Solutions Properties of Water Solutions

Predict the % water in the following foods

Predict the % water in the following foods

Water in the Body liquids 1000 mL urine 1500 mL water gain water loss liquids 1000 mL urine 1500 mL food 1200 mL perspiring 300 mL cells 300 mL exhaling 600 mL feces 100 mL Calculate the total water gain and water loss Total ______ mL _____ mL

Water Most common solvent A polar molecule O - a hydrogen bond H +

Hydrogen Bonds Attract Polar Water Molecules

Explore: Surface Tension HW Fill a glass to the brim with water How many pennies can you add to the glass without causing any water to run over? Predict _________________ Actual _________________ Explain your results

Explore 1. Place some water on a waxy surface. Why do drops form? 2. Carefully place a needle on the surface of water. Why does it float? What happens if you push it through the water surface? 3. Sprinkle pepper on water. What does it do? Add a drop of soap. What happens?

Surface Tension Water molecules within water hydrogen bond in all directions Water molecules at surface cannot hydrogen bond above the surface, pulled inward Water surface behaves like a thin, elastic membrane or “skin” Surfactants (detergents) undo hydrogen bonding

Solute and Solvent Solutions are homogeneous mixtures of two or more substances Solute The substance in the lesser amount Solvent The substance in the greater amount

Nature of Solutes in Solutions Spread evenly throughout the solution Cannot be separated by filtration Can be separated by evaporation Not visible, solution appears transparent May give a color to the solution

Types of Solutions air O2 gas and N2 gas gas/gas soda CO2 gas in water gas/liquid seawater NaCl in water solid/liquid brass copper and zinc solid/solid

Discussion Give examples of some solutions and explain why they are solutions.

Learning Check SF1 (1) element (2) compound (3) solution A. water 1 2 3 B. sugar 1 2 3 C. salt water 1 2 3 D. air 1 2 3 E. tea 1 2 3

Solution SF1 (1) element (2) compound (3) solution A. water 2 B. sugar 2 C. salt water 3 D. air 3 E. tea 3

Learning Check SF2 Identify the solute and the solvent. A. brass: 20 g zinc + 50 g copper solute = 1) zinc 2) copper solvent = 1) zinc 2) copper B. 100 g H2O + 5 g KCl solute = 1) KCl 2) H2O solvent = 1) KCl 2) H2O

Solution SF2 A. brass: 20 g zinc + 50 g copper solute = 1) zinc solvent = 2) copper B. 100 g H2O + 5 g KCl solute = 1) KCl solvent = 2) H2O

Learning Check SF3 Identify the solute in each of the following solutions: A. 2 g sugar (1) + 100 mL water (2) B. 60.0 mL ethyl alcohol(1) and 30.0 mL of methyl alcohol (2) C. 55.0 mL water (1) and 1.50 g NaCl (2) D. Air: 200 mL O2 (1) + 800 mL N2 (2)

Solution SF3 Identify the solute in each of the following solutions: A. 2 g sugar (1) B. 30.0 mL of methyl alcohol (2) C. 1.5 g NaCl (2) D. 200 mL O2 (1)

Like dissolves like A ____________ solvent such as water is needed to dissolve polar solutes such as sugar and ionic solutes such as NaCl. A ___________solvent such as hexane (C6H14) is needed to dissolve nonpolar solutes such as oil or grease.

Learning Check SF4 Which of the following solutes will dissolve in water? Why? 1) Na2SO4 2) gasoline 3) I2 4) HCl

Solution SF4 Which of the following solutes will dissolve in water? Why? 1) Na2SO4 Yes, polar (ionic) 2) gasoline No, nonnpolar 3) I2 No, nonpolar 4) HCl Yes, Polar

Formation of a Solution H2O Hydration Na+ Cl- Na+ Dissolved solute Cl- H2O Na+ Cl- solute

Electrolyte and Non-electrolyte Electrolyte: a substance that conducts electricity when dissolved in water. Acids, bases and soluble ionic solutions are electrolytes. Non-electrolyte: a substance that does not conduct electricity when dissolved in water. Molecular compounds and insoluble ionic compounds are non-electrolytes.

Electrolytes Some solutes can dissociate into ions. Electric charge can be carried.

Types of solutes Strong Electrolyte - 100% dissociation, high conductivity Strong Electrolyte - 100% dissociation, all ions in solution Na+ Cl-

Types of solutes Weak Electrolyte - partial dissociation, slight conductivity Weak Electrolyte - partial dissociation, molecules and ions in solution CH3COOH CH3COO- H+

Types of solutes Non-electrolyte - No dissociation, no conductivity Non-electrolyte - No dissociation, all molecules in solution Sugar C6H12O6

Types of Electrolytes Strong electrolyte dissociates completely. Good electrical conduction. Weak electrolyte partially dissociates. Fair conductor of electricity. Non-electrolyte does not dissociate. Poor conductor of electricity. A generalization is helpful: Essentially all soluble ionic compounds are strong electrolytes. Most molecular compounds are weak electrolytes or non-electrolytes

Representation of Electrolytes using Chemical Equations A strong electrolyte: MgCl2(s) → Mg2+(aq) + 2 Cl- (aq) A weak electrolyte: CH3COOH(aq) ← CH3COO -(aq) +H+(aq) → Strong – complete dissociation Weak – reversible CH3OH(aq) A non-electrolyte:

Electrolytes in Action

Strong Electrolytes Strong acids: HNO3, H2SO4, HCl, HClO4 Strong bases: MOH (M = Na, K, Cs, Rb etc) Salts: All salts dissolving in water are completely ionized. Stoichiometry & concentration relationship NaCl (s)  Na+ (aq) + Cl– (aq) Ca(OH)2 (s)  Ca2+(aq) + 2 OH– (aq) AlCl3 (s)  Al3+ (aq) + 3 Cl– (aq) (NH4)2SO4 (s)  2 NH4 + (aq) + SO42– (aq)

Writing An Equation for a Solution When NaCl(s) dissolves in water, the reaction can be written as H2O NaCl(s) Na+ (aq) + Cl- (aq) solid separation of ions in water

Learning Check SF5 Solid LiCl is added to some water. It dissolves because A. The Li+ ions are attracted to the 1) oxygen atom(-) of water 2) hydrogen atom(+) of water B. The Cl- ions are attracted to the

Solution SF5 Solid LiCl is added to some water. It dissolves because A. The Li+ ions are attracted to the 1) oxygen atom(-) of water B. The Cl- ions are attracted to the 2) hydrogen atom(+) of water

Rate of Solution You are making a chicken broth using a bouillon cube. What are some things you can do to make it dissolve faster? Crush it Use hot water (increase temperature) Stir it

How do I get sugar to dissolve faster in my iced tea? Stir, and stir, and stir Fresh solvent contact and interaction with solute Add sugar to warm tea then add ice Faster rate of dissolution at higher temperature Grind the sugar to a powder Greater surface area, more solute-solvent interaction

Learning Check SF6 You need to dissolve some gelatin in water. Indicate the effect of each of the following on the rate at which the gelatin dissolves as (1) increase, (2) decrease, (3) no change A. ___Heating the water B. ___Using large pieces of gelatin C. ___Stirring the solution

Learning Check SF6 You need to dissolve some gelatin in water. Indicate the effect of each of the following on the rate at which the gelatin dissolves as (1) increase, (2) decrease, (3) no change A. 1 Heating the water B. 2 Using large pieces of gelatin C. 2 Stirring the solution

Solubility Percent Concentration Colloids and Suspensions

Solubility The maximum amount of solute that can dissolve in a specific amount of solvent usually 100 g. g of solute 100 g water

Saturated and Unsaturated A saturated solution contains the maximum amount of solute that can dissolve. Undissolved solute remains. An unsaturated solution does not contain all the solute that could dissolve

SUPERSATURATED SOLUTION Solubility UNSATURATED SOLUTION more solute dissolves SATURATED SOLUTION no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form increasing concentration

Factors Affecting Solid Solubility Polarity Temperature Surface Area Stirring

Factors Affecting Solubility Polarity Temperature Pressure

Intramolecular Bonding Intramolecular bonding refers to the chemical bonding that holds atoms together within a molecule of a compound Covalent bonding and ionic bonding are the two main types of intramolecular bonding Covalent bonding involves the sharing of valence electrons involves the sharing of valence electrons between two atoms. POLAR- unequal sharing of electrons NON POLAR – equal sharing of electrons Ionic bonding involves the transference of valence electrons

SOLUTE POLAR SOLVENT NONPOLAR SOLVENT Ionic Soluble Insoluble Polar Nonpolar soluble

Learning Check S1 At 40C, the solubility of KBr is 80 g/100 g H2O. Indicate if the following solutions are (1) saturated or (2) unsaturated A. ___60 g KBr in 100 g of water at 40C B. ___200 g KBr in 200 g of water at 40C C. ___25 KBr in 50 g of water at 40C

Solution S1 At 40C, the solubility of KBr is 80 g/100 g H2O. Indicate if the following solutions are (1) saturated or (2) unsaturated A. 2 Less than 80 g/100 g H2O B. 1 Same as 100 g KBr in 100 g of water at 40C, which is greater than its solubility C. 2 Same as 60 g KBr in 100 g of water, which is less than its solubility

Temperature and Solubility of Solids Temperature Solubility (g/100 g H2O) KCl(s) NaNO3(s) 0° 27.6 74 20°C 34.0 88 50°C 42.6 114 100°C 57.6 182 The solubility of most solids (decreases or increases ) with an increase in the temperature.

Temperature and Solubility of Solids Temperature Solubility (g/100 g H2O) KCl(s) NaNO3(s) 0° 27.6 74 20°C 34.0 88 50°C 42.6 114 100°C 57.6 182 The solubility of most solids increases with an increase in the temperature.

Temperature and Solubility of Gases Temperature Solubility (g/100 g H2O) CO2(g) O2(g) 0°C 0.34 0.0070 20°C 0.17 0.0043 50°C 0.076 0.0026 The solubility of gases (decreases or increases) with an increase in temperature.

Temperature and Solubility of Gases Temperature Solubility (g/100 g H2O) CO2(g) O2(g) 0°C 0.34 0.0070 20°C 0.17 0.0043 50°C 0.076 0.0026 The solubility of gases decreases with an increase in temperature.

Learning Check S2 A. Why would a bottle of carbonated drink possibly burst (explode) when it is left out in the hot sun ? B. Why would fish die in water that gets too warm?

Solution S2 A. Gas in the bottle builds up as the gas becomes less soluble in water at high temperatures, which may cause the bottle to explode. B. Because O2 gas is less soluble in warm water, the fish may not obtain the needed amount of O2 for their survival.

Gas Solubility Higher Temperature …Gas is LESS Soluble CH4 O2 2.0 O2 Higher Temperature …Gas is LESS Soluble CO Solubility (mM) 1.0 He 10 20 30 40 50

Solubility Curves Show the conditions that affect states of the solution: unsaturated, saturated, supersaturated.

of solubility on temperature Solubility vs. Temperature for Solids 140 KI 130 120 gases solids NaNO3 110 Solubility Table 100 KNO3 90 80 HCl NH4Cl shows the dependence of solubility on temperature 70 Solubility (grams of solute/100 g H2O) 60 NH3 KCl 50 “Solubility Curves for Selected Solutes”   Description: This slide is a graph of solubility curves for 10 solutes. It shows the number of grams of solute that will dissolve in 100 grams of water over a temperature range of 0cC to 10 cC. Basic Concepts The maximum amount of solute that will dissolve at a given temperature in 100 grams of water is given by the solubility curve for that substance. When the temperature of a saturated solution decreases, a precipitate forms. Most solids become more soluble in water as temperature increases, whereas gases become less soluble as temperature increases. Teaching Suggestions Use this slide to teach students how to use solubility curves to determine the solubilities of various substances at different temperatures. Direct their attention to the dashed lines; these can be used to find the solubility of KClO3 at 50 cC (about 21 g per 100 g of H2O). Make sure students understand that a point on a solubility curve represents the maximum quantity of a particular solute that can be dissolved in a specified quantity of solvent or solution at a particular temperature. Point out that the solubility curve for a particular solute does not depend on whether other solutes also are present in the solution (unless there is a common-ion effect; this subject usually is covered at a later stage in a chemistry course). Questions Determine the solubilities (in water) of the following substance at the indicated temperatures: NH3 at 50 oC; KCl at 90 oC; and NaNO3 at 0 oC. Which of the substances shown on the graph is most soluble in water at 20 oC? Which is lease soluble at that temperature? For which substance is the solubility lease affected by changes in temperature? Why do you think solubilities are only shown between 0 oC and 100 oC? In a flask, you heat a mixture of 120 grams of KClO3 and 300 grams of water until all of the KClO3 has just been dissolved. At what temperature does this occur? You then allow the flask to cool. When you examine it later, the temperature is 64 oC and you notice a white powder in the solution. What has happened? What is the mass of the white powder? Compare the solubility curves for the gases HCl, NH3, and SO2) with the solubility curves for the solid solutes. What generalizations(s) can you make about the relationship between solubility and temperature? According to an article in an engineering journal, there is a salt whose solubility in water increases as the water temperature increases from 0 oC to 65 oC. The salt’s solubility then decreases at temperatures above 65 oC, the article states. In your opinion, is such a salt likely to exist? Explain your answer. What could you do to verify the claims of the article? 40 30 NaCl KClO3 20 10 SO2 0 10 20 30 40 50 60 70 80 90 100 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 517

How to determine the solubility of a given substance? Find out the mass of solute needed to make a saturated solution in 100 cm3 of water for a specific temperature(referred to as the solubility). This is repeated for each of the temperatures from 0ºC to 100ºC. The data is then plotted on a temperature/solubility graph, and the points are connected. These connected points are called a solubility curve.

How to use a solubility graph? A. IDENTIFYING A SUBSTANCE ( given the solubility in g/100 cm3 of water and the temperature) Look for the intersection of the solubility and temperature.

Learning Check SG1: What substance has a solubility of 90 g/100 cm3 in water at a temperature of 25ºC ?

Learning Check SG2: What substance has a solubility of 200 g/100 cm3 of water at a temperature of 90ºC ?

Look for the temperature or solubility Locate the solubility curve needed and see for a given temperature, which solubility it lines up with and visa versa.

Learning Check SG3: What is the solubility of potassium nitrate at 80ºC ?

What is the solubility of potassium nitrate at 80ºC ?

Learning Check SG4: At what temperature will sodium nitrate have a solubility of 95 g/100 cm3 ?

Learning Check SG4: At what temperature will sodium nitrate have a solubility of 95 g/100 cm3 ?

Learning Check SG5: At what temperature will potassium iodide have a solubility of 230 g/100 cm3 ?

Learning Check SG5: At what temperature will potassium iodide have a solubility of 130 g/100 cm3 ?

Using Solubility Curves: What is the solubility of sodium chloride at 25ºC in 150 cm3 of water ? From the solubility graph we see that sodium chlorides solubility is 36 g.

Place this in the proportion below and solve for the unknown solubility. Solve for the unknown quantity by cross multiplying.   Solubility in grams = unknown solubility in grams 100 cm3 of water other volume of water ___36 grams____ 150 cm3 water The unknown solubility is 54 grams. You can use this proportion to solve for the other volume of water if you're given the other solubility.

C. Determine if a solution is saturated, unsaturated,or supersaturated. If the solubility for a given substance places it anywhere on it's solubility curve line it is saturated. If it lies above the solubility curve line, then it's supersaturated, If it lies below the solubility curve line it's an unsaturated solution. Remember though, if the volume of water isn't 100 cm3 to use a proportion first as shown above.

Solubility how much solute dissolves in a given amt. of solvent at a given temp. SOLUBILITY CURVE Temp. (oC) Solubility (g/100 g H2O) KNO3 (s) KCl (s) HCl (g) unsaturated: solution could hold more solute; below line saturated: solution has “just right” amt. of solute; on line supersaturated:solution has “too much” solute dissolved in it; above the line

Solids dissolved in liquids Gases dissolved in liquids To Sol. To Sol. As To , solubility As To , solubility

Sometimes you'll need to determine how much additional solute needs to be added to a unsaturated solution in order to make it saturated. For example,30 grams of potassium nitrate has been added to 100 cm3 of water at a temperature of 50ºC. How many additional grams of solute must be added in order to make it saturated?

How many additional grams of solute must be added in order to make it saturated? From the graph you can see that the solubility for potassium nitrate at 50ºC is 84 grams

If there are already 30 grams of solute in the solution, all you need to get to 84 grams is 54 more grams ( 84g-30g )

of solubility on temperature 0 10 20 30 40 50 60 70 80 90 100 Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H2O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO3 KNO3 HCl NH4Cl NH3 NaCl KClO3 SO2 gases solids Solubility Table shows the dependence of solubility on temperature “Solubility Curves for Selected Solutes”   Description: This slide is a graph of solubility curves for 10 solutes. It shows the number of grams of solute that will dissolve in 100 grams of water over a temperature range of 0cC to 10 cC. Basic Concepts The maximum amount of solute that will dissolve at a given temperature in 100 grams of water is given by the solubility curve for that substance. When the temperature of a saturated solution decreases, a precipitate forms. Most solids become more soluble in water as temperature increases, whereas gases become less soluble as temperature increases. Teaching Suggestions Use this slide to teach students how to use solubility curves to determine the solubilities of various substances at different temperatures. Direct their attention to the dashed lines; these can be used to find the solubility of KClO3 at 50 cC (about 21 g per 100 g of H2O). Make sure students understand that a point on a solubility curve represents the maximum quantity of a particular solute that can be dissolved in a specified quantity of solvent or solution at a particular temperature. Point out that the solubility curve for a particular solute does not depend on whether other solutes also are present in the solution (unless there is a common-ion effect; this subject usually is covered at a later stage in a chemistry course). Questions Determine the solubilities (in water) of the following substance at the indicated temperatures: NH3 at 50 oC; KCl at 90 oC; and NaNO3 at 0 oC. Which of the substances shown on the graph is most soluble in water at 20 oC? Which is lease soluble at that temperature? For which substance is the solubility lease affected by changes in temperature? Why do you think solubilities are only shown between 0 oC and 100 oC? In a flask, you heat a mixture of 120 grams of KClO3 and 300 grams of water until all of the KClO3 has just been dissolved. At what temperature does this occur? You then allow the flask to cool. When you examine it later, the temperature is 64 oC and you notice a white powder in the solution. What has happened? What is the mass of the white powder? Compare the solubility curves for the gases HCl, NH3, and SO2) with the solubility curves for the solid solutes. What generalizations(s) can you make about the relationship between solubility and temperature? According to an article in an engineering journal, there is a salt whose solubility in water increases as the water temperature increases from 0 oC to 65 oC. The salt’s solubility then decreases at temperatures above 65 oC, the article states. In your opinion, is such a salt likely to exist? Explain your answer. What could you do to verify the claims of the article? LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 517

Solubility vs. Temperature for Solids 0 10 20 30 40 50 60 70 80 90 100 Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H2O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO3 KNO3 HCl NH4Cl NH3 NaCl KClO3 SO2 gases solids Classify as unsaturated, saturated, or supersaturated. 80 g NaNO3 @ 30oC 45 g KCl @ 60oC 50 g NH3 @ 10oC 70 g NH4Cl @ 70oC =unsaturated per 100 g H2O =saturated =unsaturated =supersaturated

Nosce te ipsum Describe each situation below. (A) Per 100 g H2O, 100 g Unsaturated; all solute NaNO3 @ 50oC. dissolves; clear solution. (B) Cool solution (A) very Supersaturated; extra slowly to 10oC. solute remains in solution; still clear. Nosce te ipsum (C) Quench solution (A) in Saturated; extra solute an ice bath to 10oC. (20 g) can’t remain in solution, becomes visible.

Soluble and Insoluble Salts A soluble salt is an ionic compound that dissolves in water. An insoluble salt is an ionic compound that does not dissolve in water

Aqueous Solutions How do we know ions are present in aqueous solutions? The solutions:_________ They are called ELECTROLYTES HCl, MgCl2, and NaCl are strong electrolytes. They dissociate completely (or nearly so) into ions.

Solubility Rules NH4+ Li+ Na+ K+ or NO3- Examples: soluble salts 1. A salt is soluble in water if it contains any one of the following ions: NH4+ Li+ Na+ K+ or NO3- Examples: soluble salts LiCl Na2SO4 KBr Ca(NO3)2

Cl- Salts positive ion is Ag+, Pb2+, or Hg22+. Examples: 2. Salts with Cl- are soluble, but not if the positive ion is Ag+, Pb2+, or Hg22+. Examples: soluble not soluble(will not dissolve) MgCl2 AgCl PbCl2

SO42- Salts 3. Salts with SO42- are soluble, but not if the positive ion is Ba2+, Pb2+, Hg2+ or Ca2+. Examples: soluble not soluble MgSO4 BaSO4 PbSO4

Other Salts 4. Most salts containing CO32-, PO43-, S2- and OH- are not soluble. Examples: soluble not soluble Na2CO3 CaCO3 K2S CuS

Learning Check S3 Indicate if each salt is (1)soluble or (2)not soluble: A. ______ Na2SO4 B. ______ MgCO3 C. ______ PbCl2 D. ______ MgCl2

Solution S3 Indicate if each salt is (1) soluble or (2) not soluble: A. _1_ Na2SO4 B. _2_ MgCO3 C. _2_ PbCl2 D. _1_ MgCl2

Solutions Molarity

Molarity (M) A concentration that expresses the moles of solute in 1 L of solution Molarity (M) = moles of solute 1 liter solution

Units of Molarity 2.0 M HCl = 2.0 moles HCl 1 L HCl solution

Molarity Calculation NaOH is used to open stopped sinks, to treat cellulose in the making of nylon, and to remove potato peels commercially. If 4.0 g NaOH are used to make 500. mL of NaOH solution, what is the molarity (M) of the solution?

Calculating Molarity = 0.20 M NaOH 1) 4.0 g NaOH x 1 mole NaOH = 0.10 mole NaOH 40.0 g NaOH 2) 500. mL x 1 L _ = 0.500 L 1000 mL 3. 0.10 mole NaOH = 0.20 mole NaOH 0.500 L 1 L = 0.20 M NaOH

Learning Check M1 A KOH solution with a volume of 400 mL contains 2 mole KOH. What is the molarity of the solution? 1) 8 M 2) 5 M 3) 2 M Drano

Solution M1 A KOH solution with a volume of 400 mL contains 2 moles of KOH. What is the molarity of the solution? 2) 5 M M = 2 mole KOH = 5 M 0.4 L Drano

Learning Check M2 A glucose solution with a volume of 2.0 L contains 72 g glucose (C6H12O6). If glucose has a molar mass of 180. g/mole, what is the molarity of the glucose solution? 1) 0.20 M 2) 5.0 M 3) 36 M

Solution M2 A glucose solution with a volume of 2.0 L contains 72 g glucose (C6H12O6). If glucose has a molar mass of 180. g/mole, what is the molarity of the glucose solution? 1) 72 g x 1 mole x 1 = 0.20 M 180. g 2.0 L

Molarity Conversion Factors A solution is a 3.0 M NaOH.. Write the molarity in the form of conversion factors. 3.0 moles NaOH and 1 L NaOH soln 1 L NaOH soln 3.0 moles NaOH

Learning Check M3 1) 15 moles HCl 2) 1.5 moles HCl 3) 0.15 moles HCl Stomach acid is a 0.10 M HCl solution. How many moles of HCl are in 1500 mL of stomach acid solution? 1) 15 moles HCl 2) 1.5 moles HCl 3) 0.15 moles HCl

Solution M3 3) 1500 mL x 1 L = 1.5 L 1000 mL 1.5 L x 0.10 mole HCl = 0.15 mole HCl 1 L (Molarity factor)

Learning Check M4 How many grams of KCl are present in 2.5 L of 0.50 M KCl? 1) 1.3 g 2) 5.0 g 3) 93 g

Solution M4 3) 2.5 L x 0.50 mole x 74.6 g KCl = 93 g KCl 1 L 1 mole KCl

Learning Check M5 1) 150 mL 2) 1500 mL 3) 5000 mL How many milliliters of stomach acid, which is 0.10 M HCl, contain 0.15 mole HCl? 1) 150 mL 2) 1500 mL 3) 5000 mL

Solution M5 0.10 mole HCl 1 L (Molarity inverted) = 1500 mL HCl 2) 0.15 mole HCl x 1 L soln x 1000 mL 0.10 mole HCl 1 L (Molarity inverted) = 1500 mL HCl

Learning Check M6 How many grams of NaOH are required to prepare 400. mL of 3.0 M NaOH solution? 1) 12 g 2) 48 g 3) 300 g

Solution M6 2) 400. mL x 1 L = 0.400 L 1000 mL 0.400 L x 3.0 mole NaOH x 40.0 g NaOH 1 L 1 mole NaOH (molar mass) = 48 g NaOH

Percent Concentration Solution Percent Concentration

Percent Concentration Describes the amount of solute dissolved in 100 parts of solution amount of solute 100 parts solution

Mass-Mass % Concentration mass/mass % = g solute x 100% 100 g solution

Mixing Solute and Solvent Solute + Solvent 4.0 g KCl 46.0 g H2O 50.0 g KCl solution

Calculating Mass-Mass % g of KCl = 4.0 g g of solvent = 46.0 g g of solution = 50.0 g %(m/m) = 4.0 g KCl (solute) x 100 = 8.0% KCl 50.0 g KCl solution

Learning Check PC1 A solution contains 15 g Na2CO3 and 235 g of H2O? What is the mass % of the solution? 1) 15% (m/m) Na2CO3 2) 6.4% (m/m) Na2CO3 3) 6.0% (m/m) Na2CO3

Solution PC1 mass solute = 15 g Na2CO3 mass solution = 15 g + 235 g = 250 g %(m/m) = 15 g Na2CO3 x 100 250 g solution = 6.0% Na2CO3 solution

Mass-Volume % mass/volume % = g solute x 100% 100 mL solution

Learning Check PC2 An IV solution is prepared by dissolving 25 g glucose (C6H12O6) in water to make 500. mL solution. What is the percent (m/v) of the glucose in the IV solution? 1) 5.0% 2) 20.% 3) 50.%

Solution PC2 1) 5.0% %(m/v) = 25 g glucose x 100 500. mL solution = 5.0 %(m/v) glucose solution

Writing Factors from % A physiological saline solution is a 0.85% (m/v) NaCl solution. Two conversion factors can be written for the % value. 0.85 g NaCl and 100 mL NaCl soln 100 mL NaCl soln 0.85 g NaCl

% (m/m) Factors NaOH and NaOH soln Write the conversion factors for a 10 %(m/m) NaOH solution NaOH and NaOH soln NaOH soln NaOH

% (m/m) Factors NaOH and NaOH soln 10 g 100 g 100 g 10 g Write the conversion factors for a 10 %(m/m) NaOH solution NaOH and NaOH soln NaOH soln NaOH 10 g 100 g 100 g 10 g

Learning Check PC 3 Write two conversion factors for each of the following solutions: A. 8 %(m/v) NaOH B. 12 %(v/v) ethyl alcohol

Solution PC 3 Write conversion factors for the following: A. 8 %(m/v) NaOH 8 g NaOH and 100 mL 100 mL 8 g NaOH B. 12 %(v/v) ethyl alcohol 12 mL alcohol and 100 mL 100 mL 12 mL alcohol

Using % Factors How many grams of NaCl are needed to prepare 250 g of a 10.0% (m/m) NaCl solution? Complete data: ____________ g solution ____________% or (______/_100 g_) solution ____________ g solute

Clculation Using % Factors 250 g solution 10.0% or (10.0 g/100 g) solution ? g solute 250 g NaCl soln x 10.0 g NaCl = 25 g NaCl 100 g NaCl soln

Learning Check PC4 How many grams of NaOH do you need to measure out to prepare 2.0 L of a 12%(m/v) NaOH solution? 1) 24 g NaOH 2) 240 g NaOH 3) 2400 g NaOH

Solution PC4 2.0 L soln x 1000 mL = 2000 mL soln 1 L 12 % (m/v) NaOH = 12 g NaOH 100 mL NaOH soln 2000 mL x 12 g NaOH = 240 g NaOH

Learning Check PC5 1) 30 mL 2) 3000 mL 3) 7500 mL How many milliliters of 5 % (m/v) glucose solution are given if a patient receives 150 g of glucose? 1) 30 mL 2) 3000 mL 3) 7500 mL

Solution PC5 5% m/v factor 150 g glucose x 100 mL = 3000 mL

Preparing a Solution by Dilution

Units of Concentrations amount of solute per amount of solvent or solution g solute g solution x 100 g solute g solute + g solvent x 100 = Percent (by mass) = moles of solute volume in liters of solution Molarity (M) = moles = M x VL

Colloids and Suspensions Solutions Colloids and Suspensions Osmosis and Dialysis

Solutions Have small particles (ions or molecules) Are transparent Do not separate Cannot be filtered Do not scatter light.

Colloids Have medium size particles Cannot be filtered Separated with semipermeable membranes Scatter light (Tyndall effect)

Examples of Colloids Fog Whipped cream Milk Cheese Blood plasma Pearls

Suspensions Have very large particles Settle out Can be filtered Must stir to stay suspended

Examples of Suspensions Blood platelets Muddy water Calamine lotion

Osmosis In osmosis, the solvent water moves through a semipermeable membrane Water flows from the side with the lower solute concentration into the side with the higher solute concentration Eventually, the concentrations of the two solutions become equal.

Osmosis semipermeable membrane 4% starch 10% starch H2O

Equilibrium is reached. water flow becomes equal 7% starch 7% starch H2OO

Osmotic Pressure Produced by the number of solute particles dissolved in a solution Equal to the pressure that would prevent the flow of additional water into the more concentrated solution Increases as the number of dissolved particles increase

Osmotic Pressure of the Blood Cell walls are semipermeable membranes The osmotic pressure of blood cells cannot change or damage occurs. The flow of water between a red blood cell and its surrounding environment must be equal

Isotonic solutions Exert the same osmotic pressure as red blood cells. Medically 5% glucose and 0.9% NaCl are used their solute concentrations provide an osmotic pressure equal to that of red blood cells H2O

Hypotonic Solutions Lower osmotic pressure than red blood cells Lower concentration of particles than RBCs In a hypotonic solution, water flows into the RBC The RBC undergoes hemolysis; it swells and may burst. H2O

Hypertonic Solutions H2O Has higher osmotic pressure than RBC Has a higher particle concentration In hypertonic solutions, water flows out of the RBC The RBC shrinks in size (crenation) H2O

Dialysis Occurs when solvent and small solute particles pass through a semipermeable membrane Large particles retained inside Hemodialysis is used medically (artificial kidney) to remove waste particles such as urea from blood

Colligative Properties On adding a solute to a solvent, the properties of the solvent are modified. Vapor pressure decreases Melting point decreases Boiling point increases Osmosis is possible (osmotic pressure) These changes are called COLLIGATIVE PROPERTIES. They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles.

Change in Freezing Point Ethylene glycol/water solution Pure water The freezing point of a solution is LOWER than that of the pure solvent

Change in Freezing Point Common Applications of Freezing Point Depression Ethylene glycol – deadly to small animals Propylene glycol

Change in Freezing Point Common Applications of Freezing Point Depression Which would you use for the streets of Bloomington to lower the freezing point of ice and why? Would the temperature make any difference in your decision? sand, SiO2 Rock salt, NaCl Ice Melt, CaCl2

Change in Boiling Point Common Applications of Boiling Point Elevation