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Chem 266: Structure and Bonding Review Notes

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University of Waterloo
CHEM 266
Steve Forsey

- - Effective Nuclear Charge = # protons - # shielding e (e in inner orbital)  Increases from left to right of periodic table Ionization energy – minimum energy required to remove an electron  Increases from left to right (atomic radius decreases and effective nuclear charge increases)  Decreases as you go down (force of attraction by nucleus decreases) Electronegativity – measure of the tendency for an atom in a molecule to attract a pair of electrons  Increases from left to right (greater # protons = stronger nuclear charge)  Decreases as you go down (decreasing distance from nucleus)  Derived from gas phase bond dissociation energies (amount of energy released when bond forms (-ve value) or the amount of energy needed to break a bond (+ve value)  Fluorine (F) is the most electronegative element (top-right corner of periodic table) Bone Type  look at the difference in electronegativity (between the 2 bonded atoms) ex. Cl-Cl ᵟ C-O ᵟ- Na Cl - 0.0 to 0.4 0.5 to 1.6 > 1.7 nonpolar polar ionic covalent covalent Polar covalent bonds: the sharing of electrons unequally due to differences in electronegativity  electron density is pulled toward the most electronegative atom (greatest nuclear charge) Ionic bonds: atoms gain or lose electrons (via electron transfer) to achieve the electron configuration of the nearest noble gas (a stable octet)  generally happens between atoms of very different electronegativities (metal + non-metal) Formal Charges – created from the imbalance between the # protons in the nucleus and electrons surrounding the nucleus  compare total # protons to total # electrons o total # protons = atomic number o total # electrons = valence electrons + electrons from inner orbital(s) o if # protons > # electrons, atom is positively charged o if # protons < # electrons, atom is negatively charged OR
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