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Lecture 3

CHEM 1A03 Lecture 3: Bonding and Lewis Structures (Chapter 10).docx

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McMaster University
Jeff Landry

Chapter 10: Chemical Bonding Bonding - Involves transfer or sharing of outer electrons, usually to acquire a stable configuration (Lewis) - Ionic boinding: transfer of electrons, usually between a metal and non-metal 1 + 2 5 - - Na [Ne]3s becomes Na [Ne] and Cl [Ne]3s 3p becomes Cl [Ar] - Covalent Bonding - sharing of electrons, often to attain an octet of electrons o Often between 2 non-metals o lewis scructure shows all e’s at equivalent - Coordinate covalent bond o One atom provides both e- for a bond o eg. NH3 + H+  NH4+ - Electronegativity - Atom’s ability to compete for e- in a bond - Trend: EN increases across a period and up a group - Pauling scale: F 4.0 (highest EN) Bond Polarity - Unequal sharing of e- - Indicated by polar arrow and partial charges - Dictated by deltaEN between atoms - Percent ionic character of a bond involving a certain element decreases across a period EN Bonding Example Large Ioni NaCl Intermediate Polar Covalent PCl5 Zero (small) Pure covalent Cl2 - 0-0.4 covalent (weakly polar covalent ) 0.4-1.9 polar covalent, >1.9 ionic Lewis Structures - Show bonding (b) and non-bonding (nb) e=, and formal charges - Octet* can be achieved by combination of bonding and nonbonding e= *(not all atoms will have an octet) - Bonding e- can be single, double, triple, bonds BF3 F - B: 1x3 = 3, F: 3x7 = 21, Total e- = 24 - Initial e- : 24, Bonds: -6  18 bonds B - Outer e-: -18  18-18 = 0 left F F - Formal charge: F: 7-1-6 = 0, B: 3-3-0 = 0 - Calculating formal charge: FC = Group A # - #non-bonding e- - #bonds - Minimize formal charges NO3 - N: 1 x 5 = 5 O: 3 x 6 = 18 charge = 1 ___________ Totale e- = 24 - Initial e-: 24 Bonds: -6 ___________ 18 electrons left BrOF2+ see note Rules - Octet is not exceeded in period 2 o CNOF (can not over fill) o Cannot exceed the “octet rule“ even if there are adjacent formal charges that could be minimized - For elements in period 3, 4, etc., minimizing formal charges is the priority, even if it means breaking the “octet rule” - Usually don’t have adjacent atoms with same formal charge - Negative formal charges usually appear on the most electronegative atoms, positive charges on the most electropositive atoms - ClO3 o Cl is in the period 3 and can accommodate more than an octet of electrons o A charge-minimized Lewis structure shows that formal charges have been made as close to zero as possible; for elements in period 3, 4, 5… this can often be achieved by creating multiple bonds in order to minimize formal charges towards zero o Structure (b) on previous shows this - Period 3 – s, p, & d orbitals – can fill those with electrons therefore for elements n=3 or below you can break octet rule Hemoglobin - Carboxyhemoglobin: the silent killer - Lone pair of electrons on o2 molecule that help bind to hemoglobin - Lone pair of e- on CO2 binds too but is not nearly as strong as O2 – competitive process but O2 is stronger so it replaces CO2 - Carbon monoxide – CO binds competitively to heme unit and displaces O2 – CO poisoning! - CO does irreversible binding to heme – suffocated on molecular level, body becomes incapable of transporting O2 around - <50ppm upper safetly limit, 50-200ppm: headache/nausea, >200ppm: dizziness/convulsions - Treatment: saturation with O2 can reverse the process (the binding is an equilibrium process – ch 15) - CO : C has an unusual lone pair of electrons and neg formal charge b/c has to fulfil octet rule it’s unusual, reactive – kind of takes / pulls away e’s from O; C can kind of steal e’s away but is very capable of donating those e’s away – O2 is more electroneg, will hold on to its electrons much more than CO will; CO will give up the lone pair on the C to bind to the Fe - Higher CO emissions in Africa – forest fires, china – population, heating, coal Resonance Structures - Average formal charge - Average bond order 3- - Recall: teeth! – fluorapatite, Ca5(PO 4 3, look at PO 4 ions - PO 43equivalent resonance structures – why the bond lengths are all equal since it switches between the diff resonance structures extremely fast.. Resonance structures for PO 4 3- - Average formal charge for an atom = total charges on atom / total # of that atom - Average formal charge on O = (0 + -1 + -1 + -1 )/ 4= -3/4 - Average bond order = total # of 1 type of bond / # of places where the bond is found - Average P-O bond order = (1 + 1 + 1 + 2 ) / 4 = 5/4 or 1.25 o Take all the bonds, add them all up – 3 singles and one double – 5 total; then look at # of P-O connections, so 5 / 4 Molecular Shape VSEPR - Electron pairs repel one another - Repulsion decreases o Lone pair/lone pair > bonded pair/lone pair > bonded pair – bonded pair o Note: double bonds occupy slightly more space than a lone pair Classics: - AXn m
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