Understanding Solutions: Properties, Solubility, and Separation Methods

1. Properties of Solutions

2. Changes In State

3. A Review of Matter • Types of matter • Pure substances • Elements – cannot be broken down by chemical means • Compounds – two or more elements combined in definite proportions • Mixtures • Homogeneous – evenly mixed, uniform throughout • Heterogeneous – unevenly mixed, varied composition

4. Increasing Surface Area Increases Dissolution

5. What is a Solution? • Homogeneous mixture of substances in the same state. • A solution may be; • A solid n solid. Ex. Zn + Cu  brass (alloy) • Gas n gas. Ex. Air. • Liquid n liquid. Ex. alcohol + water • And, most familiar to you, a substance, usually a solid, dissolved in water.Ex. NaCl + water.(sea water)

6. Components of a Solutions • A solute is the substance being dissolved. • The solvent is the substance that exists in the greater amount, into which the solute is mixed. • Ex. In sea water, salt is the solute, while water is the solvent. • Water’s properties make it a universal solvent.

7. The Secret To Waters Dissolving Power is it’s Polarized Nature The polar water molecules surround the solute molecules, forming H-bonds and dipole attractions

8. Aqueous Solutions • Solutions of a substance dissolved in water are termed aqueous and are written as follows; • NaCl (s)  Na+ (aq) + Cl- (aq) (where s = solid, aq = aqueous, and g = gas) • The particles of a homogeneous aqueous solution are very small, will not settle and will pass through a filter.

9. Mixtures Can Be Separated

10. Properties of Mixtures • Light will pass directly through a solution. • The light beam will not be visible in the body of the solution. • Light will be scattered by the larger particles of heterogeneous solutions. • This is called the Tyndall Effect, after the scientist who discovered it.

11. Colloids and Suspensions Scatter Light. Solutions Do Not.

12. Summary of Solution Properties • Aqueous solutions are homogeneous mixtures • They are clear and do not disperse light. • Can have color (ex. CuSO4 is blue). • Particles will not settle when left standing. • Particles pass readily through a filter.

13. Solubility • Ability of a substance to dissolve in another substance. • “Like dissolves like.” • Polar molecules dissolve in polar substances • Non-polar (lipids, oils) dissolve in non-polar. • All materials do not dissolve in equal concentrations.

14. Solubility • Highly soluble materials are said to be soluble or miscible. • Materials that dissolve very little are said to be insoluble or immiscible. • A material dissolves because the attraction of the solvent molecules is greater than the attractive force holding the solute molecules together. • Insoluble materials have a greater affinity for their own molecules. • These last 2 points is why oil pools together in water

15. Temperature and Pressure • Solubility is relative to changing temperature. • Generally, solubility increases with increasing temperature. • Pressure has NO EFFECT on the solubility of solids and liquids (that compressibility thing again) ONLY GASES. • Gases behave oppositely, they are more soluble athigher pressure and lower temperature (open a warm seltzer bottle and a cold one at the same time).

16. Gas Mixtures React Opposite to Liquids in Terms of Pressure and Temperature

17. Solubility Graphs • We can look at a graph of solubility curves to determine how much solute can dissolve in a specific quantity of solvent. • We need to know; • Which solute • What temperature • See reference tables F and G.

18. Table F • Guidelines for solubility (Table F) • Table F is broken down into Ionic compounds that are soluble. Try learning the exceptions, there are fewer and you’ll know the rules by default. • There are many Ionic compounds that are insoluble and many exceptions (use the table).

20. Table G • Graph of solubility curves • Broken down into; • X axis = Temperature in Celsius • Y axis = amount of solute per 100g H2O. • How to read graph • Trace given info, grams or °C, along curve of X and Y axis to find desired data point for given solute.

22. Saturation • There is an upper limit to how much solute a given solvent can hold at a certain temperature. • When holding the maximum amount of solute the solution is said to be saturated. • When a solution is saturated the addition of any more solute will cause a precipitate to form. • Ex. Have you ever found un-dissolved substance on the bottom of your coffee or ice tea?

23. Using Table G (solubility curves) To Predict Solubility of Substances • Find KNO3 (potassium nitrate) in table G. • At approximately 32°C, 100 grams H20 can hold a maximum of ~50grams of KNO3. It is saturated at this mass and temperature. • If the temperature was raised to 50°C, 100g of solvent could hold up to ~83 grams of KNO3. • One could add an additional 33 grams of solute to re-saturate the solution. • If you do not add more solute, the solution would be called unsaturated.

24. Supersaturation • This occurs when a solution is heated, saturated and then slowly cooled, creating a new saturation point. • The solution maintains it’s homogeneous appearance. • Supersaturated solutions are very unstable. • Addition of a single crystalline molecule will cause the solution to form a precipitate.

25. Determining Saturation • Visual inspection: Does the solution contain un-dissolved particles? • Addition of more solid solute: • Does the solid fall to the bottom? • Yes  solution is saturated • No  solution is not saturated • Does a precipitate form (material comes out of solution)? • Yes  solution is supersaturated • No  return to (1)

26. Concentration of Solutions • Compositions of homogeneous mixtures can vary. • Two terms are commonly used; • Dilute • Concentrated • The above terms are relative, that is, they are not very precise and tell little about how much solute is present. • More precise terms are; molarity, % mass, % volume & ppm (parts per million).

27. Molarity • Molarity is the number of moles of solute per liter of solution. M = # mol solute L solution • Molarity can be used as a conversion factor to find Volume and # of mols. M (mol/L) x V (L) = # of mol

28. Percent by Mass • Mass of a certain ingredient divided by the total mass. % Mass = Mass of part x 100 Mass of Whole

29. Parts Per Million • PPM is used when very small amounts are present. • Generally used to report acceptable levels of toxins, pollutants, poisons or other trace amounts of substances. PPM = grams of solute x 106 grams of solution

30. Dilution • The process of preparing a less concentrated solution from a more concentrated one. M1 V1 = M2 V2 • Where M1 and M2 equal the old and new molarities and… • V1 and V2 equal the old and new volumes.