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Molecular Compounds

Molecular Compounds. Binary molecular compounds = 2 nonmetals covalently bonded No ions, so cannot use ionic naming Can combine in more combinations than ions  need a new system Prefix: tells # of atoms of each element are present All end in “–ide”.

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Molecular Compounds

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  1. Molecular Compounds • Binary molecular compounds = 2 nonmetals covalently bonded • No ions, so cannot use ionic naming • Can combine in more combinations than ions  need a new system • Prefix: tells # of atoms of each element are present • All end in “–ide”

  2. You need to remember Table 8.3 (p248) • CO versus CO2 • N2O = laughing gas • If there is one atom of first element, then you do not need to put “mono-”

  3. Naming Molecular Compounds • Make sure it is a binary molecular compound (2 nonmetals) • Name elements in order • Use prefixes to designate # atoms • If 1st element has only one atom, then omit “mono-” • End 2nd element w/ “-ide”

  4. Practice 8-1 Write the names for the following molecular compounds: • SF6 • Cl2O8 • N2O6 • PF6 • Br2Cl

  5. Practice 8-2 Write the formulas for the following molecules: • Silicon monocarbide • Dinitrogen tetroxide • Selenium hexafluoride • Sulfur decabromide • Hexaphosphorus tricarbide

  6. Practice 8-3 1. Write the formulas for the following compounds: • ammonium dichromate • tin(II) perchlorate • carbon tetrafluoride • triiodine pentabromide • heptanitrogen nonaphosphide 2. Write the names for the following compounds: • CrPO4 • CuC2H3O2 • Se3S • N2F5 • P3Cl2

  7. Why do atoms bond? • Gaining or losing electrons makes atoms more stable by forming ions with noble-gas electron configurations (ions  ionic compounds) • Sharing valence electrons with other atoms also results in atoms having noble gas configurations (covalent bonds  molecules)

  8. Molecules • Nonmetals share electrons • Covalent bond: chemical bond caused by atoms sharing electrons • Molecule: atoms joined by covalent bonds • Diatomic: made of 2 atoms • 7 elements that exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2

  9. Formation of Diatomic Elements

  10. Single Covalent Bonds • When one pair of electrons is shared between 2 atoms

  11. Multiple Covalent Bonds • Double bonds: when 4 electrons are shared between two atoms • Triple bonds: when 6 electrons are shared between two atoms

  12. Strength of Covalent Bonds • Bond dissociation energy: amount of energy needed to break a bond • Stronger bonds need more energy to be broken • Double bonds are stronger than single bonds • Triple bonds are stronger than double bonds

  13. Length of Covalent Bonds • In stronger bonds, the atoms are closer together • Double bonds are shorter than single bonds • Triple bonds are shorter than double bonds

  14. Electronegativity • How much an atom pulls electrons in a covalent bond towards itself • Noble gases are not listed because they do not form compounds

  15. Nonpolar Covalent Bonds • Electrons are pulled equally by the atoms • Nonpolar molecules are not attracted to an electric field • Examples: H2, Cl2, O2

  16. Polar Covalent Bonds • Electrons are pulled unequally to the atoms in a covalent bond • The atom with the higher electronegativity value pulls the electrons more • Polar molecules align in an electric field • Dipole: polar molecule

  17. Electronegativity Difference and Bond Type • We can calculate the difference in EN to determine whether a bond is a nonpolar, polar, or ionic.

  18. Higher electronegativity = pulls electrons more • δ = partial charge (<1) • Polar molecules: one end of molecule is δ- and the other end is δ+ • Can also be shown w/ arrow pointing towards more electronegative atom

  19. Practice 8-4 Determine the bond type that exists between the following atoms. Show your work! • C-H • Na-F • Br-Cl • H-F • Fe-O

  20. Representing Molecules

  21. Lewis Structures • Lines represent covalent bonds (2 electrons), dots represent electrons • To draw Lewis Structures: • Use the periodic table to add up all valence electrons in the substance (these e- are available for bonding)  if it’s an ion then add or remove e- accordingly • Join all atoms by a single bond • Put in lone pairs of electrons • Count up electrons and make sure all atoms obey the octet, if they do not, then make double/triple bonds when necessary

  22. Exceptions to the Octet • H can only have 2 electrons MAX • Be can have 4 electrons • B can have 6 electrons • Any nonmetal in Period 3 or below can have more than 8 electrons, but it must be an even # of electrons

  23. Lewis Structure Example Draw Lewis structures for the following molecules: • CH4 • NH3 • CCl4 • CO2

  24. Practice 8-5 Draw Lewis structures for the following: • water • BF3 • SO22- • H3O+ • sulfur dichloride • C2H6

  25. Practice 8-6 Draw Lewis structures for the following: • Carbonate ion • Carbon dioxide • Sulfur trioxide • Sulfur hexafluoride • Carbon monoxide

  26. Practice 8-7 Draw Lewis structures for the following: • O3 (ozone) • Boron trifluoride • Phosphorus pentafluoride • Sulfur hexafluoride • Nitrate ion

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