+ _ This process is called dissociation. 2+ _ 2+ _

49 Slide 49 / 142 Electrolytes and Nonelectrolytes An electrolyte is a substances that dissociates into ions when dissolved in water. A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

50 Slide 50 / 142 Electrolytes and Nonelectrolytes Soluble ionic compounds tend to be electrolytes. Molecular compounds tend to be nonelectrolytes, except for acids and bases. Strong Weak Nonelectrolyte electrolyte electrolyte Ionic All None None Molecular strong acids weak acids All other weak bases compounds

51 Slide 51 / 142 Electrolytes A strong electrolyte dissociates completely when dissolved in water. A weak electrolyte only dissociates partially when dissolved in water.

52 Slide 52 / 142 Electrolytes The following are examples of chemicals that are strong and weak electrolytes. Nonelectrolytes do not dissociate in water. Strong Electrolyte Weak Electrolyte Nonelectrolyte HCl HNO 3 HClO 4 H 2 SO 4 NaOH Ba(OH) 2 Ionic Compounds CH 2 COOH HF HNO 2 NH 3 H 2 O (NH 2 ) 2 CO (urea) CH 3 OH (methanol) C 2 H 5 OH (ethanol) C 6 H 12 O 6 (glucose) C 12 H 22 O 11 (sucrose)

53 Slide 53 / A strong electrolyte is one that completely in solution. A reacts B associates C disappears D ionizes(dissociates) E solidifies

54 Slide 54 / A weak electrolyte exists predominantly as in solution. A B C D E atoms ions molecules electrons an isotope

55 Slide 55 / Which of the following would make the most effective electrolyte when dissolved in water? A CO 2 (g) B NaCl(s) C C 6 H 12 O 6 (s) D C(s) E N 2 (g)

56 Slide 56 / Which of the following would make the LEAST effective electrolyte when dissolved in water? A C 2 H 5 OH(l) B LiBr(s) C NaNO 3 (s) D MgCl 2 (s) E All are effective electrolytes

57 Slide 57 / 142 Solutions The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles. 2+ _ 2+ _

58 Slide 58 / 142 How Does a Solution Form? As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them solute + + solvent water The solute is added to the solvent The negative ions are pulled away by the positive pole of the solvent molecule The positive ions are pulled away by the negative pole of the solvent molecule

59 Slide 59 / 142 How Does a Solution Form? If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

60 Slide 60 / The process of solute particles being surrounded by solvent particles is known as. A B C D E salutation agglomeration solvation agglutination dehydration

61 * Three processes affect the energetics of solution: Slide 61 / 142 Energy Changes in Solution separation of solute particles separation of solvent particles H1- Separation of solute molecules new interactions between solute and solvent The heat content of a system includes the internal energy of a system and the pressure and temperature and is referred to as H. H 2 - Separation of solvent molecules + H 3 - Formation of solute-solvent interactions

62 * Slide 62 / 142 Energy Changes in Solution Enthalpy Net exothermic process Separated Separated solvent + solute particles particles Separated Solvent + solute particles ΔH 1 #H 2 Solvent + Solute ΔH solution Solution ΔH 3 The enthalpy change of the overall process depends on H for each of these steps. Solution can occur when the process is endothermic or exothermic. When heat is released or when it is pulled in from the surroundings. Why?

63 Slide 63 / 142 * Gibbs Free Energy Reactions, including solution, will occur spontaneously as long as the change in Gibbs Free Energy is negative. When the process, is endothermic (heat is taken in from the surroundings), the increase in enthalpy is offset by an increase in entropy. Separated Separated solvent + solute particles particles Separated Solvent + solute particles ΔH 1 ΔH 2 ΔH solution Solvent + Solute Solution ΔH 3 Net endothermic process

64 Slide 64 / 142 Solutions Just because a substance disappears when it comes in contact with a solvent, it doesn t mean the substance dissolved. Dissolution is a physical change you can get back the original solute by evaporating the solvent. If you can t, the substance didn t dissolve, it reacted.

65 Slide 65 / 142 Saturated Solutions + _ + + _ + + _ _ + _ + In a saturated solution, the solvent holds as much solute as is possible at that temperature. Dissolved solute is in dynamic equilibrium with solid solute particles.

66 Slide 66 / 142 Unsaturated Solutions In an unsaturated solution, there is less solute dissolved in the solvent at that temperature. Solid solute is not in dynamic equilibrium with dissolved solute

67 Slide 67 / 142 Supersaturated Solutions In supersaturated solutions, the solvent holds more solute than is normally possible at that temperature. These solutions are unstable; crystallization can usually be stimulated by adding a seed crystal or scratching the side of the flask. Click here for a video on Rapid Crystallization

68 Slide 68 / A saturated solution. A B C D E contains as much solvent as it can hold contains no double bonds contains dissolved solute in equilibrium with undissolved solute will rapidly precipitate if a seed crystal is added cannot be attained

69 Slide 69 / An unsaturated solution is one that. A B C D E has no double bonds contains the maximum amount of solute possible, and is in equilibrium with undissolved solute has less solute dissolved than the maximum solubility at that temperature contains more dissolved solute than the solubility allows contains no solute

70 Slide 70 / A solution with a concentration higher than the solubility is. A B C D E is not possible is unsaturated is supercritical is saturated is supersaturated

71 Slide 71 / A supersaturated solution. A B C D E is one with more than one solute is one that has been heated is one with more amount of solute than its solubility must be in contact with undissolved solids exists only in theory and cannot actually be prepared

72 Slide 72 / 142 Factors Affecting Solubility Chemists use the axiom like dissolves like." Polar substances tend to dissolve in polar solvents. Nonpolar substances tend to dissolve in nonpolar solvents. Alcohol CH 3 OH methanol CH 3 CH 2 OH ethanol CH 3 CH 2 CH 2 OH propanol CH 3 CH 2 CH 2 CH 2 OH butanol CH 3 CH 2 CH 2 CH 2 CH 2 OH pentanol CH 3 CH 2 CH 2 CH 2 CH 2 CH 2 OH hexanol Solubity in water Solubility in hexane # 0.12 # # # # 0.11 # # # solubility expressed in mol/100g solvent # = completely miscible

73 Slide 73 / 142 Factors Affecting Solubility The more similar the intermolecular attractions, the more likely one substance is to be soluble in another. Glucose (which has hydrogen bonding) is very soluble in water, while cyclobutane (which only has dispersion forces) is not. Hydrogen bonding sites Glucose- has hydroxyl groups and is highly soluble in water Cyclobutane-has no polar OH groups and is essentially insoluble in water

74 Slide 74 / 142 Factors Affecting Solubility Vitamin A Vitamin C soluble in nonpolar compounds (like fats) soluble in water

75 Slide 75 / The phrase "like dissolves like" refers to the fact that. A B C D E gases can only dissolve other gases polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes solvents can only dissolve solutes of similar molar mass condensed phases can only dissolve other condensed phases polar solvents dissolve nonpolar solutes and vice versa

76 Slide 76 / Which one of the following is most soluble in water? A B C D E CH 3 OH CH 3 CH 2 CH 2 OH CH 3 CH 2 OH CH 3 CH 2 CH 2 CH 2 OH CH 3 CH 2 CH 2 CH 2 CH 2 OH

77 Slide 77 / Which one of the following is most soluble in hexane (C 6 H 14 )? A B C D E CH 3 OH CH 3 CH 2 CH 2 OH CH 3 CH 2 OH CH 3 CH 2 CH 2 CH 2 OH CH 3 CH 2 CH 2 CH 2 CH 2 OH

78 Slide 78 / Which of the following substances is more likely to dissolve in CH 3 OH? A CCl 4 B Kr C N 2 D CH 3 CH 2 OH E H 2

79 Slide 79 / Which of the following substances is more likely to dissolve in water? A HOCH 2 CH 2 OH B CHCl 3 C D CH 3 (CH 2 ) 9 HCO CH 3 (CH 2 ) 8 CH 2 OH E CCl 4

80 Slide 80 / Which one of the following substances is more likely to dissolve in CCl 4? A CBr 4 B HBr C HCl D CH 3 CH 2 OH E NaCl

81 Slide 81 / 142 Temperature and Solubility A solubility chart can be used to determine the amount of solute that can be dissolved by a particular solvent at a range of temperatures. The line of a solubility chart represents a saturated solution. A point above the line represents a supersaturated solution at that temperature.

82 Slide 82 / 142 Temperature and Solubility A point above the line represents a supersaturated solution at a specific temperature. The line of a solubility chart represents a saturated solution.

83 Slide 83 / The point on the graph represents a solution that is: A B C D Unsaturated Saturated Supersaturated Cannot be Determined

84 Slide 84 / The point on the graph represents a solution that is: A B C D Unsaturated Saturated Supersaturated Cannot be Determined

85 Slide 85 / The point on the graph represents a solution that is: A B C D Unsaturated Saturated Supersaturated Cannot be Determined

86 Slide 86 / The point on the graph represents a solution that is: A B C D Unsaturated Saturated Supersaturated Cannot be Determined

87 Slide 87 / The point on the graph represents a solution that is: A B C D Unsaturated Saturated Supersaturated Cannot be Determined

88 Slide 88 / The point on the graph represents a solution that is: A B C D Unsaturated Saturated Supersaturated Cannot be Determined

89 Slide 89 / The change in concentration show on the graph below is most likely due to (assume there is no phase change and the amount of water remains constant) A More solute being added to the solution at constant temperature B C D No extra solute added and the solution being cooled The solution heated, more solute added, then the solution is cooled None of the above

90 Slide 90 / The change in concentration show on the graph below is most likely due to (assume there is no phase change and the amount of water remains constant) A More solute being added to the solution at constant temperature B C D No extra solute added and the solution being cooled The solution heated, more solute added, then the solution is cooled None of the above

91 Slide 91 / The change in concentration shown on the graph below is most likely due to (assume there is no phase change and the amount of water remains constant) A B C D More solute being added to the solution at constant temperature No extra solute added and the solution being cooled The solution heated, more solute added, then the solution is cooled None of the above

92 Slide 92 / 142 Temperature and Solubility Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

93 Slide 93 / 142 Temperature and Solubility of gases The opposite is true of gases. Carbonated soft drinks are more bubbly if stored in the refrigerator. Warm lakes have less O 2 dissolved in them than cool lakes.

94 Slide 94 / 142 Gases in Solution The solubility of liquids and solids does not change appreciably with pressure. The solubility of a gas in a liquid is directly proportional to its pressure. In general, the solubility of gases in water increases with increasing molar mass. Larger molecules have stronger dispersion forces.

95 Slide 95 / Increasing the temperature the solubility of solids and the solubility of gases in a liquid. A B C D E decreases, increases doesn't affect, increases increases, decreases increases, increases doesn't affect, doesn't affect

96 Slide 96 / Increasing the pressure on a liquid the solubility of solids and the solubility of gases in a liquid. A B C D E decreases, increases doesn't affect, increases increases, decreases increases, increases doesn't affect, doesn't affect

97 Slide 97 / Pressure has an appreciable effect on the solubility of in liquids. A B C D E gases solids liquids salts solids and liquids

98 Slide 98 / 142 Expressing Concentrations of Solutions Recall that solutions are homogeneous mixtures of two or more pure substances. In a solution, the solute is dispersed uniformly throughout the solvent. State of Solution State of Solvent State of Solute Example Gas Gas Gas Air Liquid Liquid Gas Oxygen in water Liquid Liquid Liquid Alcohol in water Liquid Liquid Liquid Salt in water Solid Solid Gas H 2 in Palladium Solid Solid Liquid Hg in Silver Solid Solid Solid Silver in Gold

99 Slide 99 / 142 Mass Percentage of solute mass of A in solution Mass % of solute A = x 100% total mass of solution

100 Slide 100 / The concentration of urea in a solution prepared by dissolving 16 g of urea in 39 g of H 2 O is % by mass. A 29 B 41 C 0.29 D 0.41 E 0.48

101 Slide 101 / A solution contains 11% by mass of sodium chloride. This means that. A B C D there are 11 g of sodium chloride in in 1.0 ml of this solution 100 g of the solution contains 11 g of sodium chloride 100 ml of the solution contains 11 g of sodium chloride the density of the solution is 11 g/ml E the molality of the solution is 11

102 Slide 102 / 142 Mole Fraction (X) Assume a solute A is dissolved in a solvent B X A = moles of A total moles (A+B) in solution In some applications, one needs the mole fraction of solvent, not solute make sure you find the quantity you need!

103 Slide 103 / What is the mole fraction of Nitrogen in a mixture of gas containing 5 moles of Nitrogen and 15 moles of Oxygen. A 0.25 B 4 C 3 D 0.75

104 Slide 104 / The mole fraction of He in a gaseous solution prepared from 4.0 g of He, 6.5 g of Ar, and 10.0 g of Ne is. A 0.60 B 1.5 C 0.20 D 0.11 E 0.86

105 Slide 105 / The mole fraction of urea (MW = 60.0 g/mol) in a solution prepared by dissolving 16 g of urea in 39 g of H 2 O is. A 0.58 B 0.37 C 0.13 D 0.11 E 9.1

106 Slide 106 / 142 Molarity (M) Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different. Molarity is one way to measure the concentration of a solution.since volume is temperature-dependent, molarity can change with temperature. Molarity (M) = moles of the solute volume of solution in liters

107 Slide 107 / When mol of HC 2 H 3 O 2 is combined with enough water to make a ml solution, the concentration of HC 2 H 3 O 2 is M. A 3.33 B 1.67 C D E 0.150

108 Slide 108 / What is the concentration (M) of CH 3 OH in a solution prepared by dissolving 11.7 g of CH 3 OH in sufficient water to give exactly 230 ml of solution? A 11.9 B 1.59 x 10-3 C D 1.59 E 11.9 x 10-3

109 Slide 109 / 142 Molality (m) Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature-dependent. m = mol of solute kg of solvent

110 Slide 110 / The concentration of a benzene solution prepared by mixing 12.0 g C 6 H 6 with 38.0 g CCl 4 is molal. A 4.04 B C D E 0.508

111 Slide 111 / The concentration of HCl in a solution that is prepared by dissolving 5.5 g of HCl in 200g of C 2 H 6 O is molal. A 27.5 B 7.5 x 10-4 C 3.3 x 10-2 D 0.75 E 1.3

112 Slide 112 / Which one of the following concentration units varies with temperature? A B C D E molality mass percent mole fraction molarity all of the above

113 Slide 113 / Which one of the following is a correct expression for molarity? A B C D E mol solute/l solvent mol solute/ml solvent mmol solute/ml solution mol solute/kg solvent μmol solute/l solution

114 Slide 114 / 142 Colligative Properties Colligative propertiesdepend only on the number of solute particles present, not on the identity of the solute particles. Among colligative properties are: Vapor pressure lowering Boiling point elevation Melting point depression Osmotic pressure

115 Slide 115 / 142 Vapor Pressure Lowering Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase. Therefore, the vapor pressure of a solution is lower than that of the pure solvent. Solvent alone Solvent + Solute

116 Slide 116 / 142 Boiling Point Elevation and Freezing Point Depression Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent. The Boiling point elevation and freezing point depression depend on the number solute particles in the solution( colligative property) Pressure Solution T MP Temperature Solvent T BP 1 atm Solution

117 Slide 117 / 142 Colligative Properties and Ionization We said earlier that colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. However, it's important to note that it's the number of particles in solution, not the number of particles before they are dissolved. If a solute ionizes, you can get more particles in solution than you started with...depending on the substance.

118 For instance, Slide 118 / 142 Colligative Properties and Ionization 1 mol NaCl becomes 2 moles of particles in solution: 1 mol Na mol Cl - 1 mol CaCl 2 becomes 3 moles in solution: 1 mol Ca mol Cl - 1 mol C 6 H 12 O 6 (glucose) stays 1 mol since it doesn't disassociate, it's stays a single molecule because it is a molecular compound. So in terms of colligative properties; you get about three times the effect with CaCl 2 (and two times the effect with NaCl) than you do with C 6 H 6.

119 Slide 119 / Which of the following will have the highest boiling point? A B C pure H 2 O 0.10 m aqueous glucose 0.20 m aqueous sucrose ( table sugar) D 0.20 m CaCl 2 E 0.20 m NaCl

120 Slide 120 / Which of the following will have the lowest freezing point? A 0.10 m aqueous sucrose (C 12 H 22 O 11 ) B 0.20 M Pb(NO 3 ) 2 C 0.20 M KOH D 0.20 M NaNO 3 E 0.20 M KCl

121 Slide 121 / Which of the following will have the lowest vapor pressure? A pure H 2 O B 0.20 M Pb(NO 3 ) 2 C 0.20 M AlCl 3 D 0.20 M SrCl 2 E 0.20 M MgF 2

122 Slide 122 / Which of the following aqueous solutions will have the lowest vapor pressure? A 0.25 M glucose, C 6 H 12 O 6 B 0.50 M glucose C 0.50 sucrose, C 12 H 22 O 11 D E 1.0 M sucrose All of these aqueous solutions have equal vapor pressure.

123 Slide 123 / Which of the following aqueous solutions will have the highest vapor pressure? A 0.75 M glucose, C 6 H 12 O 6 B 0.50 M glucose C 0.25 M sucrose, C 12 H 22 O 11 D E 0.50 M sucrose All of these aqueous solutions have equal vapor pressure.

124 Slide 124 / Which of the following will have the highest vapor pressure? A B C D E pure water 1.0 m sucrose (aq) 1.0-m NaCl (aq) 1.0-m HCl (aq) 1.0-m CaCl 2 (aq)

125 Slide 125 / Which of the following will have the lowest vapor pressure? A B C D E pure water 1.0 m sucrose (aq) 1.0-m CaCl 2 (aq) 1.0-m HCl (aq) 1.0-m KCl (aq)

126 Slide 126 / Which of the following will have the highest boiling point? A B C D E pure water 1.0 m sucrose (aq) 1.0-m NaCl (aq) 1.0-m HCl (aq) 1.0-m CaCl 2 (aq)

127 Slide 127 / Which of the following will have the lowest boiling point? A B C D E pure water 1.0 m sucrose (aq) 1.0-m NaCl (aq) 1.0-m HCl (aq) 1.0-m CaCl 2 (aq)

128 Slide 128 / Which of the following will have the highest freezing point? A B C D E pure water 0.20-m glucose (aq) 0.20-m KBr (aq) 0.20-m HCl (aq) 0.20-m AlCl 3 (aq)

129 Slide 129 / Which of the following will have the lowest freezing point? A B C D pure water 0.15-m Mg(NO 3 ) 2 (aq) 0.15-m glucose(aq) 0.15-m NaF (aq) E 0.15-m HBr (aq)

130 Slide 130 / Which of the following aqueous solutions will have the highest boiling point? A B C D E 0.10 m NaCl 0.15 m NaCl 0.20 m NaCl 0.25 m NaCl pure water

131 Slide 131 / As the concentration of a solute in a solution increases, the freezing point of the solution and the vapor pressure of the solution. A B C D E increases, increases increases, decreases decreases, increases unaffected, decreases decreases, decreases

132 Slide 132 / Which of the following solutions will have the lowest freezing point? A B C D E pure H 2 O 0.10 m aqueous glucose 0.15 m aqueous glucose 0.20 m aqueous glucose 0.25 m aqueous glucose

133 Slide 133 / Colligative properties of solutions include all of the following except. A B C D E depression of vapor pressure upon addition of a solute to a solvent elevation of the boiling point of a solution upon addition of a solute to a solvent depression of the freezing point of a solution upon addition of a solute to a solvent an increase in the osmotic pressure of a solution upon the addition of more solute the increase of reaction rates with increase in temperature

134 Slide 134 / 142 Colligative properties Credit to Tom Greenbowe

135 Slide 135 / 142 Osmosis Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles. In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

136 Slide 136 / 142 Osmosis In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration ) to the area of lower solvent concentration (higher solute concentration ).

137 Slide 137 / 142 Osmotic Pressure The pressure required to stop osmosis, known as osmotic pressure, P is PV = nrt P = nrt/v = MRT where M is the molarity of the solution. If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.

138 Slide 138 / 142 Osmosis in Cells If the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic. Water will flow out of the cell, and crenation results.

139 Slide 139 / 142 Osmosis in Cells If the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic. Water will flow into the cell, and hemolysis results.

140 Slide 140 / Osmosis is best defined as the movement of: A B C D E Molecules from an area of high concentration to an area of lower concentration Molecules from an area of low concentration to an area of higher concentration Water molecules across a membrane from an area of low water to an area of higher concentration Water molecules across a membrane from an area of high concentration to low area of concentration Water molecules inside a container

141 Slide 141 / Which of the following will pass through a cell membrane most easily? A B C D E small polar molecules small nonpolar molecules large polar molecules large nonpolar molecules large neutral molecules

142 Slide 142 / 142