1 / 26

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”.

page
Download Presentation

Molecular Compounds

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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

More Related