1. Laboratory 06 MOLECULAR GEOMETRY AND POLARITY

2. Background- Lewis structure Diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule A Lewis structure can be drawn for any covalently bonded molecule, as well as coordination compounds.

3. Construction of Lewis Structures Two Rules 1.Total # of valence electrons – the total number of valence electrons must be accounted for, no extras, none missing. 2.Octet Rule – every atom should have an octet (8) electrons associated with it. Hydrogen should only have 2 (a duet).

4. Determining the number of valence electrons Full d-orbitals do not count as valence electrons. They belong to the inner shell. For example: Pb [Xe]4f 14 5d 10 6s 2 6p 2 This is four (4) valence electrons. The 5d is part of the inner shell (n=5) which is full.

5. The total number of available valence electrons is just the sum of the number of valence electrons that each atom possesses (ignoring d-orbital electrons) For H 2 O, The total number of valence electrons = 2 x 1 (each H is 1s 1 ) + 6 (O is 2s 2 2p 4 ) = 8 For CO 2 The total number of valence electrons = 4 (C is 2s 2 2p 2 ) + 2 * 6 (O is 2s 2 2p 4 ) = 16

6. Central Atom In a molecule, there are only 2 types of atoms: 1.“central” – bonded to more than one other atom. 2.“terminal” – bonded to only one other atom. Almost always the least electronegative atom is the central atom. For example, in ClO 2, the Cl is the central atom; in SF 5 the S is the central atom. You can have more than one central atom in a molecule. Hydrogen never is the central atom. It forms only one bond, so it must generally be in the outer layer of atoms.

7. Bonds Bonds are pairs of shared electrons. Each bond has 2 electrons in it. You can have multiple bonds between the same 2 atoms. For example: C-O C=O C O Each of the lines represents 1 bond with 2 electrons in it.

8. Lewis Dot Structure Each electron is represented by a dot in the structure. :Cl: ¨ That symbol with the dots indicate a chlorine atom with 7 valence electrons.

9. Drawing Lewis Dot Structures 1.Determine the total number of valence electrons. 2.Determine which atom is the “central” atom. 3.Stick everything to the central atom using a single bond. 4.Fill the octet of every atom by adding dots. 5.Verify the total number of valence electrons in the structure. 6.Add or subtract electrons to the structure by making/breaking bonds to get the correct # of valence electrons. 7.Check the “formal charge” of each atom.

10. Formal Charge of an Atom Formal charge = number of valence electrons – number of bonds – number of non-bonding electrons.

11. Dot structure for H 2 O 1. Total number of valence electrons: 6 + (2 x 1) =8 2. Central Atom – O 3. Stick all terminal atoms to the central atom using a single bond.

12. Dot structure for H 2 O H – O - H

13. Dot structure for H 2 O.. H – O – H ¨ That is a total of 8 valence electrons used: each bond is 2, and there are 2 non-bonding pairs. FC (H) = 1-1-0 = 0 FC (O) = 6 – 2 – 4 = 0

14. Expanded Octets Example PCl 5 :.... :Cl: :Cl: Total valence e - = 40.... :Cl – P - Cl : FC(P) = 5 – 5 – 0 =0 ¨ | ¨ : Cl: FC (Cl) = 7 – 1 – 6 = 0 ¨

15. Background - Covalent Bonds The simplest covalent bond is that in H 2 –the single electrons from each atom combine to form an electron pair –the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom The number of shared pairs –one shared pair forms a single bond –two shared pairs form a double bond –three shared pairs form a triple bond

16. Polar and Nonpolar Covalent Bonds Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing We divide covalent bonds into –nonpolar covalent bonds –polar covalent bonds Background – Polarity

17. An example of a polar covalent bond is that of H-Cl –the difference in electronegativity between Cl and H is 3.0 - 2.1 = 0.9 + - –Polarity can be shown by using the symbols + and -, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end HCl  +  - HCl  

18. Polar bonds and polar molecules

19. Laboratory 07 QUALITATIVE ANALYSIS : TESTING THE SOLUBILITY RULES

20. Background : Ionic Compounds 1. Most ionic compounds are also called salts. 2.Most ionic compounds exist as solids and many dissolve to form aqueous solutions. Example : AgCl insoluble in water but AgNO 3 is soluble 3. An ionic compound is made up of a metal and a nonmetal; metals are located on the left side of the periodic table and nonmetals are on the right side. 4. The cation (positive ion) is written first followed by the anion (negative ion).

21. Nomenclature of binary ionic compounds SymbolAnion Symbol Anion Name BrBr - Bromide ClCl - Chloride FF-F- Fluoride HH-H- Hydride II-I- Iodide NN -3 Nitride OO -2 Oxide PP -3 Phosphide SS -2 Sulfide NaCl NaF H2SH2SH2SH2S Sodium chloride Sodium fluoride Hydrogen sulfide BaCl 2 K2OK2OK2OK2O Mg 3 N 2 Barium chloride Potassium oxide Magnesium nitride

22. Polyatomic anions NO 3 - = nitrateNO 2 - = nitrite SO 4 2 - = sulfateSO 3 2- = sulfite PO 4 3- = phosphateCO 3 2- = carbonate HCO 3 - = hydrogen carbonate or bicarbonate OH - = hydroxideCN - = cyanide C 2 H 3 O 2 - = acetateC 2 O 4 2- = oxalate

23. NaHCO 3 = sodium hydrogen carbonate or sodium bicarbonate K 2 SO 3 = potassium sulfite MgSO 4 = magnesium sulfate KCN = potassium cyanide H 2 PO 4 = hydrogen phosphate Ca(OH) 2 = calcium hydroxide NH 4 NO 3 = ammonium nitrate Zn(NO 3 ) 2 = zinc nitrate Li 3 PO 4 = lithium phosphate HNO 3 = hydrogen nitrate

24. Solubility Rules for Ionic Compounds in Water Soluble Ionic CompoundsInsoluble Ionic Compounds 1. All common compounds of Group 1A ions (Li +, Na +, K +, etc.) and ammonium ion (NH 4 + ) are soluble. 1. All common metal hydroxides are insoluble, except those of Group 1A and the larger members of Group 2A (beginning with Ca 2+ ). 2. All common nitrates (NO 3 - ), acetates (CH 3 COO - ), and most perchlorates (ClO 4 - ) are soluble. 2. All common carbonates (CO 3 2- ) and phosphates (PO 4 3- ) are insoluble, except those of Group 1A and NH4 +. 3. All common chlorides (Cl - ), bromides (Br - ), and iodides (I - ) are soluble, except those of Ag +, Pb 2+, Cu +, and Hg 2+. All common fluorides (F - )are soluble, except those of Pb 2+ and Group 2A. 3. All common sulfides are insoluble except those of Group 1A, Group 2A, and NH 4 +. 4. All common sulfates (SO 4 2- ) are soluble, except those of Ca 2+, Ba 2+, Ag +, and Pb 2+.

25. Precipitation reactions General form: Solution A + Solution B → Insoluble Solid C + Solution D. In a precipitation reaction two solutions are mixed together to produce an insoluble solid which is called the precipitate. This type of reaction is also called a double displacement reaction Lead nitrate(aq) + Potassium iodide(aq) → Lead iodide(s) + potassium nitrate(aq) Pb(NO 3 ) 2 (aq) + 2KI(aq) → PbI 2 (s) + KNO 3 (aq)

26. Barium chloride(aq) + Sodium sulfate(aq) → Barium sulfate(s) + Sodium chloride(aq) BaCl 2 (aq) + Na 2 SO 4 (aq) → BaSO 4 (s) + NaCl(aq) Copper sulfate(aq) + Sodium hydroxide(aq) → Copper hydroxide(s) + Sodium sulfate(aq) CuSO 4 (aq) + 2NaOH(aq) → Cu(OH) 2 (s) + Na 2 SO 4 (aq)

27. Silver nitrate(aq) + sodium chloride(aq) →Silver chloride(s) + sodium nirate(aq) AgNO 3 (aq) + NaCl(aq) → AgCl(s) + NaNO 3 (aq) Mercury(II) nitrate(aq) + Potassium iodide(aq) → Mercury iodide(s) + Potassium nitrate Hg(NO 3 ) 2 (aq) + 2KI(aq) → HgI 2 (s) + KNO 3 (aq)