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Chemical Bonding

Chemical Bonding. Bonds form in 2 main ways atoms share electrons electrons are transferred between atoms Type of bond depends on the atom’s electronegativity and electron configuration. 3 Main Types of Bonds. Ionic Bonds electrostatic force atoms transfer e - to become ions

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Chemical Bonding

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  1. Chemical Bonding

  2. Bonds form in 2 main ways • atoms share electrons • electrons are transferred between atoms Type of bond depends on the atom’s electronegativity and electron configuration

  3. 3 Main Types of Bonds • Ionic Bonds • electrostatic force • atoms transfer e- to become ions • usually between a metal and a nonmetal • Covalent Bonds • electrons are shared by atoms • usually between two nonmetals

  4. Metallic Bonds • forces that hold metals together • metals have many freely moving electrons that attract positive metal ions • Ex. What type of bonding would exist in solid aluminum?

  5. Ionic Bonding valence electrons: outermost s and p electrons of an atom Dot Diagrams: show valence electrons Examples: isoelectronic: having the same electron configuration

  6. Characteristics of Ionic Compounds • high melting point • able to conduct electricity in molten state • tend to be water soluble • crystallize in definite patterns (crystal lattice)

  7. A closer look at ionic bonding

  8. Naming BINARY Ionic Compounds • name the metal first – do not change ending • name nonmetal second – change ending to –ide • Examples:

  9. Writing Formulas for Ionic Compounds • Sum of all ion charges MUST equal ZERO! • Use the “criss-cross” method • Examples:

  10. Covalent Bonding molecule: name for a covalently bonded particle

  11. Characteristics of Covalent Compounds • low melting point • do not conduct electricity • usually brittle solids, liquids, or gases

  12. A closer look at covalent bonding

  13. Naming Covalent Molecules • make sure the bond is covalent (usually 2 nonmetals) • first element’s name does not change • second element’s ending becomes –ide • Use Greek prefixes to indicate the # of atoms of each element mono = 1 di = 2 tri = 3 tetra = 4 penta = 5 hexa = 6 hepta = 7 octa = 8

  14. Writing Formulas for Covalent Molecules • Prefixes tell the # of atoms of each element • Examples:

  15. Molecular Geometry VSEPR Theory

  16. Lewis Structures • Show arrangement of atoms in molecules • Show shared (bonding) and free electrons

  17. Drawing Lewis Structures Used for covalently bonded molecules ONLY! • Determine the atoms in the molecule • Count valence electrons for each atom. • Find total # of valence electrons • Arrange atoms in skeleton structure. *Least electronegative atom in center!* • Add electrons to structure.

  18. The number of covalent bonds normally formed by an atom in a Lewis structure depends on its group in the periodic table. • H is expected to form one bond. • F, Cl, Br, I, all in group 17 are expected to form one bond each. • O, S, Se, in group 16, are expected to form two bonds each. • N, P, As, in group 15, are expected to form three bonds each. • C, Si, Ge, in group 14 are expected to form four bonds each.

  19. Octet Rule • Atoms try to achieve Noble Gas configuration (8 outer e-) • Hydrogen – forms “duet” instead • Some atoms exceed octet – more than 8 bonding e-

  20. VSEPR Theory Valence Shell Electron Pair Repulsion Theory • electron groups arranged to minimize repulsion

  21. Molecular Shapes FILL IN SHAPE CHART! • Show relative positions of atomic nuclei • MUST determine Lewis structure to determine shape!

  22. Resonance • Equivalent Lewis structures • Shows possible locations of double bonds

  23. Resonance

  24. Polarity of Bonds • polar: having opposite ends • polar bond: e- shared unequally • caused by difference in electronegativity • nonpolar bond: e- shared equally

  25. ALL COVALENT BONDS ARE POLAR EXCEPT: • C – H • any atom bonded to itself

  26. A closer look at polar bonds

  27. Polarity of Molecules • Bonds must be polar for molecule to be polar. • Molecule must have a definite top and bottom with opposite charges in order to be polar.

  28. Intermolecular Forces (Weak Bonds) Three main types • Dispersion forces (London, van der Waals) • Dipole-Dipole Interactions • Hydrogen bonding

  29. Dispersion Forces (van der Waals) • Very weak • Between nonpolar molecules • Induces momentary (temporary) dipole • Ex. – occurs in Cl2, CO2, CH4, etc.

  30. Dipole-Dipole Interactions • Stronger than dispersion • Occur between molecules with permanent dipoles (aka – polar) • Partially + end of one molecule attracted to partially – end of another

  31. Hydrogen bonding • Stronger type of dipole-dipole interactions • Results from H being covalently bonded to either F, O, or N • Stronger because… • H is so small • F, O, & N are very EN • Partial +/- charges are more intense

  32. Intermolecular Forces H-bonds > Dipole-dipole > Dispersion • Affect BP, MP, solubility • More E required to boil/melt substances w/ stronger intermolecular forces • Why?

  33. Pop Quiz • Name CaCO3 • Write a formula for sodium sulfite. • Draw a dot diagram for boron (B). • How many valence electrons does carbon have? • What is the oxidation number of potassium?

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