Chem202

LEARNING GUIDE

PROPERTIES OF SOLUTIONS

July 05, 2014

I – OBJECTIVES: At the end of the lesson, the students will be able to:

a. Define operationally solute and solvent

b. Explain the effect of pressure, temperature, particle size and nature of    

    substance on  solubility and  the rate of dissolution

c. Express solution concentration in percent by weight, mole fraction, molality      

    and molarity

d. Describe hemolysis, crenation, isotonic solution, hypertonic solution, hypotonic

    solution, and  physiological saline solution

e. Recognize and appreciate the importance of some properties of solutions

II – STATEMENT OF THE PROBLEM:

a. What is the difference between each pair of terms: solute and solvent, saturated

     solution  and unsaturated solution, hypertonic and hypotonic ?

            b. How to solve solution concentrations in percent by weight, mole fraction,  

                 molality and molarity?

            c. Why we must drink fresh water solution and not sea water?

            d. Why solutions important in our lives?

III – TOPIC: Solutions and Their Properties

a. Types of solutions

b. Solubility

c. Factors Affecting Solubility and Rate of Dissolution

d. Concentration of Solutions

e. Application of Solutions

IV – LESSON PROPER:

A. Motivation:

Show and animation coming from the Phet Colorado Simulation and let the learners interact, give comments with the presentation.

B. Discussion:

Ø  To describe a solution, two terms are commonly used: solute and solvent.

Ø  There are three types of solutions:

o    gaseous solution- made by mixing one gas with another

o   liquid solution – made by dissolving a gas, solid or liquid in a liquid

o   solid solution – are solids in which one component is randomly dispersed in an atomic or molecular scale throughout another component.

Ø  There are three types of solutions that can occur in your body based on solute concentration: isotonic, hypotonic, and hypertonic.

o   An isotonic solution is one in which the concentration of solutes is the same both inside and outside of the cell.

o   A hypotonic solution is one in which the concentration of solutes is greater inside the cell than outside of it.

o   A hypertonic solution is one where the concentration of solutes is greater outside the cell than inside it.

Ø  The substance that dissolves in the solvent is said to be soluble. An insoluble substance does not dissolve in the solvent. Solubility refers to the maximum amount of solute, expressed in grams that can be dissolved in 100 g of water at a specific temperature.

Ø  There are several factors affecting the solubility of a solute in a solvent. Among these are the natures of the solute, temperature, pressure and particle size. Pressure appreciably in affects the solubility of gases in liquids.

Ø  The rate of dissolution may be increased by stirring, crushing the solute and heating.

Ø   The relative amount of solute in a given amount of solvent or solution is referred to as concentration. There are varieties of ways of expressing concentration. These are percentage concentration, mole fraction, molarity and molality.

Ø  Solutions are essentials to life and throughout the living world, solutions are necessary for maintenance and survival.

C. Generalization:

            Let the learners state what are they have learned from the lesson.

D. Evaluation:

Five minutes paper and pencil test

V – ASSIGNMENT

Make a reflection on how solutions help you in your day to day life

References

Ben-Naim, Arieh (1974). Water and Aqueous Solutions: Introduction to a Molecular Theory. New    York: Plenum Press.

Estrella E. Mendoza, T. F. (2004). Chemistry. Quezon City: Phoenix Publishing House, Inc.

Whitten, K. W.; Davis, R. E.; Peck, M. L.; and Stanley, G. G. (2004). General Chemistry. Pacific               Grove, CA: Brooks/Cole.

Read more: http://www.chemistryexplained.com/Ru-Sp/Solution-Chemistry.html#ixzz36N0kg1WL

http://www.chem.wisc.edu/deptfiles/genchem/sstutorial/Text11/Tx112/tx112.htmhttp://education-portal.com/academy/lesson/hypertonic-solution-definition-effect-example.html#lesson

http://www.ehow.com/info_8542616_three-factors-affect-dissolving-process.html

  

PROPERTIES OF SOLUTIONS

A solution is a homogenous mixture of two or more substances that exist in a single phase. There are two main parts to any solution.

The solute is the component of a solution that is dissolved in the solvent; it is usually present in a smaller amount than the solvent.

The solvent is the component into which the solute is dissolved, and it is usually present in greater concentration.

For example, in a solution of salt water, salt is the solute and water is the solvent. In solutions where water is the solvent, the solution is referred to as an aqueous solution.

Activity No. 1 The purpose of this lab is to: ·         Classify different solutions as one of nine types depending on its solute and solvent phase. ·         Provide other examples of the nine solution types. ·         Determine which substance in a solution is acting as the solute and which is acting as the solvent.

TYPES OF SOLUTION

Introduction:

Consider for a moment what is meant by the word solution. What is an example of a solution that first comes to mind?

What would be your response if you were told that smoke particles in air are a solution?

Solutions are of extreme importance in everyday life in that they comprise some of the foods we eat and many of the drinks we consume, and they also form the basis of many household products such as cleansers, disinfectants, coatings and medications. We often think that only things like salt dissolving in water are a solution when, in fact, things like smoke particles suspended in the air are also solutions.

A solution is defined as a homogeneous mixture containing a solute and a solvent.  The solute (the substance being dissolved) is present in the lesser amount in the solution, while the solvent (the substance that dissolves the solute) is present in the larger amount.  At the molecular level, molecules and ions of a solute are completely mixed with and interact with those of the solvent when a solute dissolves in a solvent. This type of mixing is homogeneous because no boundary is visible in the entire solution.  This interaction occurs because of electrostatic attractions between the solute and solvent.

An example is table salt (NaCl) dissolving into water.  Click on the link below to view the dissolving process at the molecular level. Note that the molecular nature of the water and sodium and chloride ions allows the dissolving to occur and the solution to result.

http://www.northland.cc.mn.us/biology/Biology1111/animations/dissolve.html

Another animation showing the dissolving process (by Tom Greenbowe):

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/thermochem/solutionSalt.html

Some solutions appear to be homogeneous, but in fact they are colloidal dispersions, where tiny droplets of solute are suspended in a solvent because the electrostatic attractions are minimal.  These are not true solutions but are sometimes referred to as such.  Examples include paints, ceramics and blended food. 

Since material can exist in three states (solid, liquid, and gas), there are nine possible combinations of solute and solvent that can make up a solution:

Gaseous mixtures are usually homogeneous and all gaseous mixtures are gas-gas solutions.  Most gases can form solutions with each other (unless they react with each other).  The air is a natural gas-gas solution.  Air is made up primarily of nitrogen (~78%) and oxygen (~21%) with trace amounts of argon, carbon dioxide and water vapour.

 Liquid mixtures are our most easily recognized mixtures. When molecules of gas, solid or liquid are dispersed and mixed with those of liquid, the homogeneous states are called liquid solutions. Solids, liquids and gases dissolve in liquids to form liquid solutions. 

Solid Solutions are much lesson common. Many alloys, ceramics, and polymer blends are solid solutions. Copper and zinc dissolve in each other and harden to give solid solutions called brass. Silver, gold, and copper form many different alloys with unique colors and appearances.

Student Instructions: 

View the item or read the description at each of the 18 stations and complete column two of the student worksheet by including both the station number and example observed at the station.  It is important to consider what the constituents of the solution are (solid, liquid or gas solute in solid liquid or gas solvent). You must think about what these constituents are at the molecular level. Provide additional examples of the nine types of solutions in column four of the student worksheet.  Finally, in column five, determine which substance is acting as the solute and which is acting as the solvent at the station observed.

Station Description:

Station #1: A bottle of household vinegar (5% liquid acetic acid dissolved in water)

Station #2:  Sugar cubes dissolving in water

Station #3:  Solid Air Freshener subliming into the air when heated

Station #4:  Hydrogen dissolving in a transition metal such as platinum (a hydrogen storage strategy for releasing hydrogen for H_2-powered vehicles).

Station #5:  A brass candle holder  (brass is an alloy of the solids copper and zinc)

Station #6:  A bottle of rubbing alcohol   formed from liquid alcohol and water

Station #7:  A mercury-gold amalgam  (formed when liquid mercury is used to extract gold from gold-bearing rock) or mercury-zinc amalgam used in tooth fillings

Station #8:  Seawater

Station #9:  Soap suds

station #10:  A Styrofoam cup

Station #11:   The atmosphere above the water when water evaporates in a sealed water cooler  container

Station #12:  Smoke (solid particulates being released into the  air) resulting from a fire 

Station #13:  A Car tire (made up of solid carbon in rubber)

Station #14:  A piece of (cheap) jewelry made up of copper and gold 

Station #15:  Methane gas released from the valve in a chemistry lab into the air

Station #16:   A bottle of antifreeze (liquid ethylene glycol in water)

Station #17:  Kool-Aid powder poured into water

Station #18:  Food coloring diffusing into water

Student Worksheet

Type of Solution (solute-solvent)	Station Number Example	Example #1	Example #2	State which substance is the solute and solvent for Station # listed
gas-solid		air in solid food (ex: generic ice-cream)
gas-liquid		carbon dioxide in water (Pepsi™)
gas-gas		oxygen in nitrogen (air)
liquid-solid		mercury in silver-tin dental alloy
liquid-liquid		various hydrocarbons in each other (petroleum)
liquid-gas		water vapour in air (humidity)
solid-solid		carbon in iron (steel)
solid-liquid		ammonium nitrate in water (ice-packs)
solid-gas		paradichloro-benzene in air (mothballs)

Summary:

Solutions are everywhere.  They comprise many of the things that we eat and drink, and many of things we use in our lives from medicines to cleaning supplies to car fluids.  All of the air and water around us is comprised of solutions.  Thus, the study of solutions has become important.

In this lab, many of the common solutions around us were classified according to their solute and solvent phases, and additional examples of solutions were given.  Although we often think solutions are only solids dissolved in liquids, there are many solution combinations, in fact a total of nine possibilities exist. A solution is defined as a homogeneous mixture containing a solute and a solvent.  Homogeneous solutions are formed when both the solute and solvent are in the gas phase (liquid in gas and solid in gas combinations form colloidal dispersions); when a solvent in the liquid phase is combined with a solid, liquid or gas solute; or when a solid solvent is combined with a solid, liquid or gas solute.

TYPES OF SOLUTIONS THAT CAN OCCUR IN OUR BODY BASED ON SOLUTE CONCENTRATION

            Though the cell is the basic unit of all life, there is nothing basic about it. Cells require very specific conditions to be able to function properly. Temperature and the amount of water and nutrients must all be just right in order for a cell to be healthy, and these optimal conditions vary depending on the organism.

The amount of fluid both inside and outside a cell is one condition that is very important, and this fluid amount is often determined by the amount of solutes outside of the cell. Solutes are the particles that are dissolved in a solvent, and together they form a solution. In your body, these solutes are ions like sodium and potassium.

There are three types of solutions that can occur in your body based on solute concentration: isotonic, hypotonic, and hypertonic.

An isotonic solution is one in which the concentration of solutes is the same both inside and outside of the cell.

A hypotonic solution is one in which the concentration of solutes is greater inside the cell than outside of it.

A hypertonic solution is one where the concentration of solutes is greater outside the cell than inside it.

How Different Solutions Affect Your Cells

For the cells in your body, the ideal solution is an isotonic solution. This is because water (which is the major solvent in your body) likes to diffuse from an area of low-solute concentration to an area of high-solute concentration. This process is called osmosis. Water does this because, by diffusing to where there are more solutes, it essentially evens out the ratio of solvent and solute.

When human cells are in a hypotonic solution, water will rush into the cell by osmosis, which is not good for the cell because it will fill with water and burst, or lyse. Plant cells actually prefer hypotonic solutions because they have a rigid cell wall that needs the pressure from extra water inside the cell to stay rigid and firm.

A hypertonic solution will do just the opposite to a cell, since the concentration of solutes is greater outside of the cell than inside. For both human and plant cells, the water will rush out of the cell and it will shrivel up. When this happens to a plant cell it is called a plasmolyzed cell.

PROPERTIES OF SOLUTIONS

What are Colligative Properties?

As we have discussed, solutions have different properties than either the solutes or the solvent used to make the solution. Those properties can be divided into two main groups--colligative and non-colligative properties.

 Colligative properties depend only on the number of dissolved particles in solution and not on their identity. Non-colligative properties depend on the identity of the dissolved species and the solvent.

To explain the difference between the two sets of solution properties, we will compare the properties of a 1.0M aqueous sugar solution to a 0.5 M solution of table salt (NaCl) in water. Despite the concentration of sodium chloride being half of the sucrose concentration, both solutions have precisely the same number of dissolved particles because each sodium chloride unit creates two particles upon dissolution--a sodium ion, Na⁺, and a chloride ion, Cl^-.

 Therefore, any difference in the properties of those two solutions is due to a non-colligative property. Both solutions have the same freezing point, boiling point, vapor pressure, and osmotic pressure because those colligative properties of a solution only depend on the number of dissolved particles.

The taste of the two solutions, however, is markedly different. The sugar solution is sweet and the salt solution tastes salty. Therefore, the taste of the solution is not a colligative property. Another non-colligative property is the color of a solution. A 0.5 M solution of CuSO₄ is bright blue in contrast to the colorless salt and sugar solutions. Other non-colligative properties include viscosity, surface tension, and solubility.

Laboratory Activity

 I Scream, We All Scream for …Colligative Properties!?

Introduction:

When a solute is added to water the physical properties of freezing point and boiling point change. Water normally freezes at 0^oC and boils at 100^oC. As more solute is added, the freezing point drops (“freezing point depression”) and the boiling point increases (“boiling point elevation”). This property is useful in our lives. Antifreeze is added to car radiators to keep the water used as a coolant from freezing in winter and boiling in the summer. Antifreeze is added to a solvent to “de-ice” planes before take-off to prevent ice forming on the wings. When roads ice over in the winter, salt is added along with sand. The sand is used to give cars traction and the salt is used to lower the freezing point of the water so the ice melts. In the kitchen we can use this same property to make ice cream. Water freezes at 0^oC. But milk, although it contains mostly water, normally freezes at about  –4^oC. In order to get the milk to freeze, we add salt to the ice. This changes the melting point of ice to –6^oC. This is low enough to freeze the milk and make ice cream. The purpose of this lab is to use the property of lowering freezing points by adding salt to water to produce ice cream.

The traditional method of ice cream making, using ice and rock salt to freeze a milk and sugar mixture, involves the application of several chemistry principles:

Colligative Properties:  Physical properties determined by the concentration of dissolved particles in a mixture, and not affected by the type of dissolved particles.  Osmotic pressure (tendency of water to flow from high water concentration to low water concentration as seen when applying salt to a slug), boiling point elevation and, most importantly in this activity, freezing point depression are examples of colligative properties. 

When 1 mole of the ionic compound NaCl dissolves in a liter of water it dissociates into 1 mole of sodium ions and 1 mole of chloride ions, or a total of 2 moles of particles per liter of mixture (this is called a 2M or 2-Molar  concentration).  When 1 mole of the covalent compound sugar, C₁₁ H₂₂ O₁₁ , dissolves in a liter of water there is just 1 mole of sugar molecules per liter (onlya 1-Molar concentration) because sugar molecules remain intact when dissolved in water and do not dissociate like ionic compounds do.  The greater the particle concentration (molarity M): the stronger the colligative effect.

Ice cream, a mixture of milk, sugar and vanilla, is really an aqueous mixture with many particles in each liter.  Because of this, the freezing point of ice cream is lower that of water. To freeze ice cream, or to keep it frozen, you must keep its temperature significantly below 0⁰ Celsius.  Adding rock salt to ice produces a melting ice and saltwater mixture, with a depressed freezing point in which the ice cream can be frozen.

Heat of Fusion:  Melting a solid to form a liquid is an endothermic process.  The heat of fusion of a substance is the amount of energy needed to change 1 gram of solid to liquid.  The heat of fusion of water is 334 joules/gram. Since the room temperature and the milk mixture are both warmer than the freezing point of the ice mixture, energy will be transferred into the ice mixture on all sides.  This energy will be used in melting the ice (334 joules per gram melted) and the temperature of the ice mixture will stay at its freezing point as long as there is still ice to melt.  This process will effectively draw energy out of the milk mixture, lowering its temperature and, eventually, freezing it!

Hess' Law and Enthalpy Change:  The energy absorbed by the ice mixture is transferred from the room and the milk mixture so DH ice = DH milk mix + DH room.  This is an application of Hess' Law.

Pre-Lab.

Define COLLIGATIVE PROPERTIES:

Give four examples of COLLIGATIVE PROPERTIES

1.                                             2.                                             3.                                             4.

 Increasing the concentration of a solution will __________ the boiling point of the solution

Increasing the concentration of a solution will __________ the freezing point of the solution

Procedure:  (BE SURE TO RECORD ANY OBSERVATIONS!)

1. Measure 1/2 cup (c) milk and 1/2 c whipping cream; put into a small Zip-lock baggie.

2. Add 1/4 tsp. vanilla to the milk/cream mixture.

3. Measure 1/4 c sugar; add to milk/cream mixture in baggie.

4. Close the bag SECURELY, squeezing most of the air out before sealing.

5. Place the small plastic bag inside a larger Zip-lock baggie.

6. Surround the smaller bag with a few cups of crushed ice.

7. Pour 3/4 to 1 c salt over the crushed ice and seal the larger bag SECURELY.

8. Put on mittens or wrap the bag in a thick towel.  Knead or roll back and forth on the lab table.  Be careful not to put too much pressure on the bag.

9. After 10 minutes, see if the mixture is frozen.  If not, continue kneading.

10. When the mixture is frozen, remove the smaller bag.  Wipe the brine from the zipped edges of the bag. 

11. Enjoy the fruits of your labors.

Analysis:

1.      Calculate  D T_f (change in the freezing point): ____________________________

2.      Using the equation  DT_f = k_f m, calculate  m (molality) ____________________

Remember k_f = 1.86

Post lab Questions:

A.    Why was salt added to the ice? (Why did your milk/cream/sugar mixture freeze?) 

B.     Where did the heat the flow from (What happened to the temperature of the ice after salt was added?)

C.     What is the minimum temperature that pure water can exist as a liquid at standard pressure?

D.    What do you think would happen to the temperature of the ice if you added 6 tablespoons of salt instead of 2 tablespoons?

E.     If there are charged particles that can move around in a substance, that substance is able to conduct electricity.  So, do ionic compounds or covalent compounds conduct electricity when dissolved in water?  Explain.  Use the idea of what happens when these two different types of compounds dissolve in water in your answer.

F.      Which compound types (ionic or covalent) produce more particles when dissolved in water and why?

G.    Answer the following questions:

How many particles are produced when a mole of sugar dissolves? _______

How many particles are produced when a mole of CO₂ dissolves?   _______

How many particles are produced when a mole of sodium chloride (NaCl) dissolves?  _____

How many particles are produced when a mole of copper (II) chloride (CuCl₂ ) dissolves? __

How many particles are produced when a mole of iron (III) sulfate [Fe₂(SO₄)₃] dissolves? __     

H. Which type of compound (ionic or covalent) will have a greater effect on the colligative  properties of a solution?  Explain.  Use the answers to the previous questions in your answer.

I. Why was salt used in the outside bag in this lab?  Be as specific as possible.  Use the  

     idea of colligative properties and the idea of covalent and ionic compounds in your

    answer.

J.  Explain in detail, using the idea of colligative properties, why salt is used to make icy

     roads safe for driving on in the winter.

SOLUBILTY

The solubility of a substance is the amount of that substance that will dissolve in a given amount of solvent. Solubility is a quantitative term. The terms soluble and insoluble are relative.

A substance is said to be soluble if more than 0.1 g of that substance dissolves in 100 mL solvent. If less than 0.1 g dissolves in 100 mL solvent, the substance is said to be insoluble or, more exactly, sparingly soluble.

The terms miscible and immiscible may be encountered when considering the solubility of one liquid in another. Miscible means soluble without limits; for example, alcohol is miscible with water Immiscible and insoluble mean the same; oil is immiscible with water, as in oil and vinegar salad dressing.

A. Determining Solubility 

How is the solubility of a substance determined? A known amount of the solvent--for example, 100 mL--is put in a container. Then the substance whose solubility is to be determined is added until, even after vigorous and prolonged stirring, some of that substance does not dissolve. Such a solution is said to be saturated because it contains as much solute as possible at that temperature. In this saturated solution, the amount of solute is the solubility of that substance at that temperature in that solvent.

Doing this experiment with water as the solvent and sodium chloride as the solute, we find that, at 20°C, 35.7 g of the salt dissolve in 100 mL water. The solubility of sodium chloride is, then, 35.7 g/100 mL water at 20°C. Sodium chloride is a moderately soluble salt. The solubility of sodium nitrate is 92.1 g/100 mL water at 20°C; sodium nitrate is a very soluble salt. At the opposite end of the scale is barium sulfate, which has a solubility of 2.3 X 10 ^-4 g/100 mL water at 20°C. Barium sulfate is an insoluble salt.

B. Saturated Solutions 

A saturated solution is one in which the dissolved solute is in equilibrium with the undissolved solute. The container in figure 11.2a holds a saturated solution of sucrose (cane sugar); at the bottom of the container is some undissolved sucrose. If we could see the individual molecules in this solution, we would see that some of the molecules of sucrose are moving away from the solid crystals at the bottom of the container as they dissolve. The same numbers of molecules are coming out of solution to become part of the undissolved sucrose. 

Dissolution (dissolving) and precipitation are occurring at the same rate, thereby satisfying the requirement for a dynamic equilibrium. We can express the equilibrium in the sucrose solution with the equation                      sucrose(s) sucrose(aq)

Two statements can be made about this solution: (1) the two processes dissolution and precipitation are going on at the same time, and (2) the number of molecules in the solution remains constant.

A saturated solution of an ionic compound is slightly different from that of a covalent compound like sucrose. The ionic compound dissolves and exists in solutions as ions; the covalent compound dissolves and exists in solutions as molecules. The equilibrium of sodium chloride with its ions in a saturated solution would be shown by the equation

NaCl(s)  Na⁺(aq) + Cl^-(aq)

An unsaturated solution contains less solute than does a saturated solution. No equilibrium is present. When additional solute is added to an unsaturated solution, it dissolves. When additional solute is added to a saturated solution, the amount of dissolved solute does not increase, because the limit of solubility has already been reached. Adding more solute to a saturated solution simply increases the amount of undissolved solute.

Remember in discussing these solutions that solubility changes with temperature. A solution that is saturated at one temperature may be unsaturated at a different temperature.

This spot is perhaps appropriate for a few comments on preparing solutions. Dissolution requires interaction between the molecules (or ions) of the solute and the molecules of the solvent. A finely divided solute will dissolve more rapidly than one that is in large chunks, because more contact is made between solute and solvent. Constant stirring increases the rate of dissolution, because stirring changes the particular solvent molecules that are in contact with undissolved solute. Because the solubility of solids and liquids generally increases with temperature, solids and liquids are often dissolved in warm solvents.

FACTORS AFFECTING SOLUBILITY

Forces between Particles

One factor that affects solubility is the nature of the intermolecular forces or interionic forces in both the solute and the solvent. There are kinds of forces operating in solids and liquids. When one substance dissolves in another, the attractive forces in both must be overcome. The dissolving solute must be able to break up the aggregation of molecules in the solvent (that is, the hydrogen bonds between molecules or the dispersion forces between molecules in a nonpolar solvent), and the molecules of the solvent must have sufficient attraction for the solute particles to remove them one by one from their neighbors in the undissolved solute. If the solute is ionic, only a very polar solvent like water provides enough interaction to affect dissolution.

In those ionic compounds like barium sulfate that are called insoluble, the interaction between the ions is greater than can be overcome by interaction with the polar water molecules. If the solute particles are polar molecules, polar solvents such as alcohols can usually affect dissolution. If the solute is nonpolar, it may dissolve only in nonpolar solvents, not because polar solvent molecules are unable to overcome the weak dispersion forces between the solute molecules, but because these dispersion forces are too weak to overcome the dipole-dipole interaction between the solvent molecules.

A general rule is that like dissolves like. Ionic and polar compounds are most apt to be soluble in polar solvents like water or liquid ammonia. Nonpolar compounds are most apt to be soluble in nonpolar solvents, such as carbon tetrachloride, and hydrocarbon solvents like gasoline.

The solubility of gases in water depends a great deal on the polarity of the gas molecules. Those gases whose molecules are polar are much more soluble in water than are nonpolar gases. Ammonia, a strongly polar molecule, is very soluble in water (89.9 g/100 g H₂O); so is hydrogen chloride (82.3 g/100 g H₂O). Helium and nitrogen are nonpolar molecules. Helium is only slightly soluble (1.8 X 10^-4 g/100 g H₂O), as is nitrogen (2.9 X 10^-3 g/100 g H₂O).

Table shows specific examples of different kinds of compounds and their relative solubilities in water, a polar solvent; in alcohol, a less polar solvent; and in benzene, a nonpolar solvent.

Solubility and Interparticle Bonds
		Solubility in
Kinds of bonds	Example	Water	Alcohol	Benzene
ionic	sodium chloride	very soluble	slightly soluble	insoluble
polar covalent	sucrose (sugar)	very soluble	soluble	insoluble
nonpolar covalent	napthalene	insoluble	soluble	very soluble

Temperature 

The next table shows the solubility of several substances in water at 20°C and 100°C. Notice that the solubilities of solids and liquids usually increase as the temperature increases, but the solubilities of gases decrease with increasing temperature. This property of gases causes our concern for the fish population of lakes, oceans, and rivers threatened with thermal (heat) pollution. Fish require dissolved oxygen to survive. If the temperature of the water increases, the concentration of dissolved oxygen decreases, and the survival of the fish becomes questionable.

Solubility and temperature
		Solubility (g/100mL water)
Compound	Type of bonding	At 20°C	At 100°C
sodium chloride	ionic	35.7	39.1
barium sulfate	ionic	2.3 X 10-4	4.1 X 10-4
sucrose	polar covalent	179	487
ammonia	polar covalent	89.9	7.4
hydrogen chloride	polar covalent	82.3	56.1
oxygen	nonpolar covalent	4.5 X 10-3	3.3 X 10-3

The solubility of solids and liquids is closely related to the heat of solution.

Heat of solution – the heat involved when a solute dissolves to give a saturated solution. Exothermal or exothermic - Heat is evolved to the surrounding and the heat of the solution is negative. Endothermal or exothermic - Heat is absorbed from the surroundings and the heat of the solution is positive.

Pressure 

The pressure on the surface of a solution has very little effect on the solubilities of solids and liquids. It does have an enormous effect on the solubility of gases. As the partial pressure of a gas in the atmosphere above the surface of a solution increases, the solubility of that gas increases. A carbonated beverage is bottled and capped under a high partial pressure of carbon dioxide so that a great deal of carbon dioxide dissolves. When the bottle is uncapped, the partial pressure of carbon dioxide above the liquid drops to that in the atmosphere, the solubility of carbon dioxide decreases, and the gaseous CO₂ comes out of solution as fine bubbles.

This same phenomenon is of concern in scuba diving. As divers descend, they are under ever-increasing pressure due to the increasing mass of water above them. Each 33 ft of water exert a pressure of 1 atm. As the diver descends, the pressure on the air in the lungs increases rapidly, and the blood dissolves more than normal amounts of both nitrogen and oxygen. As the diver comes up, the pressure decreases and the gases come out of solution. If the change in pressure is too swift, the bubbles of gas are released into the blood stream and tissues, instead of into the lungs. The bubbles cause sharp pain wherever they occur, most notably in the joints and ligaments. This dangerous and sometimes fatal condition is called bends, the name deriving from the temporary deformities caused by the affected diver's being unable to straighten his or her joints. Because helium, at all pressures, is less soluble than nitrogen, divers often breathe a mixture of helium and oxygen instead of air, a nitrogen-oxygen mixture.

The term hyperbaric means greater-than-normal atmospheric pressure. In hyperbaric medical treatments, patients are subjected to a pressure greater than atmospheric. This pressure increases the amount of oxygen dissolved in the blood. Treatment in hyperbaric, or high-pressure, chambers is of particular value for patients who are suffering from severe anemia, hemoglobin abnormalities, or carbon monoxide exposure or undergoing skin grafts, for such treatment increases the amount of oxygen transported by the blood to the tissues.

Lab Activity: Preparing Saturated Solutions & Measuring Solubility

Objective:  You will prepare a saturated solution of two substances and measure the solubility (concentration of the solute when it is saturated).

Substances: Copper Sulfate (CuSO₄) and one other assigned substance

Directions:

Day #1: Prepare Saturated Solution

1.      Measure & record the mass of an empty evaporation dish.

2.      Put 10.0 mL of water in a test tube.

3.      Add 1 scoop of copper sulfate.

4.      Shake the test tube until all of the solid has dissolved.  If all is dissolved, proceed to step # 7.

5.      Add another scoop of copper sulfate, and shake until it has all dissolved.  If all is dissolved, proceed to step # 7.

6.      Continue to add scoops and shake until you have prepared a solution that is beyond saturated.  It is beyond saturated when you can see solid particles on the bottom.  Record the number of scoops in your lab notebook.

7.      The liquid portion of your solution is saturated.  Decant (pour out the liquid portion) into a graduated solution and record the volume.

8.      Pour the solution into an evaporation dish, label and let sit out for the night.

9.      Repeat this procedure with the other assigned substance.

Day #2: Measuring and Calculate Solubility.

1. If all the liquid has evaporated, measure the mass of the evaporation dish and the dry solid. 
2. If there is still some liquid left, gently heat the solution until all of the liquid has evaporated.
3. Calculate the mass of the dry solid and calculate solubility.

Data table #1: Group Data -  mass & volume

Substance	Mass empty evaporation dish	Volume of Solution	Mass dish & dry solid

Data table #2: Concentration & Solubility

Substance	Mass Dissolved Solid	Volume Solution	Concentration (m / v)	Solubility (m / v) x 100

 Measurements & Calculations:

1.      Measure the volume of saturated solution in a graduated cylinder.

2.      Calculate the mass of the dissolved solid: mass of crucible with dry solid – mass of empty crucible.

3.      Calculate the concentration: divide the mass of the solid dissolved by the volume of the water boiled away

4.      Calculated the solubility. Multiply the concentration of the saturated solution by 100.  In other words, if you had used 100 mL of water, what would be the maximum amount of this solid you could have dissolved?

Data table #3: Class data comparing solubility

	Sodium Chloride (NaCl)	Sodium Nitrate (NaNO3)	Glucose (C6H12O6)
Group #1
Group #2
Group #3
Group #4
Group #5

FACTORS AFFECTING THE RATE OF DISSOLUTION

The dissolution process of a solute in a solvent takes place at a certain rate. This rate of dissolution is usually dependent on factors that can alter this rate. These factors can be modified to increase or decrease the rate of dissolution to a certain extent. The effects of these factors on dissolution can be described in three ways.

Dissolution: A Chemical Process

Dissolution is chemical phenomena in which the solvent and solute act as reactants and the solution so formed is the product of the reaction. Like any chemical reaction, dissolution is also affected by factors like temperature, surface area of the reactants and agitation or turbulence in the solution.

Temperature

The temperature at which dissolution process takes place also affects the rate of dissolution. Increase in temperature accelerates the rate at which dissolution takes place. During dissolution, when the temperature is increased, more collisions take place between the particles of the reactants and, hence, the rate of dissolution increases.

Agitation

Agitation produced by stirring or mixing in a solution increases the rate if dissolution process. Agitation increases the chemical reaction rate by increasing the rate of diffusion in the solution. For example; a teaspoon of salt dissolves faster in water when stirred, as compared to the salt left    unstirred.

Surface Area of Solute

As the surface area of the solute increases, the rate at which the solute dissolves into the solvent also increases. More surface area of solute particles enables the solvent to react quickly with the solute particles, increasing the rate of dissolution.

Laboratory Activity - Rates of Dissolving

Problem: What affects the rate at which a solute dissolves?

Hypothesis: _______________________________________________________________

_______________________________________________________________________________________________________________________________________________

Materials:

 8 clear plastic cups and Labeled A, B, C, D, E, F, G, H          100 mL graduated cylinder

 8 sugar cubes  3 paper towels or pieces of paper                       Stirring rod

 500 mL beaker  Hot plate                                             Water

 Methods:

1. Fill the 500 mL beaker with water and begin heating on the hotplate. The water needs to

   be warm/hot, but not boiling.

 2. Get 8 sugar cubes. Crush 4 of the sugar cubes.

 3. Measure and pour 100 mL hot water in cups A and B. Measure and pour 100 mL cold

     water in cups C through F.

 4. One at a time, add sugar to each cup. Follow the instructions below for each cup.

     Carefully observe the time it takes for the sugar to dissolve completely in each cup of

      water. Record your observations.

A. Crushed cube in hot water

B. Whole cube in hot water

C. Crushed cube in hot water, stir until sugar is completely gone

D. Whole cube in cold water, stir until sugar is completely gone

E. Crushed cube in cold water, stir until sugar is completely gone

F. Whole cube in cold water, stir until sugar is completely gone

G. Crushed cube in cold water

H. Whole cube in cold water

 5. When all observations have been recorded, empty and rinse each cup.

 6. Rate the sugar samples from fastest to slowest in dissolving. Give the fastest dissolving

     sample a rating of 1. The slowest dissolving rate should be rated 6. Record your ratings.

Safety:

1. ______________________________________________________________________

2. ______________________________________________________________________

Independent Variable: ______________________________________________________

Dependent Variable: ________________________________________________________

Controls: ___________________________________________

Data and Observations

Cup                    Sugar Sample              Water Conditions        Time (s)           Rating

A                        Crushed                                   Hot

B                        Cube                                        Hot

C                        Crushed                                   Hot, Stir

D                        Cube                                        Hot, Stir

E                        Crushed                                   Cold, Stir

F                        Cube                                        Cold, Stir

G                        Crushed                                   Cold

H                        Cube                                        Cold

Things to write about in your discussion:

1. Explain how particle size affects the rate at which sugar dissolves in water?

2. Explain how temperature affects the rate at which sugar dissolves.

3. Explain how stirring affects the rate at which sugar dissolves.

Discussion: _________________________________________________________________________

__________________________________________________________________________________________________________________________________________________

Conclusion: _________________________________________________________________________________________________________________________________________________

CONCENTRATION OF SOLUTIONS

The relative amount of solute in a given amount of solvent or solution

Percent Composition (by mass)

We can consider percent by mass (or weight percent, as it is sometimes called) in two ways:

• The parts of solute per 100 parts of solution.
• The fraction of a solute in a solution multiplied by 100.

We need two pieces of information to calculate the percent by mass of a solute in a solution:

• The mass of the solute in the solution.
• The mass of the solution.

Use the following equation to calculate percent by mass:

Molarity

Molarity tells us the number of moles of solute in exactly one liter of a solution. (Note that molarity is spelled with an "r" and is represented by a capital M.)

We need two pieces of information to calculate the molarity of a solute in a solution:

• The moles of solute present in the solution.
• The volume of solution (in liters) containing the solute.

To calculate molarity we use the equation:

Molality

Molality, m, tells us the number of moles of solute dissolved in exactly one kilogram of solvent. (Note that molality is spelled with two "l"'s and represented by a lower case m.)

We need two pieces of information to calculate the molality of a solute in a solution:

• The moles of solute present in the solution.
• The mass of solvent (in kilograms) in the solution.

To calculate molality we use the equation:

Mole Fraction

The mole fraction, X, of a component in a solution is the ratio of the number of moles of that component to the total number of moles of all components in the solution.

To calculate mole fraction, we need to know:

• The number of moles of each component present in the solution.

The mole fraction of A, X_A, in a solution consisting of A, B, C, ... is calculated using the equation:

To calculate the mole fraction of B, X_B, use:

SAMPLE PROBLEMS

þ  What is the molality of a solution prepared by dissolving 0.385 g of cholesterol, C₂₇H₄₆O in

40.0 g of chloroform, CHCl₃?  Cholesterol = 386.7 amu  ;  chloroform = 119.4 amu

              

 What is the mole fraction of cholesterol    in the solution?

þ  What is the molality of a solution prepared by  dissolving 0.385 g of cholesterol, C₂₇H₄₆O in   40.0 g of chloroform, CHCl₃? 

Cholesterol = 386.7 amu  ;  chloroform = 119.4 amu 

                 moles cholesterol                0.385g / 386.7

Molality = ----------------------------  = -----------------------  = 0.0249 mol/kg

                        kg chloroform                         0.04 kg

What is the mole fraction of cholesterol in the solution?

                             0.385g / 386.7

X_chloroform = ------------------------------------------  =  2.96 x 10^-3

                              0.385g/386.7 + 40/119.4

þ  Assuming that seawater is an aqueous solution of NaCl what is its molarity?

  The density of seawater is 1.025 g/mL at 20C and the NaCl concentration is 3.50 mass %

þ  Assuming that seawater is an aqueous solution of NaCl what is its molarity?  The density of seawater is 1.025 g/mL at 20C and the NaCl concentration is 3.50 mass %

              3.5 mass% = 35 g NaCl in 1 kg solution

                     1 kg solution = m/d = 1000/1.025 = 975.6 mL

                    

    35 g / (23+35.5)                    0.598

Molarity =  ----------------------------- = ------------- = 0.61 mol/L = 0.61 M

                        0.9756 L                            0.9756

þ  What is the mass % concentration of a saline sol’n prepared by dissolving 1.00 mol of NaCl in 1.00 L of water? Density_H2O=1.00 g/mL  MM_NaCl = 58.443 g/mol

þ  Assuming that seawater is an aqueous solution of NaCl what is its molarity?  The density of seawater is 1.025 g/mL at 20°C and the NaCl concentration is 3.50 mass %

Assume 1 L to make easier, 1000 mL

þ  What is the molality of a solution prepared by dissolving 0.385 g of cholesterol, C₂₇H₄₆O in 40.0 g of chloroform, CHCl₃?  What is the mole fraction of cholesterol in the solution?

þ  What mass in grams of a 0.500 m solution of sodium acetate, CH₃CO₂Na, in water would you use to obtain 0.150 mol of sodium acetate?

þ  The density at 20°C of a 0.258 m solution of glucose in water is 1.0173 g/mL and the molar mass of glucose is 180.2 g/mol.       What is the molarity of the solution?

þ  The density at 20°C of a 0.500 M solution of acetic acid in water is 1.0042 g/mL.  What is the concentration of this solution in molality?  The molar mass of acetic acid, CH₃CO₂H, is     60.05 g/mol. Assume 1 L

þ  1.36 g of MgCl₂ are dissolved in 47.46 g of water to give a solution with a final volume of 50.00 mL. Calculate the concentration of the solution in mass %, ppm, mole fraction, molarity, and molality.

þ  Hydrochloric acid is sold as a concentrated aqueous solution.  The concentration of commercial HCl is 11.7 M and its density is 1.18 g/mL.  Calculate the mass percent of HCl in the solution.  Calculate the molality of the solution. 

þ  Concentrated sulfuric acid has a density of 1.84 g/mL and is 18 M.  What is the mass % H₂SO₄ in the solution?  What is the molality of H₂SO₄ in the solution?

SOME APPLICATIONS OF SOLUTIONS

a.       A common application of this effect in some parts of the world is in the use of antifreeze solutions in the cooling systems of automobiles in cold climates. "Antifreeze" compounds are usually organic liquids that are miscible with water so that large freezing point effects can be attained.

b.      Knowledge about the right concentration of solute will help us to prevent diseases such as gout. Excessive or too much uric acid in the plasma

c.       Laboratories prepared different  solutions in varied concentrations depending on what is in need ( dilution )

d.      Understanding the properties of solution will help us in our day to day life. Example adding  salt to an ice to prevent it from melting right away

e.        Normal Saline is given to increase the amount of fluid in the blood vessels (intravascular space), without significantly changing the balance of electrolytes in the body.  This is useful for making sure a patient remains in a well-hydrated state of homeostasis.

Why should athletes drink solutions that are isotonic to body fluids rather than ones that are hypotonic?

Cells have a tendency to try and reach equilibrium--- the same amount of water outside the cell as inside. If a cell absorbs too much water from its surroundings in effort to reach equilibrium, it may explode which can be very dangerous even deadly for the athlete. By drinking isotonic solutions, their cells won't need to absorb huge amounts of water and therefore won't explode.