1. Part I: Basic properties of pure water.

2. Water Molecule Triatomic (3 atoms) Bent Shape (104.45°) Polar Molecule δ+δ+ δ+δ+ δ-δ-

3. δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ+δ+ δ-δ-

4. Hydrogen Bonding The only Inter-Molecular Force (IMF) you must know by name. When the δ + hydrogen of one water molecule are attracted to the δ - oxygen of a different water molecule.

5. Water Properties hydrogen bonding high surface tension low vapor pressure

6. Surface Properties a glass so filled with water that the surface isn’t flat but bulges above the rim Meniscus water forms nearly spherical droplets

7. The surface of water acts like a skin, as the water strider demonstrates. The molecules within the liquid form hydrogen bonds with other molecules that surround them on all sides.

8. surface tension The inward force, or pull, that tends to minimize the surface area of a liquid

9. Surfactant is any substance that interferes with the hydrogen bonding between water molecules and thereby reduces surface tension used in soaps, detergents, and any cleaning solution.

10. Vapor pressure of a liquid is the result of molecules escaping the surface of the liquid and becoming vapor below the boiling point. Water has a relatively low vapor pressure Crack out you reference table, turn to table H.

11. Higher the vapor pressure, the faster a liquid evaporates ALWAYS at least ONE question on this graph, ALWAYS!

12. Part II: Definitions used with aqueous solutions

13. Solvents and Solutes Solute - dissolved particles in a solution, usually a solid into a liquid like water. In the case of two liquids, the one in the lesser amount. Solvent - the dissolving medium in a solution (usually water. In the case of two liquids, the one in the greater amount. Aqueous Solution - water that contains dissolved substances

14. A true solution can be seen through easily. It may be colored but there should be no particles to reflect light in a solution. IF you can see bits floating, it isn’t a solution, it is a suspension. (a) is a solution, (b) is not

15. Not a terribly important definition, but it could crop up on the Regents, so jot it down… When a salt is added to water it dissociates and forms aqueous ions in solution. Solvation - a process that occurs when an ionic solute dissolves; in solution, solvent molecules surround the positive and negative ions.

16. Electrolytes and Non-electrolytes Electrolyte - conducts an electric current when it is in an aqueous solution or in the molten state. all ionic compounds are electrolytes, but most covalent compounds are not Non-electrolyte - does not conduct an electric current in aqueous solution or in the molten state (i.e. pure water)

17. Suspensions a mixture where some of the particles settle out slowly upon standing. differs from a solution because the particles of a suspension are much larger and do not stay suspended indefinitely. are heterogeneous because at least two substances can be identified. Ex. Mud water

18. Colloids are rarely asked on exam, but are common in life. Learn this and impress your family with you superior science knowledge. a heterogeneous mixture whose particles are intermediate in size (particles that range in size from 1 nm to 1000 nm) between those of a suspension and a true solution. The particles are spread throughout the medium, which can be a solid, liquid, or gas. The first substances to be identified as colloids were glues. Other colloids include such mixtures as gelatin, paint, aerosol spray, ice cream

19. Emulsions ditto from colloids a colloidal dispersion of a liquid in a liquid. An emulsifying agent forms an emulsion and maintains its stability. e.g.oils and greases are not soluble in water. form a colloidal dispersion if soap or detergent is added to the water. Soaps and detergents are emulsifying agents. Mayonnaise is a heterogeneous mixture of oil and vinegar. quickly separate without the presence of egg, which is the emulsifying agent.

20. Soaps and Detergents this ya gotta know One end of a large soap or detergent molecule is polar and is attracted to water molecules. “hydrophilic” The other end of the soap or detergent molecule is non-polar and is soluble in oil or grease. “hydrophobic” Soaps and other emulsifying agents thus allow the formation of colloidal dispersions between liquids that do not ordinarily mix.

21. Part III: Colligative Properties of Water

22. Colligative Properties of Water Colligative properties are the physical changes that result from adding solute to a solvent. Colligative Properties depend on how many solute particles are present as well as the solvent amount, but they do NOT depend on the type of solute particles. ¡Muy importante! Tres important! Sehr wichtig! immer in der Prüfung!

23. Boiling point elevation: difference between the boiling points of the pure solvent and the solution is proportional to the concentration of the solute particles i.e. greater concentration of solute, the higher the boiling point will be elevated boiling points of solutions are all higher than that of the pure solvent.

24. Freezing Point Depression freezing points of solutions are all lower than that of the pure solvent. freezing point depression is directly proportional to the concentration of the solute. Greater concentration, lower the freezing point (the GREATER the FREEZING POINT DEPRESSION) CaCl2 will lower freezing point more than NaCl. Why?

25. Vapor Pressure Lowering The vapor pressure of a solvent in a solution is always lower than the vapor pressure of the pure solvent. vapor pressure lowering is directly proportional to the concentration of the solute.

26. Part IV: Water & Energy

27. Measurement of energy Energy in various forms may be converted to heat, heat units (calories or Joules) are used to measure the energies involved in chemical reactions.

28. Joule – DON’T copy ‘til I explain the amount of heat required to raise the temperature of one gram of water one Kelvin is 4.18 Joules. 4.18 Joules = 1 calorie 1 kilocalorie is equivalent to 1000 calories 1 food Calorie = 1 kilocalorie = 1000 calories Q = m c ΔT where Q = heat measured in calories or Joules where m = mass measured in grams where c = Specific Heat measured in Joules/g 0 C or Joules/gK (Regents have changed over to all Joules and Kelvin degrees). where ΔT = change in temperature

29. Specific Heat Is the amount of heat required to change the temperature of a 1 gram sample 1 K (°Celsius). is the amount of heat a sample can hold. The Specific Heat for water is: 4.18 Joules / gram K See reference table B

30. Evaporation vs. Vaporization Evaporation is the process that describes a change from a liquid phase to a gaseous phase below the substance’s boiling point. Vaporization is the process of liquid to gas at the boiling point.(see Vapor Pressure) These changes are Endothermic (requires energy in)

31. Condensation is the process that describes a change from a gaseous phase to a liquid phase. This change is Exothermic (releases energy out)

32. Heat during a phase change During a phase change temperature remains the same Heat of vaporization (use at liquid-gas phase change) Heat of Vaporization for water is 2260 Joules / gram See reference table B Heat of fusion (use at solid-liquid phase change) Heat of Fusion for water is 334 Joules / gram See reference table B