Unit 5 Chemical Bonding PDF

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Unit 5 Chemical Bonding.docx
Dr Chen - Prof...

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Unit 5 Chemical Bonding Bonding - Involves transfer or sharing of outer electrons, usually to acquire a stable configuration\ - Ionic bonding (transfer of electrons: Usually between a metal and non-metal Covalent Bonding - Sharing of electrons - Often to attain a complete shell of electrons - Often between 2 nonmetals - Lewis structure shows all electrons as equivalent - Bonds depicted as lines Electronegativity (EN) – The Final Trend EN: Atom’s ability to compete for e in a bond, with larger values assigned to elements which are more able to pull bonding electrons towards themselves. Derived from the behaviour of atoms Trend: EN increases across a period (left to right) and up a group. Mirroring trend of ionization energy and electron affinity. Pauling scale: F 4.0 (highest EN) Bond Polarity - Dictated by the difference in electronegativity (EN) EN Bonding Small (1.9) Ionic

Example CL2 (EN = 0) PCL5 (EN = 0.9) NaCl (EN =2.1)

Polar covalent bonds: - Unequal sharing of e- Indicated by polar arrow and partial charges Electrostatic Potential Maps – useful to visualize bond polarity - These maps show surface around the molecule which encompasses most of the electron density in the molecule. (Typically, 95% of the electron density and each point on this surface is assigned a colour based on the electrostatic potential at this point) - Electrostatic potential: the work done in moving a positive point charge from an infinite distance to each point on the electron density surface. - For regions of delta + on the electron density surface, the point charge will be repelled. - Regions of delta - on the electron density surface, the point charge will be attracted. - Regions are depicted using a rainbow colour scale, with the most delta plus regions shown in blue and the most delta minus regions shown in red.

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For Cl2, the electron density surface around the molecule, looks like a nice symmetrical peanut. Image on the right, uses a transparent surface in order to see the position of the chlorine nuclei relative to electron density surface. In Cl2, the surface is coloured entirely yellow, since the bond is purely covalent. For HCL, the electron density surface around the molecule is a different shape, because hydrogen is much smaller than chlorine. The surface is mostly yellow, with some green around hydrogen, consistent with a polar covalent bond. NaCl has blue colour around sodium and red around chlorine. This bond is purely ionic. Sodium cation and chloride anion held together by electrostatic forces (attraction between opposite charge). A line bond should not be shown between Na and Cl.

Lewis Structures - Show bonding (b) and non-bonding (nb) e-, and formal charges - Bonding e- can be involved ins single, double, triple bonds - Complete shells can be achieved by combination of bonding and nonbonding e- A complete shell is an octet for most elements except hydrogen which is satisfied with 1 electron pair - Exceptions: where the octet rule is not followed are: - Beryllium can be satisfied with 2 electron pairs - Boron/Aluminum can be satisfied with 3 electron pairs - Elements in period 3 and beyond can have an expanded octet if involved as the central atom. Drawing Lewis Structures 1) Count total # of valence e- including charge of structure - Add e- for negative charge, subtract e for positive charge. 2) Draw skeletal structure (central and terminal atoms) - Least electronegative atom is usually the central atom - Hydrogen and Fluorine are always terminal 3) Complete the octet of terminal atoms (or 2 e- for hydrogen) 4) Subtract all e- used in previous steps and place any remaining e- on central atom. - This sometimes led to an expanded octet. 5) Calculate formal charges (FC) on each atom - FC = (Valence e- - number of bonds – number of lone pair e-) 6) Minimize formal charges by creating multiple bonds using nonbonding electrons. - Typically happens when neighbouring atoms have opposite charges. - Only minimizes when it will not give an element from the 2nd period > 8e-. 7) Ensure all atoms have an allowed electron count (e.g. C, N, O,F must obey octet rule, and only Be and group 13 elements are allowed to have less than an octet). Resonance Structures for PO4^3- For PO4^3-, there are 4 equivalent charge minimized structures (resonances structures) - The molecule exists as a hybrid of all formal resonance structures, known as the resonance hybrid.

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Resonance hybrid bond lengths, orders and charges are the average of all equivalent resonance states. - Resonance structures contribution to a resonance hybrid must have the same atomic arrangement; they differ in how the electrons are arranged. - Most polyatomic anions have resonance structures. Average formal charge for an atom: total charges on atom / total # of that atom Average formal charge on O: 0 +(-1) + (-1) + (-1) / 4 = ¾ = -0.75 Bond order: single (1), double (2) Average bond order: total number of bonds involved in the resonance structure/ total # of places the is formed. Average P-O bond order = 2+1+1+1/4 = 5/4 = 1.25 Bond Order, Length and Energy Covalent bond length - Distance between 2 nuclei involved in a covalent bond Bond dissociation energy (homolysis) - Approximate energy required to break 1 mole of bonds in gas phase. As bond order increases, bond length decreases. As bond length decreases, bond energy increases. Bond Order Length (pm) Energy (kJ mol) C-C 1 154 347 C=C 2 134 611 C---C 3 120 837 N---N 3 109.8 946 Molecular Shape - VSEPR (valence shell electron pair repulsion) theory - AKA Gillespie Nyholm theory - (Ron Gillespie, McMaster Chemistry) - Electron groups repel one another - As electron group is defined as a lone pair or a bond (whether or not the bond is a single, double or triple bond) - Repulsions decreases in the order: lone pair – lone pair > bonded pair – lone pair > bonded pair – bonded pair - Note: double bonds occupy slightly more space (cause more repulsion) than a lone pair. VSEPR Classes AXnEm A = central atom X = atoms bonded to central atom E = lone electron pairs

2 electron groups - Electron Group Geometry – Linear Chemical Formula Molecular Geometry Angles Symmetry Example

AX2 Linear 180 degrees “symmetrical” BECl2, C02

3 electron groups Electron Group Geometry – Trigonal Planar Chemical Formula AX3 Molecular Geometry Trigonal Planar Angles 120 degrees Symmetry “symmetrical” Example BF3 -

AX2 Bent...