Solutions

1. Solutions

2. Mixtures • Heterogeneous mixture: substances in mixture are not spread uniformly throughout mixture. • Homogeneous mixture: components uniformly mixed in solution • The smaller the particles the moreuniform the mix.

3. Solutions: • Homogeneous mixture of two or more substances, the composition of which may vary within limits. • Consists of two parts: • Solute: (solid, liquid, gas) dissolved substance or substance present in the smaller amount. • Solvent: (solid, liquid, gas) substance that the solute is dissolved in –dissolving medium or substance present in the larger amount.

4. Solutes and solvents • Usually substances are two phases. Solute changes phase when placed in solvent. • Examples: • Aqueous: solvent is water salt in water NaCl in water or NaCl (aq) NaCl is solute and water is solvent Tincture: solvent is alcohol (C2H5OH) tincture of iodine iodine is solute and alcohol is solvent

5. Rate of Dissolving • The quantity solute dissolved per unit of time. • How fast a solute dissolves • Factors that affect the rate of dissolving: (think of a sugar cube in coffee) • Surface area (powder is faster) • Agitation (stir) • Temperature (hot faster than cold)

6. Solubility • Quantity (mass) of solute which can be dissolved in a given volume of solvent at equilibrium, under specified conditions of temperature and pressure. • Table F tells you if something is soluble in water • Table G tells you how much at a particular temperature. • What factors affect the solubility of a substance?

7. Factors that Affect Solubility • Nature of solute and solvent: “Like dissolves like” • Polar and ionic solutes are soluble in polar solvents. (ex water) • Nonpolar solutes are soluble in nonpolar solvents. • Note: Alcohols are soluble in both polar and nonpolar solvents. But ionic solutes are insoluble in alcohols.

8. Factors that Affect Solubility • Temperature: • For Ionic solids: as T solubility example: Jell-O in boiling water • For gases: as T solubility  example: warm soda goes flat

9. Factors that Affect Solubility • Pressure: • For solids/liquids: as P changes, solubility does not change • For gases: as P solubility  (Effervescence: escape of gas from solution)

10. How many types of solutions are there?Give an example of each type? • There are nine different types of solutions. • Solid in solid: Alloys (ex: brass mixture of Cu/Zn) • Solid in liquid: Seawater • Solid in gas: Soot in air • Liquid in Solid: Hg on copper • Liquid in liquid: Alcohol in water • Liquid in gas: fog • Gas in solid: Hydrogen on platinum • Gas in liquid: Carbonated beverage • Gas in gas: Air

11. There are three common solutions. Type 1: Gases in Liquids: • In a closed system an equilibrium exists between the gas dissolved in the liquid and the undissolved gas above the liquid. • The equilibrium is affected by temperature and pressure. • An increase in temperature decreases the solubility. • An increase in pressure increases the solubility. (Henry’s law: the mass of a gas which dissolves in a liquid at a given temperature is proportional to the partial pressure of the gas over the solution)

12. Type 2: Liquids in liquids • There are other liquid solvents besides water. • Some liquids mix together well while others do not. • Miscible: Liquids that are soluble in one another. Mix well together. Ex: Gasoline and oil, water and alcohol • Immiscible: Liquids that are insoluble in each other. Do not mix well together. Ex: Oil and water (oil floats on top)

13. Type 3: Solids in liquids • A solution equilibrium exists when the opposing processes of dissolving the solute in the solvent and of crystallizing the solute from the solvent occur at equal rates. • At this point no further solute can be dissolved and the solution is known as saturated. • General rule: the solubility of a solid increases as the temperature increases. (table G)

14. Looking at solubility • Solubility curves show the relationship of grams of solute that may be dissolved in a solvent at various temperatures. • The solubility curves on Table G in your reference table show the number of grams of a substance that can be dissolved in 100 grams of water at temperatures between 0oC and 100oC. • Each line represents the maximum amount of a substance that can be dissolved at a given temperature.

15. Table G: Solubility Curves • All of the lines that show an increase in solubility as temperature increases represent solids being dissolved in water. • Three lines show decreasing solubility with increasing temperature. These three lines represent gases NH3, HCl, and SO2. • Remember the solubility of all gases decreases with increasing temperature.

16. Concentration • The concentration of a solution may be expressed in a variety of ways • Concentrated and dilute: because the terms are vague they are used for comparison only. • Concentrated: contains a relatively large amount of solute. • Dilute: contains a relatively small amount of solute • For example: orange juice made from concentrate

17. Unsaturated, Saturated and Supersaturated • When reading the solubility curves on table G you will need to recognize three positions with respect to the line of maximum solubility. • When a solution holds less solute than the maximum it can hold it is said to be unsaturated. In this case the amount dissolved will be below the line of solubilty. • When a solution contains the maximum amount of solute that will dissolve at a specific temperature it is saturated. In this case the amount dissolved will be directly on the line of solubility. Solution equilibrium. • When a solution contains more than the maximum amount of solute that will dissolve at a specific temperature it is supersaturated. In this case the amount dissolved will be above the line of solubility. Made from a saturated solution at a higher temperature and then cooling it. Supersaturated solutions are unstable.

18. Table G Supersaturated Saturated Unsaturated

19. Recognizing Unsaturated, Saturated, and Supersaturated • One method of recognizing the type of solution is if it contains some undissolved solute, it must be a saturated solution. • The addition of more solute crystals can also determine the conditions. • If it dissolves, the original solution was unsaturated. • If it simply falls to the bottom of the container then it is saturated. • If it causes additional crystals to form, the original solution was supersaturated.

20. Concentration of Solution • The concentration of a solution is a measurement of the amount of solute dissolved in solution. • There are several ways of expressing the specific concentration of solute in a solution. • Percent by Mass • Percent by Volume • Parts per million • Molarity • Molality

21. Percent by Mass & Percent by Volume • Percent by massis simply the mass of the solute divided by the total mass of the solution, expressed as a percentage (x100). Use a % sign as the unit. • Percent by volumeis simply the volume of the solute divided by the total volume of the solution, expressed as a percentage (x100). Use a % sign as the unit. Mass of Part (solute) Mass of Whole (solution) Percent by mass = X 100 Volume of part (solute) Volume of whole (solution) Percent by volume = X 100

22. Parts Per Million • Parts per million(ppm) is similar to percent by mass because it compares masses. It represents the ratio between the mass of a solute and the total mass of a solution. Instead of multiply by 100 you multiply by 1,000,000. • ppm is often used to measure concentrations of solutes that are present in very small amounts. For example if you wanted to measure the concentration of chloride ions in tap water. • The units for parts per million are ppm • Table T Mass of Solute Mass of Solution X 1,000,000 Parts Per Million =

23. Molarity • Molarity measures the concentration of a solution in terms of moles of solute in a given volume of solution. • The Molarity (M)of a solution is the number of moles of solute in 1L of solution. • When calculating the Molarity of a solution you may need to make conversions before you solve for and answer. • For example if you are given the mass of the solute you will need to convert to moles using the mole calculation equation on Table T. • You may also need to convert the volume of solution given if it is not in liters. • The units for molarity are mol./L or M for short Moles of Solute (mol) Liters of solution (L) Molarity (M) =

24. Molarity by Dilution • Most acids are purchased from laboratory supply houses in concentrated form. • If you want to make a different concentration of acid use the formula below: • M1V1 = M2V2 Where: M1=initial concentration V1= initial volume M2= final concentration V2= final volume

25. Molality • Molality is the number of moles of solute per kg pf solvent. • m= n/kg

26. Helpful Hint about Concentration Problems. • Remember that a solution is made up of the solute and solvent combined. • So if you are performing a concentration problem you may need to add the masses or volumes of your “solute” and “solvent” to solve for the mass or volume of the “solution”, which is what you need for all concentration equations. • If you are given the mass or volume of the “solution” then you do not need to worry about adding the solute and the solvent. Solution = Solute + Solvent

27. Colligative Properties • Properties that depend on the number of particles rather than the nature of particles are called colligative particles. • For example: Boiling point, freezing point, osmotic pressure and vapor pressure.

28. Non-electrolytes • Non-electrolytes are molecular substances that do not break up into ions (i.e. sugar, alcohols CxHyOH) • Do not conduct an electric current (no mobile ions) • Non-electrolytes have a dissociation factor (d.f.) of 1. • C12H22O11 C12H22O11 (aq) 1 mole 1 mole

29. Electrolytes • Electrolytes are ionic substances that break up into ions when put in solution. • For example: acids (H-ion), bases (metal-OH) and salts (+ ions and – ions) • Conduct electricity due to mobile ions • NaCl  Na+ + Cl- 1 mole NaCl  1mole Na++ 1 mole Cl- The dissociation factor is equal to 2. The more moles of ions produced, the greater the effect on colligative properties.

30. Freezing Point Depression • Freezing point is lowered when a non-volatile substance (non-electrolyte) is added to a solvent. • When one mole of a nonelectrolyte is added to 1 kg of water the freezing point is lowered by 1.86 degrees celcius. • For electrolytes the more particles the lower the freezing point. • Example: salt on walkways in icy conditions, making ice cream • Change in f.p. = m x d.f. X 1.86 m=molality and d.f =dissocation factor

31. Compare the following: • Which will have the lowest freezing point? • Which will have the highest freezing point? 1 mole of MgCl2 in 500 g water 1 mole of C6H12O6 in 500 g water Freezing point will always be lower than 0°C.

32. Boiling Point ELevation • Boiling point is raised when a non-volatile substance (non-electrolyte) is added to a solvent. • When one mole of a nonelectrolyte is added to 1 kg of water the boiling point goes up by .52 degrees celcius. • For electrolytes the more particles the higher the boiling point. • Example: water boils at a higher temp when salt is added • Change in b.p. = m x d.f. x .52 where m=molality and d.f. =dissociation factor

33. Compare the following: • Which will have the highest boiling point? • Which will have the lowest boiling point? 1 mole of MgCl2 in 500 g water 1 mole of C6H12O6 in 500 g water Boiling point will always be higher than 100°C.