honors - chemical bonding

honors - chemical bonding - CHEMICAL BONDING I. Ionic Bonds...

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Unformatted text preview: CHEMICAL BONDING I. Ionic Bonds A. Ionic Bonds form so atoms may have the same electronic configuration as inert gases. 1. Ions form when atoms gain or lose electrons. 2. Positive ions are very strongly attractive to negative ions. 3. Ionic bonds are classified as strong bonds. B. Ionic bonds are formed when a metallic atom transfers an electron to a nonmetallic atom. 1. Ionic compounds are classified as salts. C. Ionic substances form a lattice structure, which are repeating units with strong intermolecular attractive forces. D. Physical Properties of ionic compounds. 1. Ionic compounds are solids at room temperature. 2. Ionic compounds are conductors in the molten (liquid) state, due to free flowing ions. 3. When they are in solid state, they are non-conductors or non-electrolytes. 4. Water soluble ionic compounds are electrolytes in the aqueous state. E. Atomic Size 1. Positively charged ions will have a smaller radius. 2. Negatively charged ions will have a larger radius. F. Electro-negativities 1. Ionic bonds generally have an electronegativity difference that is greater than 1.7. 2. Ionic compounds are electrically neutral. G. Must review all nomenclature rules. H. Writing net ionic equations Precipitation Rxn (reaction) 2NaCl(aq) + Pb(NO3)2(aq) ⎯⎯ 2NaNO3(aq) + PbCl2 (s) → + +2 2Na (aq) + 2Cl (aq) + Pb (aq) + 2NO3-(aq) ⎯⎯ 2Na+(aq) + 2NO3-(aq) + PbCl2(s) → 2Na+(aq) + 2Cl-(aq) + Pb+2(aq) + 2NO3-(aq) ⎯⎯ 2Na+(aq) + 2NO3-(aq) + PbCl2(s) → +2 2Cl (aq) + Pb (aq) ⎯⎯ PbCl2(s) → II. Covalent Bonding A. Covalent Bonding occurs when electron pairs are shared between two atoms. 1. Electrons are shared when 2 orbitals slightly overlap. B. The elements that form covalent bonds are the group 6 nonmetals, group 7 nonmetals, N, P, C, Si, and H 1. Covalent bonds exist between non-metals. 2. Polyatomic ions are covalently bonded, which is why they must be treated as a unit. *** Polyatomic ions DO NOT dissociate in water *** Examples: Ba(NO3)2 Ba+ 2NO3- K2SO4 2K+ SO4-2 C. Types of covalent bonds 1. Single bond (sigma bond) C-C 2. Double bond (pi bond π ) C=C 3. Triple bond (1sigma δ and 2 pi π ) C≡C 4. Bond distance decreases as you increase the number of bonds. D. Lewis Dot Structures 1. Used to indicate the valence electrons only 2. The assignment of valence electron dots in the Lewis dot structure generally follows the octet rule (with the exception of He) 3. Examples: E. Properties of molecular (covalently bonded) substances 1. They have a low m.p. (lack the strong intermolecular attractive forces) 2. Non-electrolytes (except for acids) 3. They do not dissociate (except for acids) 4. Most molecular substances are liquids at room temperature. F. Molecular Geometry 1. Due to repulsive forces (within molecules), bond angles are altered for maximum electron distancing. 2. Lone electron pairs on the central atoms are more repulsive than bonded atoms. 3. Molecular geometry is determined by the electron repulsive forces in the molecular form. a. VSEPR Theory (Valence Shell Electrons Pair Repulsion Theory 4. Examples of common molecular geometry: (See Next Page) G. Polarity 1. Molecular Polarity a. Polar molecules exist when there is an uneven distribution of electrons about the central atom. 1. Polar molecules lack symmetry about the central atom. 2. Any molecule with lone electron pairs on the central atom is polar 3. Polar molecules will “tip the plate”. 2. Bond Polarity- Determined through electronegativity differences. a. If electronegativity difference is ≤ 0.3, there is equal sharing of electrons in the bond, therefore, the bond is non-polar covalent. b. If the electronegativity is between 0.3 and 1.7, there is an uneven sharing of electrons in the bond. Therefore the bond is polar covalent. c. If electronegativity is ≥ 1.7, there has been a transfer of electrons and it is ionic. Examples Formula NO MgO Br2 LiBr CuF CCl4 HAt EN Difference 0.5 2.3 0 1.8 2.1 1 0.1 Bond Type Greatest electronegativity Polar Covalent Oxygen Ionic Oxygen Non-polar covalent none Ionic Bromine Ionic Fluorine Polar Covalent Oxygen Non-polar covalent Hydrogen (Continued) Formula NI3 BCl3 CL2O OF2 CH3Cl Lewis Dot Structure (Not Given): Make Lewis Dot Structures to help determine the molecular Geometry Geometry Trigonal Pyramidal Trigonal Planar Angular / Bent Angular / Bent Tetrahedral Bond Polarity Molecular Polarity 0.5 Polar Covalent Polar 1 Polar Covalent Non-Polar 0.5 Polar Covalent Polar 0.5 Polar Covalent Polar Polar Covalent Polar H. Van der Waals forces (intermolecular forces) 1. Hydrogen bonding (STUPID NAME!!!) a. a. Hydrogen bond exists between any hydrogen, which is bonded to an N, O, or F atom and an N, O, or F atom of an adjacent molecule. 2. Dipole-Dipole Forces a. They exist between polar covalent molecules. 3. London or dispersion forces (weakest) a. Exist between non-polar covalent molecules. 4. Van der Waal forces decrease in strength from Hydrogen bonding to Dipole Dipole to London dispersion. 5. They increase melting points, boiling points, and activation energy (EA) for all molecular substances. 6. The Van der Waal forces are the intermolecular forces between molecules. I. Pot Pourri 1. With respect to bonding between atoms a. As bonds are formed, kinetic energy decreases and potential energy increases. b. Orbitals overlap in bonding, they do not merge completely c. Attractive forces decrease as atoms approach each other, due to the increases of impulsive forces between the nuclei. 2. Solubility a. Like dissolves like. ...
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This document was uploaded on 11/03/2011 for the course CH 101 at Montgomery College.

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