1. Solutions

2. Pure Substances & Mixtures o What's the matter? o Pure Substances have a definite set of physical properties. o (melting point, density, reactivity, etc) o Elements – cannot be broken down o Compounds – can be broken down o Mixtures are two or more substances physically combined but not chemically combined o Heterogeneous- a mixture that is not uniform throughout. o Looks opaque – particles are large enough to block light o Examples: dirt, chocolate chip cookies, milk, you o Homogeneous- uniform throughout o Solution- a special homogeneous mixture where one substance is dissolved into another. Almost always clear. o Examples: salt water (ionic compound + water), syrup, rubbing alcohol Everything! (har har)

3. Mixtures Activity: 1. Left-hand side of your notebook, draw 4 gas tanks and label them “H 2 Gas,” “O 2 Gas,” “Hydrogen & Oxgen compound,” and “H 2 & O 2 Mixture.” 2. Fill each tank with the appropriate molecules. Use open circles for hydrogen atoms and filled-in circles for oxygen atoms. H 2 GasO 2 Gas O 2 & H 2 Mixture Oyxgen & Hydrogen Compound Pure Not pure

4. Separating Mixtures Separating Mixtures: distillation – using different boiling points to separate substances in a mixture Examples: desalination filtration – Using particle size to separate substances through a semi-permeable membrane. Examples: air filter in a car, tea bag, coffee filter Chromatography- using a molecule’s ability to move through a medium. (demo) Examples: gel electrophoresis, HPLC

5. Solutions Solutions have two components: Solvent – major component of a solution, dissolves the solute Solute – minor component of a solution,dissolved by solvent The solute is the active ingredient of a solution. Solutions used in chemical equations have (aq) added aq = aqueous Steps of solution process: Known as solvation. (video) Solute particles separate – energy is absorbed in order to separate solute particles Solvent molecules separate – energy also absorbed Solvent particles surround solute particles – energy is released Solutions that release more energy than absorbed are called exothermic (energy EXiting). Solutions that absorb more energy than released are called endothermic (energy ENters)

6. Solubility Solubility- able to be dissolved Solids, liquids and gases can be dissolved. “Like dissolves Like” Ionic & polar solutes dissolve in polar solvents. Non-polar solutes dissolve more in non-polar solvents. Immiscible - liquids that are not soluble in each other ● Ex: Oil and water do not dissolve in each other because of different polarities: Water is polar and oil is nonpolar

7. Solubility  Not all ionic compounds are soluble in water  Insoluble compounds will fall out of solution, forming a precipitate  Solubility for ionic compounds can be predicted using solubility rules

8. Temperature & Pressure Temperature and solubility Solids dissolve better at higher temperatures Faster solvent particles allow solute particles more movement Example: hot tea can dissolve more sugar than cold tea Gases do not dissolve well at higher temperatures Gas molecules tend to disperse with higher temperatures Example: soda bottle at room temperature will be less fizzy than when cold Pressure and solubility Pressure has no effect on solid solutes High pressure allows gases to dissolve better High pressure does not allow gas molecules to disperse easily Example: sealed soda bottles stay fizzy longer

9. Saturation Saturated solution- solution containing the maximum amount of solute for a specific amount of solvent at a given temperature and pressure. More solute will not dissolve Unsaturated solution- does not contain the maximum amount of solute in solution. More solute can dissolve until the maximum is reached Supersaturated solution- an unstable solution that contains more solute than it can theoretically hold at a given temperature Created by cooling a saturated solution Adding more solute causes the excess solute to precipitate Precipitate- a solid that forms and settles down in a liquid mixture

10. Solubility Curves Solubility Curve– graph of solubility versus temperature compares the solubilities of multiple compounds in water as a function of temperature. solubility is expressed in terms of grams of solute per 100 grams of H 2 O.

11. Solubility Curves How many grams of KNO 3 can dissolve in 100 grams of water at 30ºC? Answer: About 48 grams.

12. Solubility Curves How many grams of KClO 3 can dissolve in 100 grams of water at 55ºC? Answer: About 25 grams.

13. Solubility Curves How many grams of NH 4 Cl can be dissolved in 200 grams of water at 40ºC? Answer: About 96 grams.

14. Solubility Curves A saturated solution of NaNO 3 is prepared at 70ºC in 100 grams of water. If the temperature is lowered to 40ºC, how many grams of NaNO 3 will precipitate? Answer: About 30 grams will precipitate out.

15. Solubility Curves If 35 grams of NH 4 Cl are dissolved in 100 grams of water at 30ºC, is the solution unsaturated, saturated, or supersaturated? Answer: The solution is unsaturated.

16. Solubilities of Gases Henry’s Law – The solubility of a gas in a liquid increases as the pressure above the liquid increases. Explains decompression sickness (the “bends”). Nitrogen is forced to dissolve in scuba divers’ blood by high pressure. If the diver ascends too quickly, the nitrogen bubbles come out of solution and form painful and possibly fatal gas embolisms.