1. Solutions

2. SOLUTE + SOLVENT SOLUTION State of Matter homogenous mixture of gases (Air…) Gas:

3.  solvent: liquidsolute: liquid (ethanol in water)  solvent: liquidsolute: gas solid (soda water: CO 2 /H 2 O; brine: NaCl/H 2 O) Liquid:

4.  solvent: solidsolute: liquid (Dental-filling alloy)  solvent: solidsolute: solid (gold-silver alloy) Solid:

5. Dissolve: solute + solvent  solution. Crystallization: solution  solute + solvent. Saturation: crystallization and dissolution are in equilibrium. Solubility: amount of solute required to form a saturated solution. Supersaturated: a solution formed when more solute is dissolved than in a saturated solution. Miscible: two liquids that mix. Immiscible: two liquids that do not mix. Terms to Know…

6. Like Dissolves Like “Rule”: polar solvents dissolve polar solutes. Non-polar solvents dissolve non-polar solutes. Why? If ΔH soln is too endothermic a solution will not form. NaCl in octane (C 8 H 18 ): the ion- London forces are weak because octane is non-polar. Therefore, the ion-London forces do not compensate for the separation of ions. NaCl dissolves nicely in water. +– +– +– +– +– +– NaCl octane water

7. Terminology: Solubility: is the maximum amount of the solute that will dissolve in a definite amount of solvent (at a given t°) g/100 mL Concentration: ratio of the solute and the solvent Dilute solutions Concentrated solutions Saturated solutions: solute pure solute dissolved

8. Supersaturated solutions: g/100 mLKNO 3 PbCl 2 NaCl Solubility curves t°

9. Temperature Effects: Solids Experience tells us that sugar dissolves better in warm water than cold. As temperature increases, solubility of solids generally increases. Sometimes, solubility decreases as temperature increases (e.g. Ce 2 (SO 4 ) 3 ).

10. Properties of Water  Most abundant liquid  Vital to life  Universal solvent 1. High melting and boiling points M.p.: 0.0°C B.p.: 100.0°C(0.1 MPa)

11. Hydrogen Bonding  Special case of dipole-dipole forces.  By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high.  Intermolecular forces are abnormally strong. CH 4 SnH 4 GaH 4 SiH 4 H2OH2O H2SH2S H2SeH2Se H2TeH2Te

12. Water Molecule

14. Solutions

15. Phase Diagram of H 2 O  The melting point curve slopes to the left because ice is less dense than water.  Triple point occurs at 0.0098°C and 4.58 mmHg.  Normal melting (freezing) point is 0°C.  Normal boiling point is 100°C.  Critical point is 374°C and 218 atm.

16. 2. Density 0.0°C (ice)0.91680 g/cm 3 0.0°C (liquid)0.99984 g/cm 3 3.98°C0.99997 g/cm 3 25.0°C0.99704 g/cm 3 3. Surface tension: high Bottom of meniscus

17. 4. High heat of vaporization 40.70 kJ/mol 5. High heat of fusion 6.02 kJ/mol 6. High specific heat 75.20 kJ/mol

18. 1. Chemically pure water 2. Water of crystallization or hydration CuSO 4 5H 2 O (s) CuSO 4(s) + 5H 2 O (g) (CaSO 4 ) 2 H 2 O(plaster of paris) CaSO 4 2H 2 O (gypsum) Na 2 CO 3 10H 2 O 

19. 3. Groundwater CationsAnions Ca 2+ HCO 3 -, CO 3 2- Na + OH - Mg 2+ SO 4 2- K+K+ Cl - Fe 2+, Fe 3+ NO 3 - NH 4 + F -, PO 4 3-

20. Hard water Soft water Water Softening  Temporary hard water HCO 3 - (bicarbonate)  Permanent hard water  Boiling  Softening agentsprecipitation complex formation  Distillation  Ion-exchange Mineral water Thermal water