Unit 2 Water and Solutions.

Unit 2 Water and Solutions

Water and Heterogeneous systems
Objectives: Explain surface tension Describe the structure of ice and liquid water. Distinguish between a suspension and a solution Identify the characteristics of a colloid.

Water in the liquid state:
Water is a polar molecule. d+ H2O d-

Hydrogen bonding: the negative end of the water molecule is attracted to the positive ends of other molecules- resulting in hydrogen bonding. O H Hydrogen bond d- d+

Water molecules “stick”
together to create surface tension to support light weight objects. It also tends to hold a drop of a liquid in a spherical shape. Surfactants interfere with hydrogen bonding reducing surface tension. This attraction between water molecules slows the tendency of water evaporation. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Suspensions: a mixture from which particles settle out upon standing.
Have very large particles Can be filtered Ex. Muddy water some medicines (antibiotics)

Colloids Have medium sized particles Can not be filtered.
Many colloids are cloudy or milky in appearance. Show Tyndall effect (the path of light is visible)

Types of colloids Colloid type Dispersion medium Dispersed substance
Examples Aerosol Gas Liquid Solid Solid emulsion

Types of colloids Colloid type Dispersion medium Dispersed substance
Examples Aerosol Gas Liquid Clouds, fog,aerosol spray Solid Smoke, dust Foam Shaving cream, whipped cream Emulsion Mayonaisse, milk Sol Paint, ink Solid aerosol Marshmallow, styrofoam Solid emulsion gel Solid sol pearl

Classwork: Complete the table on the last slide, giving examples for the different types of colloids.

Solutions Objectives: Distinguish between a solvent and a solute
Describe what happens in the solution process Explain what electrolytes are. Identify the factors that determine the rate at which a solute dissolves. Identify the factors that determine the mass of solute that will dissolve in a given mass of solvent.

Definitions Solute - substance being dissolved
Solution - homogeneous mixture Solute - substance being dissolved Solvent - present in greater amount

Solution = Solute + Solvent
Solute - gets dissolved Solvent - does the dissolving Aqueous (water) Tincture (alcohol) Amalgam (mercury) Organic Polar Non-polar Dental filling Solutions are always homogeneous – evenly mixed. Solutions – In all solutions, whether gaseous, liquid, or solid, the substance present in the greatest amount is the solvent, and the substance or substances present in lesser amounts are the solute(s). – Solute does not have to be in the same physical state as the solvent but the physical state of the solvent determines the state of the solution. – If solute and solvent combine to give a homogeneous solution, solute is said to be soluble in the solvent. The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems. – Vitamins can be classified as either fat soluble or water soluble. 1. Fat-soluble vitamins (Vitamin A) are nonpolar, hydrophobic molecules and tend to be absorbed into fatty tissues and stored there. 2. Water-soluble vitamins (Vitamin C) are polar, hydrophilic molecules that circulate in the blood and intracellular fluids and are excreted from the body and must be replenished in the daily diet.

Types of Solutions Solute Solvent Solution Gaseous Solutions
liquid air (nitrogen, oxygen, argon gases) humid air (water vapor in air) Liquid Solutions solid carbonated drinks (CO2 in water) vinegar (CH3COOH in water) salt water (NaCl in water) Solid Solutions dental amalgam (Hg in Ag) sterling silver (Cu in Ag), alloys Solutions are not limited to gases and liquids; solid solutions also exist. • Amalgams, which are usually solids, are solutions of metals in liquid mercury. • Network solids are insoluble in all solvents with which they do not react chemically; covalent bonds that hold the network together are too strong to be broken and are much stronger than any combination of intermolecular interactions that might occur in solution. • Most metals are insoluble in all solvents but do react with solutions such as aqueous acid or base to produce a solution; in these cases the metal undergoes a chemical transformation that cannot be reversed by removing the solvent. Charles H.Corwin, Introductory Chemistry 2005, page 369

Solvation solute particles are surrounded by solvent particles
Solvation – the process of dissolving: solute particles are surrounded by solvent particles The interactions that determine the solubility of a substance in a liquid depend on the chemical nature of the solute whether it is ionic or molecular) rather than on its physical state (solid, liquid, or gas). • Two examples: 1. Forming a solution of a molecular species in a liquid solvent 2. Formation of a solution of an ionic compound Solutions of molecular substances in liquids – London dispersion forces, dipole-dipole interactions, and hydrogen bonds that hold molecules to other molecules are weak. – Energy is required to disrupt these interactions, and unless some of that energy is recovered in the formation of new, favorable solute-solvent interactions, the increase in entropy on solution formation is not enough for a solution to form. – For solutions of gases in liquids, the energy required to separate the solute molecules is ignored (H2 = 0) because molecules are already separated – it is necessary to only consider the energy required to separate the solvent molecules (H1) and the energy released by new solute-solvent interactions (H3). 1. Nonpolar gases are most soluble in nonpolar solvents because H1 and H3 are both small and of similar magnitude. 2. Nonpolar gases are less soluble in polar solvents than in nonpolar solvents because H1 >> H3. 3. Solubilities of nonpolar gases in water increase as the molecular mass of the gas increases. – All common organic liquids, whether polar or not, are miscible; the strengths of the intermolecular attractions are comparable, the enthalpy of solution is small, and the increase in entropy drives the formation of a solution. – If predominant intermolecular interactions in two liquids are very different from one another, they may be immiscible, and when shaken with water, they form separate phases or layers separated by an interface. – Only the three lightest alcohols are completely miscible with water; as the molecular mass of the alcohol increases, so does the proportion of hydrocarbon in the molecule, which leads to fewer favorable electrostatic interactions with water Hydrophilic and hydrophobic solutes – A solute can be classified as hydrophilic, meaning that there is an electrostatic attraction to water, or hydrophobic, meaning that it repels water. 1. Hydrophilic substance is polar and contains O–H or N–H groups that can form hydrogen bonds to water; tend to be very soluble in water and other strongly polar solvents 2. Hydrophobic substance may be polar but usually contains C–H bonds that do not interact favorably with water; essentially insoluble in water and soluble in nonpolar solvents – The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems. – Vitamins can be classified as either fat soluble or water soluble. 1. Fat-soluble vitamins (Vitamin A) are nonpolar, hydrophobic molecules and tend to be absorbed into fatty tissues and stored there. 2. Water-soluble vitamins (Vitamin C) are polar, hydrophilic molecules that circulate in the blood and intracellular fluids and are excreted from the body and must be replenished in the daily diet. solute particles are separated and pulled into solution

Dissolving of NaCl - - - O + + + + Cl- Na+ hydrated ions H
Timberlake, Chemistry 7th Edition, page 287

Electrolytes Classify each compound as either a strong electrolyte or a nonelectrolyte. If a compound is a nonelectrolyte, the concentration is the same as the molarity of the solution. If a compound is a strong electrolyte, determine the number of each ion contained in one formula unit and find the concentration of each species by multiplying the number of each ion by the molarity of the solution. (a) Nonelectrolyte (b) Weak electrolyte (c) Strong electrolyte Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Electrolytes Electrolytes - solutions that carry an electric current
Electrolyte — any compound that can form ions when it dissolves in water – When strong electrolytes dissolve, constituent ions dissociate completely, producing aqueous solutions that conduct electricity very well. – When weak electrolytes dissolve, they produce relatively few ions in solution, and aqueous solutions, of weak electrolytes do not conduct electricity as well as solutions of strong electrolytes. – Nonelectrolytes dissolve in water as neutral molecules and have no effect on conductivity. strong electrolyte weak electrolyte nonelectrolyte NaCl(aq) Na+ + Cl- HF(aq) H+ + F- Timberlake, Chemistry 7th Edition, page 290

Electrolyte Imbalances
Normal range (mmol / L) Excess Defiency Sodium Na+ Hypernatremia (increased urine excretion; excess water loss) Hyponatremia (dehydration; diabetes-related low blood pH; vomiting, diarrhea) Potassium K+ 3.5 – 5.0 Hyperkalemia (renal failure, low blood pH) Hypokalemia (gastointestinal conditions) Hydrogen carbonate HCO3- Hypercapina (high blood pH; hypoventilation) Hypocapnia (low blood pH; hyper-ventilation; dehydration) Chloride Cl- Hyperchloremia (anemia, heart conditions, dehydration) Hypochloremia (acute infections; burns; hypoventilation) When you sweat, you lose electrolytes (ions in your body fluids). These electrolytes conduct electricity to tell your muscles to work. Severe dehydration and sweating will cause you to cramp and may lead to heat stroke related injuries. It is important to maintain your electrolytic balance. Drinking pop will not do this nor will only drinking water Sports drinks have electrolytes in them to restore the electrolytic balance in our body.

Isotonic Hypotonic Hypertonic (a) Cells in dilute salt solution
(b) Cells in distilled water (c) Cells in concentrated salt solution Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

Solution formation The “nature” (polarity, or composition) of the solute and the solvent will determine… Whether a substance will dissolve How much will dissolve

Factors determining rate of solution. 1
Factors determining rate of solution Stirring (agitation) moves fresh solvent into contact with the solute. 2. Smaller pieces increases the amount of surface area of the solute. 3. Higher temperature makes the molecules of the solvent move faster and contact the solute harder and more often. Speeds up dissolving. Higher Temperature ALSO Usually increases the amount that will dissolve (an exception is gases ).

Solubility- is the maximum amount of substance that will dissolve at a specific temperature. The units for solubility are: grams of solute/100 grams solvent Saturated solution- Contains the maximum amount of solute dissolved NaCl = 36.0 g/100 mL water Unsaturated solution- Can still dissolve more solute (for example 28.0 grams of NaCl/100 mL) Supersaturated- solution that is holding (or dissolving) more than it theoretically can; a “seed crystal” will make it crystallize. Very unstable.

SUPERSATURATED SOLUTION
Solubility UNSATURATED SOLUTION more solute dissolves SATURATED SOLUTION no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form Maximum amount of a solute that can dissolve in a solvent at a specified temperature and pressure is its solubility. – Solubility is expressed as the mass of solute per volume (g/L) or mass of solute per mass of solvent (g/g) or as the moles of solute per volume (mol/L). – Solubility of a substance depends on energetic factors and on the temperature and, for gases, the pressure. • A solution that contains the maximum possible amount of solute is saturated. • If a solution contains less than the maximum amount of solute, it is unsaturated. When a solution is saturated and excess solute is present, the rate of dissolution is equal to the rate of crystallization. • Solubility increases with increasing temperature — a saturated solution that was prepared at a higher temperature contains more dissolved solute than it would contain at a lower temperature, when the solution is cooled, it can become supersaturated. increasing concentration

a. The solubility of the KNO3 increases as the temperature increases.
b. Yb2(SO4)3 shows a decrease in solubility as the temperature increases, and NaCl shows the least change in solubility as temperature changes. c. Only a negligible amount of NaCl would go into solution, if any.

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 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 517

Solubility vs. Temperature for Solids
a) 80g NaNO3 at 45C. The solution is ? 140 unsaturated KI 130 120 gases solids b) 100g NaNO3 at 45C. The solution is ? NaNO3 110 100 saturated KNO3 90 80 c) 120g NaNO3 at 45C. The solution is ? HCl NH4Cl 70 Solubility (grams of solute/100 g H2O) 60 NH3 supersaturated KCl 50 d) How much more NaNO3 can you add to a solution with 40g of NaNO3 at 45C until it becomes saturated? “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 60 g 10 SO2 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 517

Classwork: solubility graph handout

Liquids Miscible means that two liquids can dissolve in each other
water and antifreeze water and ethanol Partially miscible- slightly water and ether Immiscible means they can’t oil and vinegar

Solids tend to dissolve best when:
They are heated They are stirred Crushed into smaller particles Gases tend to dissolve best when: The solution is cold The pressure is high

For gases in a liquid, the effect is the opposite of solids in liquids
For solids in liquids, as the temperature goes up-the solubility usually goes up For gases in a liquid, the effect is the opposite of solids in liquids As the temperature goes up, gas solubility goes down Think of boiling water bubbling? Thermal pollution may result from industry using water for cooling Solids dissolved in liquids Gases dissolved in liquids To Sol. As To , solubility

Concentration of Solutions
Objectives: Solve problems involving the molarity of a solution. Describe the effect of dilution on the total moles of solute in solution. Define percent by volume and percent by mass solutions.

Concentration is... a measure of the amount of solute dissolved in a given quantity of solvent A concentrated solution has a large amount of solute A dilute solution has a small amount of solute These are qualitative descriptions But, there are ways to express solution concentration quantitatively (NUMBERS!)

Concentrated vs. Dilute
Notice how dark the solutions appears. Notice how light the solution appears. Small amount of solute in a large amount of solvent. Lots of solute, in a small amount of solvent.

Making solutions Pour in a small amount of the solvent, maybe about one-half Then add the pre-massed solute (and mix by swirling to dissolve it) Carefully fill to final volume.

Molarity: a unit of concentration
Molarity = n (moles of solute) V (liters of solution) Abbreviated with a capital M, such as 6.0 M This is the most widely used concentration unit used in chemistry.

- Page 481

How many grams of sodium chloride, NaCl, do you need to prepare 250
How many grams of sodium chloride, NaCl, do you need to prepare 250.mL of a 0.5M NaCl solution? Molar mass NaCl= 58.5 g/mol Molarity= n n= mass V molar mass n= Molarity x V = 0.5 mol L =0.125 mol L Mass= molar mass x n = 58.5 g mol = 7.31 g mol

2. What volume of a 1. 0M NaCl solution can you prepare if you have 45
2. What volume of a 1.0M NaCl solution can you prepare if you have 45.0g of NaCl? Molar mass NaCl = 58.5g / mol? n= mass = g mol = mol molar mass g Molarity (M) = n V= n = mol L =0.769 L V M mol Classwork : p 54 # 19-21

Percent solutions can be expressed by a) volume or b) mass
Percent by volume: = Volume of solute x Volume of solution indicated %(v/v) Vsolution= Vsolute + Vsolvent msolution= msolute + msolvent

Percent by mass: = Mass of solute(g) x 100 Volume of solution (mL)
Indicated %(m/v) More commonly used Another way to do mass percentage is as mass/mass: Percent by mass: = Mass of solute(g) x Mass of solution (g) Indicated %(m/m)

1) 4. 8 g of NaCl are dissolved in 82 mL of solution
1) 4.8 g of NaCl are dissolved in 82 mL of solution. What is the percent of the solution? % (m/V) = mass solute x100 = 4.8 g x 100 = 5.85% volume solution mL 2) How many grams of salt are there in 52 mL of a 6.3 % solution? Mass solute = %(m/v) x volume solution = 6.3 x 52 = 3.28 g

(mass solute + mass solvent)
3. You mix g of salt with 225g of water. What is the %(m/m) concentration of the solution? %(m/m) = mass solute x 100 (mass solute + mass solvent) %(m/m) = x 100 = 10% ( ) Classwork: percent composition handout

Dilution Adding water to a solution will reduce the number of moles of solute per unit volume but the overall number of moles remains the same! Think of taking an aspirin with a small glass of water vs. a large glass of water You still have one aspirin in your body, regardless of the amount of water you drank, but a larger amount of water makes it more diluted.

Dilution The number of moles of solute in solution doesn’t change if you add more solvent! The # moles before = the # moles after Formula for dilution: M1 x V1 = M2 x V2 M1 and V1 are the starting concentration and volume; M2 and V2 are the final concentration and volume. Stock solutions are pre-made solutions to known Molarity.

1. You need to prepare 500. mL of a 0. 5M KCl solution
1. You need to prepare 500. mL of a 0.5M KCl solution. What volume of a 2.0 M KCl solution do you need? M1 x V1 = M2 x V2 M1=2.0M V1= ? M2= 0.5M V2= 500.mL V1= M2V2 = 0.5M 500mL = 125 mL M M

M1 x V1 = M2 x V2 M1=3.0M V1= 50.0mL M2= ? V2= 250.mL
2. You add 200 mL of water to 50.0 mL of a 3.0M NaCl solution. What is the new concentration of the solution? M1 x V1 = M2 x V2 M1=3.0M V1= 50.0mL M2= ? V2= 250.mL M2= M1V1 = 3.0M 50.0mL = 0.6 M V mL Classwork: Dilutions handout

Colligative Properties of Solutions
OBJECTIVES: Identify three colligative properties of solutions. Explain why the vapor pressure, freezing point, and boiling point of a solution differ from those properties of the pure solvent. Solve problems related to the molality and mole fraction of a solution. Describe how freezing point depression and boiling point elevation are related to molality.

Colligative Properties
-These depend only on the number of dissolved particles -Not on what kind of particle -Three important colligative properties of solutions are: Vapor pressure lowering Boiling point elevation Freezing point lowered

Colligative Properties
Some particles in solution will IONIZE (or split), while others may not. Colligative Properties CaCl2 will have three particles in solution for each one particle it starts with. Glucose will only have one particle in solution for each one particle it starts with. NaCl will have two particles in solution for each one particle it starts with.

Vapor Pressure is LOWERED
Surface area is reduced, thus less evaporation, which is a surface property The bonds between molecules keep molecules from escaping. So, in a solution, some of the solvent is busy keeping the solute dissolved. This lowers the vapor pressure Electrolytes form ions when they are dissolved, making more pieces. NaCl ® Na+ + Cl- (this = 2 pieces) More pieces = a bigger effect

Boiling Point is ELEVATED
The vapor pressure determines the boiling point. (Boiling is defined as when the vapor pressure of liquid = vapor pressure of the atmosphere). Lower vapor pressure means you need a higher temperature to get it to equal atmospheric pressure Salt water boils above 100ºC The number of dissolved particles determines how much, as well as the solvent itself.

Freezing Point is LOWERED
Solids form when molecules make an orderly pattern called “crystals” The solute molecules break up the orderly pattern. Makes the freezing point lower. Salt water freezes below 0ºC How much lower depends on the amount of solute dissolved.

The addition of a solute would allow a LONGER temperature range, since freezing point is lowered and boiling point is elevated.

Molality (abbreviated m)
a new unit for concentration m = Moles of solute kilogram of solvent

Mole fraction X = This is another way to express concentration
It is the ratio of moles of solute to total number of moles of solute plus solvent moles solute (moles of solute + moles of solvent) The sum of the mole fractions of all the components of a solution equals one. There is no unit for mole fraction. X =

Ex. 1 Calculate the molality of a solution made by dissolving 45
Ex.1 Calculate the molality of a solution made by dissolving 45.0g of glucose, C6H12O6, in g of water. m = Moles of solute kilogram of solvent Moles of solute = 45.0g mol = mol 180g m = moles of solute = 0.25 mol = 0.5 m kg of solvent kg

Molar mass glucose : 180.0g/mol Molar mass water: 18.0g/mol
Ex. 2 What are the mole fractions of glucose and water in a solution made of 7.59g of glucose, C6 H12O6 dissolved in 125 g of water? Molar mass glucose : 180.0g/mol Molar mass water: 18.0g/mol Mol glucose= g mol = mol mol water= 125g mol = 6.94 mol 180.0g g Xglucose = moles solute = = (moles solute + moles solvent ( ) Xwater = 1 – =

CW p 144 # 1-3, 145 # 4-5; 147 #5,6

Unit 2 Water and Solutions.

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