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Ch. 6 Chemical Bonding

Ch. 6 Chemical Bonding. 6-1 Introduction to chemical bonding. Chemical bond -a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Types of Bonds.

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Ch. 6 Chemical Bonding

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  1. Ch. 6 Chemical Bonding • 6-1 Introduction to chemical bonding

  2. Chemical bond-a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

  3. Types of Bonds • Ionic-results from the electrical attraction between large numbers of cations and anions • Covalent-results from sharing of electron pairs between two atoms

  4. Rarely are bonds ever purely ionic or purely covalent but most lie somewhere between

  5. fig. 6-3 • 1. nonpolar covalent bond-the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge • polar- uneven distribution of charge • 2. polar-covalent bond-unequal attraction for the shared electrons

  6. Fig. 6-2 • 1. nonpolar-covalent: 0-5% ionic character; electronegativity difference of 0-0.3 • 2. polar-covalent: 5-50% ionic character; electronegativity difference of 0.3-1.7 • 3. ionic: 50-100% ionic character; electronegativity difference of >1.7

  7. 6-2 • 6-2 Covalent Bonding and Molecular Compounds

  8. Molecule-neutral group of atoms that are held together by covalent bonds • Molecular compound-chemical compound whose simplest units are molecules • Chemical formula-indicates the relative numbers of atoms of each kind in a chemical compound by using symbols and subscripts

  9. Molecular formula-the types and numbers of atoms combined in a single molecule of a molecular compound • Diatomic molecule-a molecule containing only two atoms

  10. Forming of a bond • The electrons-protons of two atoms are attracted to each other while the electrons-electrons are repelled by each other • The bond forms at a distance at which these two forces become equal, potential energy is at a minimum and a stable molecule forms

  11. Covalent Bonds • Bond length-the distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms • Bond energy-the energy required to break a chemical bond and form neutral isolated atoms

  12. Octet Rule • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level • Exceptions: H, B (6), Be(4), P, S, Xe

  13. Electron-Dot Notation • An electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol • Pg. 170

  14. Lewis Structures • Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons

  15. Lone pair (unshared pair)-pair of electrons that belongs exclusively to one atom and not involved in bonding • Structural formula-indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule ex: F-F

  16. Multiple Bonds • single bond-covalent bond produced by the sharing of one pair of electrons between two atoms • Double bond-covalent bond by sharing two pair of electrons • Triple bond-bond by sharing of three pair of electrons

  17. Resonance structures Structure that cannot be correctly represented by a single Lewis structure

  18. Ionic Bonding and Compounds • Ionic compound-composed of positive and negative ions that are combined so that the charges are equal • Ions minimize their potential energy (form bonds) by combining in an orderly arrangement known as a crystal lattice • Pg. 177

  19. Ionic Bonds/Compounds • Form between a metal (+) and nonmetal (-) • + and – ions attracted to each other • Hard, brittle • High melting pt/boiling pt • Form crystal lattices • Do not vaporize at room temp • Good electrical conductors

  20. Covalent Bonds/Compounds • Share electrons • Between 2 nonmetals • Form molecules • Lower melting pts • Vaporize at room temp

  21. Polyatomic Ions • Pg. 180 • Group of covalently bonded atoms with a charge • NO3- • NH4+

  22. 6-4 Metallic Bonding • Metals outer orbital overlap each other so electrons are free to roam from atom to atom • The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

  23. Metallic Properties • Malleable-ability to be hammered or beaten into thin sheets • Ductile-ability of a substance to be drawn into a wire • Bond strength-the amount of heat required to vaporize the metal is a measure of the bond strength

  24. Metallic Bonds • Metals bond with other metals • Attraction b/t metals and a surrounding sea of electrons • High electrical and thermal conductivity • Malleable, ductile, and high luster (shiny) • Ability to absorb wide range of light frequencies

  25. 6-5 Molecular Geometry • Properties of molecules depend on the bonding and also the geometry or shape. • The polarity depends on the geometry and determines if the molecule is polar

  26. VSEPR Theory • States that repulsion of valence electrons surrounding an atom causes them to be oriented as far apart as possible • 1. Linear-2 atoms set equal distance apart (180˚) OR 3 atoms because the electrons repel each other • Ex: AB2

  27. 2. Trigonal Planar- AB3 difference of 120º • 3. Tetrahedral- AB4 difference of 109.5º VSEPR with unshared electrons 4. Trigonal Pyramidal- AB3E (107º)-lone electrons will also repel each other but shape only involves the atoms involved in the bonding

  28. 5. Bent or Angular (105º) AB2E2 or AB2E • 6. Trigonal bipyramidal- AB5 ex: PCl5 • 7. Octahedral- AB6 ex: SF6 • Pg. 186 Table 6-5

  29. Predict the geometrical shape using the VSEPR theory and by drawing the Lewis Structure for the following: • A. HI E. SO2 • B. CBr4 F. Cl4 • C. AlBr3 G. BCl3 • D. CH2Cl2

  30. Using a protractor and pg. 186, construct the geometrical shapes using gumdrops and toothpicks

  31. Intermolecular Forces • Forces of attraction b/t molecules • Higher the boiling pt, the stronger the force • These are generally weaker than bonds that join atoms in molecules • Strongest intermolecular forces are those between polar molecules

  32. 1. Dipole-dipole=equal but opposite attractions • dipoles that are additive make it more polar ex: NH3 • Dipoles in a molecule that cancel each other make it nonpolar ex: CCl4

  33. 2. Induced dipole=temporary dipole develops • Weaker than dipole-dipole force • 3. Hydrogen bond=strong type of dipole-dipole between H-F, H-O, H-N • Makes them highly polar • Ex: water H2O

  34. 4. London dispersion forces- weak, induced instantaneous dipole created by constant motion of electrons • Only among noble gases and nonpolar molecules due to their low boiling pts. • Their strength increases with the number of electrons in the molecules or with increased mass

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