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Covalent Bonding

Covalent Bonding. Ionic bond - a metal bonded to a non-metal, electrons are transferred Covalent bond - a non-metal bonded to another non-metal, electrons are shared, remember molecules try to arrange themselves into octets using each others valence electrons

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Covalent Bonding

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  1. Covalent Bonding

  2. Ionic bond - a metal bonded to a non-metal, electrons are transferred Covalent bond - a non-metal bonded to another non-metal, electrons are shared, remember molecules try to arrange themselves into octets using each others valence electrons Molecule - formed when two or more atoms are bonded covalently Diatomic Molecules - H2, N2, O2, F2, Cl2, Br2, and I2 form a 7 on the periodic table. These elements only exist as pairs in nature. They are more stable ex. F2 - has seven valence electrons and must have one more to be "happy" As 2 F atoms approach each other they both have a negative charge so they repel each other, as they move even closer the attraction of the nuclei for the other atoms electrons becomes large enough to over power the repelling forces

  3. Single Bonds - when a single pair of electrons are shared ex. H2 Lewis structures - use electron-dot diagrams to show how the bonds are formed ex. H20 each H atom attains noble gas configuration of He by sharing electrons with O, and O attains noble gas configuration of Ne by sharing electrons with 2 H's

  4. Sigma Bond - single covalent bonds · occurs when the electron pair is shared in an area centered between the two atoms · when the valence atomic orbital of one atom overlaps with the valence atomic orbital of the other atom · used to identify where the electrons will most likely be found

  5. Multiple Bonds - many molecules form double, or triple bonds to gain noble gas configuration ex. O2 N2 Pi Bonds - · formed when parallel orbitals overlap to share electrons · the shared electron pair of a pi bond occupy the space above and below the line that represents where the two atoms are joined together · Consists of one sigma bond and at least one pi bond

  6. Strength of a Covalent Bond ·Remember that the bond has both attractive and repulsive forces so it can be broken easier than an ionic bond ·Covalent bonds strength depends on how much distance separates the nuclei ·Double and Triple bonds are stronger than single bonds because the atoms are closer together ·Energy is released when a bond is formed and energy must be added when a bond is broken ·Bond Dissociation energy - the energy required to break a covalent bond (always a positive value) Indicates the strength of a chemical bond ·Endothermic reactions occur when a greater amount of energy is required to break the existing bonds then is released when the new bonds form ·Exothermic reactions occur when more energy is released forming new bonds than is required to break bonds in the initial reactants

  7. Molecular and Acid Naming Ch 6 Lecture Notes 07-08.notebook

  8. Molecular Structures - uses letters and lines to show bonds

  9. Coordinate Covalent Bond - Forms when one atom supplies both electrons to a shared pair - Occurs in many polyatomic ions - Example: NH3 & H3B Properties of Covalent Compounds - Made of molecules - Have lower melting points Resonance - When two or more structures can be written for a molecule - The actual structure is a blend of the structures - Example: NO3-

  10. Hybridization - Two or more orbitals (of nearly equal energy) combine into orbitals of equal energy - Hybridized orbitals orient themselves as far apart from one another as possible - Helps to explain the shape of molecules - VSEPR Theory (Gold Sheet) - Valence Shell Electron Pair - ex. C has 4 valence e- [He] 2s2 2p2

  11. Bond Polarity Nonpolar Covalent Bonds - Bonding electrons are shared equally - Electronegativity difference of 0.4 or less - Example: Polar Covalent Bonds - Bonding electrons are shared unequally - Electronegativity difference of 0.4 - 2.0 - The more electronegative atom has a slight negative (δ-) charge - The less electronegative atom has a slight positive (δ+) charge - Example: H2O

  12. Molecular Polarity - Depends on bond polarity and shape of molecule - If bonds are nonpolar, then molecule is nonpolar - Example: - If bonds are polar but molecule is symmetrical, then molecule is nonpolar - Example: CH4 - If bonds are polar and molecule is not symmetrical, then molecule is polar - Called a “dipole” - Example:

  13. WS examples: Use your Gold Sheet!!!!!!! SiCl4 - Bond Skeleton Hybrid Type 3-D structure

  14. TEST!!!!! 10 Molecular Naming 10 Acid Naming 3 True False 5 MC Vocab - structural formula sigma bond molecule polar covalent VSEPR model pi bond coordinate covalent covalent bond hybridization resonance oxyacid endothermic electronegativity exothermic 3 Drawings

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