Laboratory 06 MOLECULAR GEOMETRY AND POLARITY 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. 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). 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]4f145d106s26p2 This is four (4) valence electrons. The 5d is part of the inner shell (n=5) which is full. 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 H2O, The total number of valence electrons = 2 x 1 (each H is 1s1) + 6 (O is 2s22p4) =8 For CO2 The total number of valence electrons = 4 (C is 2s22p2) + 2 * 6 (O is 2s22p4) = 16 Central Atom In a molecule, there are only 2 types of atoms: 1. 2. “central” – bonded to more than one other atom. “terminal” – bonded to only one other atom. Almost always the least electronegative atom is the central atom. For example, in ClO2, the Cl is the central atom; in SF5 the S is the central atom. Hydrogen never is the central atom. It forms only one bond, so it must generally be in the outer layer of atoms. You can have more than one central atom in a molecule. 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. 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. 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. Formal Charge of an Atom Formal charge = number of valence electrons – number of bonds – number of non-bonding electrons. Dot structure for H2O 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. Dot structure for H2O H–O-H Dot structure for H2O .. H–O–H ¨ FC (H) = 1-1-0 = 0 FC (O) = 6 – 2 – 4 = 0 That is a total of 8 valence electrons used: each bond is 2, and there are 2 non-bonding pairs. Expanded Octets Example PCl5: .. .. :Cl: :Cl: .. .. :Cl – P - Cl : ¨ | ¨ : Cl: ¨ Total valence e- = 40 FC(P) = 5 – 5 – 0 =0 FC (Cl) = 7 – 1 – 6 = 0 Background - Covalent Bonds • The simplest covalent bond is that in H2 – 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 Background – Polarity 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 D i fference in El ectron eg ati vity Betw een Bo nded Ato ms Less than 0.5 0.5 to 1.9 Greater than 1.9 Typ e of Bond N on pol ar cov alent Pol ar co valent Io ns f orm 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 d+ and d-, 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 d+ dH Cl H Cl Polar bonds and polar molecules direction of dip ole moment N O H H Water  = 1.85D H H Ammonia  = 1.47D H direction of dip ole moment Laboratory 07 QUALITATIVE ANALYSIS : TESTING THE SOLUBILITY RULES 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 AgNO3 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). Nomenclature of binary ionic compounds Symbol Anion Symbol Anion Name Br Br- Bromide Cl Cl- Chloride F F- Fluoride H H- Hydride I I- Iodide N N-3 Nitride O O-2 Oxide P P-3 Phosphide S S-2 Sulfide NaCl H2S NaF BaCl2 Mg3N2 K2O Sodium chloride Hydrogen sulfide Sodium fluoride Barium chloride Magnesium nitride Potassium oxide Polyatomic anions NO3- = nitrate NO2- = nitrite SO4 2 - = sulfate SO32- = sulfite PO43- = phosphate CO32- = carbonate HCO3- = hydrogen carbonate or bicarbonate OH- = hydroxide CN- = cyanide C2H3O2- = acetate C2O42- = oxalate NaHCO3 = sodium hydrogen carbonate or sodium bicarbonate K2SO3 = potassium sulfite MgSO4 = magnesium sulfate KCN = potassium cyanide H2PO4 = hydrogen phosphate Ca(OH)2 = calcium hydroxide NH4NO3 = ammonium nitrate Zn(NO3)2 = zinc nitrate Li3PO4 = lithium phosphate HNO3 = hydrogen nitrate Solubility Rules for Ionic Compounds in Water Soluble Ionic Compounds Insoluble Ionic Compounds 1. All common compounds of Group 1A 1. All common metal hydroxides are ions (Li+, Na+, K+, etc.) and ammonium insoluble, except those of Group 1A and ion (NH4+) are soluble. the larger members of Group 2A (beginning with Ca2+). 2. All common nitrates (NO3-), acetates 2. All common carbonates (CO32-) and (CH3COO-), and most perchlorates phosphates (PO43-) are insoluble, except (ClO4-) are soluble. those of Group 1A and NH4+. 3. All common chlorides (Cl-), bromides 3. All common sulfides are insoluble (Br-), and iodides (I-) are soluble, except except those of Group 1A, Group 2A, those of Ag+, Pb2+, Cu+, and Hg2+. All and NH4+. common fluorides (F-)are soluble, except those of Pb2+ and Group 2A. 4. All common sulfates (SO42-) are soluble, except those of Ca2+, Ba2+, Ag+, and Pb2+. 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(NO3)2(aq) + 2KI(aq) → PbI2(s) + KNO3(aq) Barium chloride(aq) + Sodium sulfate(aq) → Barium sulfate(s) + Sodium chloride(aq) BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + NaCl(aq) Copper sulfate(aq) + Sodium hydroxide(aq) → Copper hydroxide(s) + Sodium sulfate(aq) CuSO4(aq) + 2NaOH(aq) → Cu(OH)2(s) + Na2SO4(aq) Mercury(II) nitrate(aq) + Potassium iodide(aq) → Mercury iodide(s) + Potassium nitrate Hg(NO3)2(aq) + 2KI(aq) → HgI2(s) + KNO3(aq) Silver nitrate(aq) + sodium chloride(aq) →Silver chloride(s) + sodium nirate(aq) AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)