Lewis Structures and Molecular Geometry PDF

MED-102 General Chemistry Lewis Structures and Molecular Geometry  LOBs covered Draw Lewis structures of molecules and molecular ions Describe how resonance structures form and identify the structure of the resonance hybrid Calculate the formal charge Discuss how th...

MED-102 General Chemistry Lewis Structures and Molecular Geometry  LOBs covered Draw Lewis structures of molecules and molecular ions Describe how resonance structures form and identify the structure of the resonance hybrid Calculate the formal charge Discuss how the formal charge can be used to determine the relative stabilities of resonance forms Determine the geometry (shape) of molecules and molecular ions  Lewis structure – the first step Distribution of the valence electrons on the atoms being bonded together Examples:  Lewis structure - steps Count the total number of valence electrons Join atoms symmetrically using single bonds Subtract electrons used up Complete the octet (8 electrons) on the exterior atoms Subtract electrons used up Place any left-over electrons on the central atom If not enough electrons to give an octet to the central atom, try using multiple bonds WATCH: https://www.youtube.com/watch?v=cIuXl7o6mAw WATCH: https://www.youtube.com/watch?v=qwqXAlvNxsU  Lewis structures - examples CH4 NH3 H2O All these molecules have a total of 8 valence electrons  Examples Showing the Procedure Five Examples Demonstrating the Procedure PART A  Gilbert N. Lewis American physical chemist Discovered the covalent bond Concept of electron (lone) pairs Lewis dot structures Contributed to valence bond theory Lewis acids and bases Coined the term “photon” in 1926 Contributed to: Chemical thermodynamics Photochemistry Isotope separation 23/10/1875 – Died in his lab of cyanide poisoning 23/03/1946  NH 4 + Lewis structure N = 5e H = 1e x 4 = 4e -1 (positive charge) Total = 8e We must place the structure in square brackets since it is an ion  PCl 4 - Lewis structure P = 5e Cl = 7e x 4 = 28e +1e (negative charge) Total = 34e  Examples Showing the Procedure Two Examples Demonstrating the Procedure PART B  CO2 Lewis structure C = 4e O = 6e x 2 = 12e The question is: why 2 double bonds? Why not a single and a triple bond?  CO2 Lewis structure There are in fact three different resonance forms However, we will show that the middle resonance form is the most important  Resonance Consider the CO32- anion Lewis structure 24e Experiments show that all three CO bonds have equal length The real structure is NOT a time average between the three forms Mixture of blue and red paint purple paint WATCH: https://www.youtube.com/watch?v=MWDL5WCZBzE Actual structure of CO 3 2- A blend of all three resonance forms  Contributions from each resonance form We need to look at the stability of each resonance form For this purpose we employ the concept of formal charge  Formal Charge The charge an atom would have in a molecule if bonding electrons were equally shared Formal Charge – An Easy Way! Let’s look at an easy way to obtain the formal charge Formal Charge = Group Number – (Dashes + Dots) F.C. (N) = 5 – (4 + 0) = +1 F.C. (O-1) = 6 – (1 + 6) = -1 F.C. (O-2) = 6 – (1 + 6) = -1 F.C. (O-3) = 6 – (2 + 4) = 0 NOTE: The sum of all formal charges should equal the total charge of the compound  Formal Charge Find the formal charge on all atoms using the easy way  Why formal charge? Use it to decide between different resonance forms Resonance form with most zeroes is preferred Negative charge should reside on most electronegative atom  Stability of Resonance Forms The last form has the most non-zero formal charges. It will be the least stable resonance form and will contribute the smallest percentage to the resonance hybrid  Stability of Resonance Forms Since N is more electronegative than S, we prefer the resonance form that places the negative charge on the N. Therefore, the second resonance form is more important.  Composition of Resonance Hybrid 5-Minute Break  VSEPR Theory Valence Shell Electron Pair Repulsion (VSEPR) Electron pairs should stay as far away from each other as possible (because electrons pairs are negative) Useful for finding the geometry (shape) of molecules We need a correct Lewis structure first Find the number of regions of electron density (bonds or lone pairs) on the central atom – One region = 1 single bond = 1 lone pair = 1 double bond = 1 triple bond This gives us the electron pair or basic geometry Then, visually ignore lone pairs to get molecular geometry WATCH: https://www.youtube.com/watch?v=Q9-JjyAEqnU  Electron Pair Geometry 2 electron pairs – linear (180o) 3 electron pairs – trigonal planar (120o) 4 electron pairs – tetrahedral (109.5o) 5 electron pairs – trigonal bipyramidal (90o and 120o) 6 electron pairs – octahedral (90o) Molecular Geometry Molecular Geometry Molecular Geometry  Finding Molecular Geometry Seven Examples Demonstrating the Procedure  Nonpolar Compounds Linear AB2 Trigonal planar AB3 Square planar and tetrahedral AB4 Trigonal bipyramidal AB5 Octahedral AB6 Anything not included here must be polar