1. UNIT IX Solution Chemistry Lesson #1

2. I NTRODUCTION Solution Chemistry is the study of chemical reactions that occur in solutions … Reactions in solutions are chemists’ favourite type of reactions…because of how easy and convenient they are. Compare to: Reactions in the gas phase are complicated  need special containers (air-tight) Solid reactions  very slow or do not occur at all

3. H OW TO MAKE SOLUTIONS …

4. L ET ’ S REMEMBER SOME DEFINITIONS Solution : is a homogeneous mixture that exists in one phase. It has uniform composition throughout and retains some of the properties of its components. Solvent: the component/substance in the solution that exists in the ___________ amount. Solute: the component in the solution that exists in the ___________ amount. It is dissolved in the solvent.

5. S OLUBIITY Solubility:the ability of one substance to dissolve within another substance. The solubility of a substance is the max. amount of solute that will dissolve in a given amount of solvent at a particular temperature.

6. SOLUBILITY Not all things are created equal… some salts dissolve better in water than others. Soluble means more than 0.1mol of solute will dissolve in 1.0L of solution at 25 C. Low solubility means less than 0.1mol of solute will dissolve in 1.0L of solution at 25 C.

8. A) Saturated solutions: a solution in which no more solid will dissolve. This solution will always have undissolved solid in it.  Additional solute will not dissolve in a saturated solution. B) Unsaturated solution: a solution in which the solvent is still capable of dissolving more of the solute.  So the solute is completely dissolved in the solvent.

9. Saturated Vs Unsaturated

10. S OLUTIONS AND ELECTRICAL CONDUCTIVITY Can some solutions conduct electricity? How do we know?

11. Requirements to conduct electricity Experiments tell us that you need… Electrical charge to be transferred in the solution in order to produce electricity (a closed circle) IONS carry electrical charge!!! We call solutions that can carry electrical charge an e lectrolyte. Electrolytes : a substance that dissolves in solutions to produce ions.

12. Requirements to conduct electricity IONS carry electrical charge! Therefore, Ionic compounds can conduct electricity but only when they’re liquid or dissolved in H 2 0 (aqeous) NaCl (aq) NaCl (s) *In the solid crystal lattice structure, the ions are too locked in, so they can’t move freely.

13. Requirements to conduct electricity Can covalent compounds conduct electricity? CS 2(aq) PCl 3(aq)

14. THE CONDUCTIVITY OF AQUEOUS SOLUTIONS A compound made up of a METAL and NONMETAL is IONIC  forms a conducting solution in water because it breaks into ________ A substance made up of a NON-METAL and a NON- METAL is COVALENT  will NOT form a conducting solution in water because it stays as a molecular compound.

15. R ULE : The greater the concentration of ions, _____________the conductivity. A 3.0M solution of NaCl (aq) is a better conducting solution than a 1.0M solution of NaCl (aq)

16. NaCl (ionic) in water The "+" and "-" ions are now free to move around. The "+" ions would be attracted to a negative electrode and the "-" ions would be attracted to a positive electrode. In this way, the ionic solution conducts a current.

17. What else can conduct electricity?? ACIDS and BASES form conducting solutions in water. Why? Because they can break into ions! HCl (aq) H 2 SO 4(s) NaOH (aq) Ca(OH) 2 (aq)

18. How to identify Acids & Bases? 1- Think of acids as any compounds that start with H 2- Think of bases as any compounds that end with OH Exception; if a compound starts with Carbon and ends with OH. It is organic, not a base and can’t conduct electricity. Ex: CH 3 OH, C 2 H 5 OH (can’t conduct electricity)

19. What else can conduct electricity?? Metals in all phases can conduct electricity Examples: Na(l), Cu(aq), Na(s)

20. Things that can NOT conduct electricity 1) Non-metals 2) Covalent compounds (non-metal + non- metal) 3) Organic compounds -usually start with Carbon(C) and contain hydrogens(H)- Ex: CH 3 OH, C 4 H 10 * CH 3 COOH (exception)- can conduct. CH 3 COOH  Acetic acid (vinegar): weak electrolyte 4) Solids !!!!!!!!!!! (except if they’re metals)

21. TO CONDUCT or NOT TO CONDUCT??? This is the question..... CONDUCTDOESN’T CONDUCT Metals –any phase- non metals Ionic (aq) solid except metals. Acids (aq) covalent compounds Bases (aq) *organic compounds * CH 3 COOH (organic exception) http://www.youtube.com/watch?v=UHYWIM8AbPE

22. T HINK ABOUT THIS For solutions to form, and to mix chemical compounds together when adding a solute to a solvent to make a solution. This process of dissolving and forming solutions requires understanding the intermolecular forces that hold molecules together… So, we need to remember the intermolecular forces that hold molecules together….

23. Back to forces between molecules… Van der Waals Forces (intermolecular) THREE main types: A. London Forces B. Dipole-dipole Forces C. HYDROGEN BONDING All are caused by dipoles!

24. R EMINDER OF DIPOLES … Dipole: A temporary separation of charges, which exists when one end of a molecule has a slight excess of negative charge and the other end has a slight excess of a positive charge

25. London Forces: LONDON FORCES exist between everything!!! London forces’ strength increases as the # of electrons is increased. Ex: Does He or Xe have a stronger london forces bond? A. LONDON FORCES

26. B. DIPOLE-DIPOLE FORCES Dipole-Dipole Forces: a permanent dipole resulting in polar molecules, because of differences in electronegativity between the atoms.

27. E XPLANATION OF DIPOLE - DIPOLE F ORCES  within any substance containing polar molecules, each molecule will have its own dipole. Therefore, when these molecules are in the liquid or solid phases, they will naturally orient themselves so that the +ve pole of the molecule is next to the –ve pole of the adjacent molecule.

28. Polar Vs Non-Polar Polar molecules HCl Non-polar molecules Cl 2

29. C. HYDROGEN BONDING

30. C. Hydrogen Bond: a strong dipole – dipole attraction between molecules containing a H – N, H – O, or H – F bond (because N,F and O are highly electronegative) *It is the strongest intermolecular bond (van der waals), but is still weaker than covalent and ionic bonds. NH 3 and NH 3 H 2 O and H 2 O

31. P REDICTING H YDROGEN B ONDS … Look for the following in compounds:  HF  NH, NH 2, NH 3 Ex: CH 3 NH 2  OH, H 2 O, H 2 O 2

32. C OMPARING STRENGTHS OF BONDS Ionic/Covalent bond  Hydrogen bond > Dipole-dipole ~ London forces Ionic bonds & Covalent bonds will always be stronger than any intermolecular bond. London forces is always the weakest.

33. E XPLAIN THIS : Boiling temperatures: ICl = 97 ℃ (70 electrons) Br 2 = 59 ℃ (70 electrons)

34. H OMEWORK Page 198,199 Questions 6,7 Question 8 (solve a,c,e, etc) Question 9 Page 202, 203, 204 Question 11 and Question 12 Question 14 and Question 16