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Covalent Bonding

Covalent Bonding. General Covalent- ness. Covalent bond - bond that results from the sharing of valence electrons. Diatomic molecule: elements that are bonded to itself in nature Br I N Cl H O F. General Covalent stuff con’t. Types of bonds: Single = sigma bond

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Covalent Bonding

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  1. Covalent Bonding

  2. General Covalent-ness • Covalent bond- bond that results from the sharing of valence electrons Diatomic molecule: elements that are bonded to itself in nature Br I N Cl H O F

  3. General Covalent stuff con’t • Types of bonds: • Single = sigma bond • Double = sigma bond + pi bond • Triple = sigma bond + 2 pi bonds Lone pair • Bonds are electron pairs that are being shared • Lone Pairs- unshared pairs of electrons

  4. Bond Lengths • Bond Length-distance between the nuclei of the 2 atoms that are bonded when they are the most attracted to one another • Determined by: • Size of the atoms • How many electron pairs are shared (how many bonds there are) • As # of bonds , the bond length 

  5. Energy and Bonds • Bond dissociation energy- amount of energy needed to break a specific covalent bond • Breaking bonds ALWAYS requires the addition of energy • Endothermic vs. Exothermic

  6. How to Draw Lewis Structures for Covalent Bonding • Add up all valence electrons • Add up all octets that need to be filled (octet=2 for H) • Subtract step 1 from step 2 • Divide value from step 3 by 2 to get how many bonds are in the molecule Ex) CH4

  7. How to Draw Lewis Structures for Covalent Bonding Some hints for how they fit together: • Element closest to the left is usually the central atom • Hydrogen is always terminal (on the end) Ex) H2O CO2

  8. Lewis Structure Practice • CO • CCl4 • CO2 • Br2 • O2

  9. Coordinate Covalent Bonds • In some cases, one atom donates both electrons to a shared pair for a bond • Ex) SO4-2

  10. Bond Polarity • 2 types: polar and non-polar • Non-polar: • Polar: • Effect of electronegativity on polarity:

  11. Dipoles • What is a dipole? • Dipole moment: H − Cl

  12. Formal Charge • Rules: • All nonbonding electrons are assigned to the atom they’re attached to • Half of the bonding electrons are assigned to each atom in the bond • Subtract # of electrons assigned from the # of actual valence electrons in the atom Ex) N2O

  13. Formal Charge con’t • Choose the Lewis structure w/ formal charges closest to zero • Choose the Lewis structure w/ negative charges on more electronegative atom • FORMAL CHARGES DO NOT REPRESENT REAL CHARGES ON ATOMS

  14. Practice • Draw the Lewis structure and assign formal charges to each atom • CO2 • CF4 • SiO2 • SO2

  15. Resonance Structures What is resonance? A condition that occurs when more than one valid Lewis structure can be written for a molecule or ion • Do NOT move the atoms; only move the bonds • Molecules behave as though there is only one possible structure

  16. Hybridization • A theory to explain bonding that doesn’t follow basic Lewis theory • blending of AO’s  results in hybridized orbitals all the same size, shape, and energy • sp3 • sp2 • sp • dsp3 • d2sp3

  17. Exceptions to the Octet Rule • Odd number of valence electrons • Some compounds form with < 8 electrons around an atom • Be and B • Electron deficient, very reactive • Expanded octet – more than 8 electrons around an atom • Period 3 or higher elements that have an empty d orbital

  18. Molecular Shape • VSEPR Model:Valence Shell Electron Pair Repulsion model • Based on an arrangement that minimizes repulsion of electrons around the central atom • lone pairs are important in determining shape: • Slightly larger orbital than bonding electrons occupy

  19. Molecular Polarity • Bond dipole- • As difference in electronegativity between 2 atoms increases, bond polarity increases • Dipole moment depends on polarities of individual bonds and geometry of m/c

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