Chem Ch.7

  1. What 2 factors account for the stability of the H2 molecule?
    • The fact that the attractive energy between the electrons and protons of opposite atoms is slightly greater than the repulsive energy between the proton-proton and electron-electron interactions.
    • The fact that the electrons are spread over the entire volume of the H2 molecule itself and not just the confines of it's atom's orbital
  2. What are the electrons in the outer-most principal energy level shell of an atom called?
    Valence Electrons
  3. In a Lewis Dot Diagram of an atom, which electrons of the atom are shown as dots around it?
    Are they the only ones involved in covalent bonding?
    • Only the valence electrons of the atom
    • Yes
  4. All molecules form covalent bonds to achieve the electron configurations of which atoms?
    Noble Gases
  5. How can you identify the amount of valence electrons in an atom on the periodic table?
    It is equal to the last digit of the group number the atom is in on the periodic table
  6. What are the 2 types of pairs possibly shown on a Lewis Structure and how are they represented?
    • Covalent Bond pairs
    • Shown as dash between the bonded atoms
    • Unshared pairs
    • Shown as separate dots outside the corresponding atom
  7. For molecules that are ionized despite covalent bonds, how are they're charges shown in Lewis structures?
    • The bonded atoms are placed in brackets and the charge is shown in the top right corner outside the bracket.
    • Ex: [ ]2-
  8. What does the octet rule state?
    What is its exception?
    • Says that atoms in covalently bonded species tend to have the electron structures of noble gases.
    • The exception is hydrogen, which need only be surrounded by 2 electrons to be stable.
  9. Which elements never form multi-covalent bonds in Lewis structures?
    Hydrogen and the halogens
  10. What is the procedure for writing Lewis structures?
    • 1) Count # of valence electrons
    • if polyatomic ion add/stubtract # of electrons equal to its charge
    • 2) Draw skeleton structure using covalent bonds
    • # of electrons to use for covalent bonds is equal to # needed for noble gas structure - # of valence e-
    • 3) Determine # of electrons still available
    • equal to # of electrons - 2 for each covalent bond
    • 4) Fill each atom with remaining unpaired electrons until octets are reached
    • remember hydrogen only needs two electrons to be stable
  11. How can you denote resonance in a Lewis structure?
    Place a double arrow between the two resonant structures
  12. Does resonance imply that there are 2 different kinds of the same molecule?
    No, the molecule is intermediate between the two forms
  13. What is the "charge an atom would have if valence electrons in bonds were distributed evenly"?
    Formal Charge
  14. If you found it possible to write more than one Lewis structure arrangement for a molecule, how would you go about telling the difference between which one would exist in nature?
    Find the formal charges of each important atom so that the more realistic arrangement of molecules would be the one where the formal charges are closest to 0
  15. What is the equation for formal charge and what doe each variable represent?
    • Cf = X-(Y + Z/2) = X-Y-Z/2
    • X = # of valence electrons in the atom alone
    • Y = # of unpaired electrons in the atom within the Lewis structure
    • Z = # of electrons used to form the covalent bonds around the atom
  16. Where should negative formal charges be in a set of formal charge equations?
    Should be on the element that is most electronegative
  17. When do there tend to be exceptions to the octet rule of Lewis Structures?
    • 1) When the amount of valence electrons in the molecule add up to an odd number
    • aka: free radicals
    • 2) When the octets must be expanded to fit more than 4 covalent bonds
  18. When are there typically Expanded Octets?
    • When a nonmetal of the Third, Fourth, or Fifth period is combined with a halogen
    • These nonmetals have d orbitals available for extra electron pairs
  19. How would you know that an Expanded Octet is needed for a Lewis Structure?
    If you find that the number of electrons needed to fill all octets is less than the amount of electrons remaining after the skeleton structure is drawn, then you need to add the left over electrons to the central atom
  20. What does the VSEPR model say about electron pairs surrounding an atom?
    It says that the pairs repel one another, and through this repulsion they are oriented to be at angles that are as far away from each other as possible
  21. For the notation AXbEc, what does each part of that designate?
    • A - Central atom in the molecule
    • X - Terminal atom surrounding the central atom
    • b - Number of terminal atoms present
    • E - Number of unshared electron pairs that central atom has
    • c - Nnumber of these unshared pairs
  22. What is the shape of a molecule with the formula AX4 and what angle are its covalent bonds at? Also give an example.
    • Tetrahedron shape and the bonds
    • are 109.5 degrees apart
    • CH4 is like this
  23. Tell me what orientation the Electron pairs in each of these molecules have:
    • Linear
    • Trigonal planar
    • Tetrahedron
    • Trigonal Bipyramid
    • Octahedron
  24. For a molecule with the formula AX2, if the central atom has unshared electrons then what shape will it take?
    What about if the central atom in AX3 had unshared electrons?
    • Bent
    • Trigonal pyramid
  25. What are the predicted VSEPR model shapes of each of these molecule types:
    • Trigonal bipyramid
    • See-Saw
    • T-Shape
    • Linear
    • Octahedral
    • Square pyramid
    • Square planar
  26. Do multiple bonds affect the predicted shape of molecule? How do they act?
    They don't affect the shape because they occupy space the exact same way a single bond would.
  27. What are the 2 types of covalent bonds?
    • Polar Covalent bonds
    • These are a result of an asymmetrical distribution of electrons that causes a positive and negative pole
    • Non-polar Covalent bonds
    • The distribution of electrons in these are symmetrical and thus have no poles
  28. When do polar and non-polar bonds typically occur?
    • Non-polar bonds happen when the atoms joined are identical or the same element (as in H2 or F2).
    • Polar bonds happen when the atoms joined aren't of the same element thus creating a higher electron density on one side than the other.
  29. How do you find whether or not a covalent bond is strongly polar?
    Calculate the difference in electronegativities of the bonded atoms
  30. What is the dipole moment?
    This is a measure of the extent to which molecules tend to orient themselves in a magnetic field.
  31. What is the dipole moment of F2?
  32. How many hybrid orbitals are created by an sp3d hibridization?
    5 orbitals
  33. True/False: All single bonds are sigma bonds?
  34. In a multiple bond, how many bonds are sigma bonds and how many are pi bonds?
    Only one bond is a sigma bond...the rest are pi bonds.
  35. Which two elements are known to need less than 8 electrons to reach stability, and how many electrons do they need specifically?
    • Be (4e)
    • B (6e)
  36. Are pi bonds made up of hybrid orbitals?
  37. If you have a molecule with 2 single bonds and 1 double bond, what is the hybridization of its orbitals?
    sp2 because two p orbitals are combined with an s orbital to create 3 sigma bonds, leaving one p orbital available for the pi bond that creates the double bond.
Card Set
Chem Ch.7
Covalent Bonding