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Chemical Bonds Modern Chemistry : Chapter 6

Chemical Bonds Modern Chemistry : Chapter 6. Why ? How? What? Where?. Why?. Electromagnetism is 1 of the 4 universal forces Balance between repulsion & attraction Protons repel protons; electrons repel electrons VSEPR = valence shell electron pair repulsion

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Chemical Bonds Modern Chemistry : Chapter 6

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  1. Chemical BondsModern Chemistry:Chapter 6 Why? How? What? Where?

  2. Why? • Electromagnetism is 1 of the 4 universal forces • Balance between repulsion & attraction • Protons repel protons; electrons repel electrons • VSEPR = valence shell electron pair repulsion • Positive nuclear charges attract electrons • Electrons are from self & nearby atom(s) • Compounds are more stable than free atoms/ions • Lower potential energy at the optimal bond length • Energy is released upon bond formation, typically • Heat &/or Light • Sound • Movement

  3. Octet Rule answers “Why?” • Bonds typically form to give 8 e- in outer shell. • This provides a “noble gas configuration.” • There are, however, exceptions. • BF3 & AlCl3 only have 6 electrons for B & Al • PF5 & SF6 have more than 8 electrons for atoms • H2 & He only have 2 electrons in outer shell

  4. What? • Types of bonds • Ionic chemical bonds • Oppositely charged ions attract one another • Covalent chemical bonds • Nonpolar = Equal sharing of electrons • Polar = Unequal sharing of electrons • Metallic bonds • Mobile sea of electrons surround cations

  5. How? • Chemical bonds are determined by differences in the degrees of electronegativity (0 – 4.0) • Ionic = difference > 2.0 • Alkali & alkaline earth metals with halogens/ nonmetals • Some transition metals with nonmetals • Polar covalent = difference of 0.6 – 1.9 • Nonmetals with one another • Some transition metals with nonmetals • Nonpolar covalent = difference of 0 – 0.5 • Diatoms (I, Br, Cl, F, O, N, H) • Nonmetals with very similar nonmetals

  6. How else? • Molecular orbitals form around atoms • Size of atoms/ ions influence bond formation • Unshared valence electrons affect the shape of the entire molecule (molecular geometry) • Bond length: minimum potential energy • Single bonds are longest • Two electrons are involved • Smallest bond energy (least repulsion) • Triple bonds are shortest • Six electrons (3 pairs) are shared

  7. What about metallic bonds? • Vacant orbitals in outer energy levels overlap • Overlapping orbitals allow outer electrons to roam freely throughout the entire metal • Malleability, Ductility, & Conductivity • Many orbitals spaced by incremental energy levels allow absorption of many frequencies • Luster

  8. Where? • Covalent bonds form in the electron cloud • s bonds: symmetrical along nuclei’s axis • p bonds: side by side overlap of p orbitals in sausage-shaped regions above & below axis • Ionic bonds form in the charged space of the electron cloud • Metallic bonds form between cations of the same element as a mobile sea of electrons

  9. Electron-Dot Notation • Illustrates only the valence electrons • Nucleus & inner-shell electrons = element symbol • Valence electrons shown as dots: E, N, W, & S • Compounds shown as Lewis structures • Shared valence electrons = dot-pairs or dash(es) • Unshared valence electrons = dot(s) • Structural formulas don’t show unshared pairs • F-F, H-Cl, K-I, etc…

  10. Resonance Structures • A single representation is inadequate • Molecule may constantly alternate between bonding structures • Molecule may form an average of 2 structures • Ozone (O3) forms identical O-O bonds that are between a single & a double bond

  11. Ionic Compounds • Formula unit = simplest collection of atoms • Crystal lattice = 3-D arrangement of ions • Cubic - Monoclinic -Triclinic • Tetragonal - Hexagonal • Orthorhombic - Rhombohedral • Lattice energy (kJ/mol) = energy released upon crystal formation from gaseous ions • Polyatomic Ions: NH4+, MnO4-, SO4=, etc… • Ions held together by covalent bonds.

  12. VSEPR Theory • Molecules form to lessen e- pair repulsion • Diatoms form linear (180o) • Group III/13  form trigonal-planar (120o) • Group IV/14  form tetrahedral (109.5o) • Group V/15  form trigonal-pyramidal (107o) • Group VI/16  form bent or angular (105o) • SF6 types  form octahedral (90o)

  13. Hybridization • Atomic orbitals mix & form equal hybrid orbitals on the same atom • s & p orbitals  sp orbital (180o) • BeF2 • s, p, & p orbitals  sp2 orbital (120o) • BF3 • s, p, p, & p orbitals  sp3 orbital (109.5o) • CCl4

  14. Intermolecular Forces • Dipole-dipole forces • Separated equal but opposite chargesDipole • Represented by arrow with head toward (-) pole • Forces of attraction between polar molecules • Hydrogen bonding • H of 1 molecule pulled to (-) charge on another • London dispersion forces • Attractions from creation of instantaneous dipoles

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