1. What is the lewis theory?

    -according to the lewis theory, when are chemical bonds formed?
    a model for chemical bonding.

    -when atoms transfer valence electrons (ionic bonding) or share valence electrons (covalent bonding) to attain NOBLE GAS electron configurations.

    • octet=8
    • duet=2 (as in the case of hydrogen)
  2. The shapes of molecules (Molecular shapes) can be predicted by combining lewis theory w/ what other theory?

    What happens in this model?
    Valence Shell Electron Pair Repulsion (VSEPR) theory.

    *In this model, electron groups-lone pairs, single bonds, double bonds, and triple bonds-around the central atom repel one another and determine the geometry of the molecule.
  3. What is an example of molecular shapes, with water (H2O)?
    Water has a bent geometry causing it to be a liquid at room temp. instead of a gas. It is also the reason ice floats on water and snowflakes have hexagonal (6 sided) patterns.
  4. What is Electronegativity and how is the power shifted through the periodic table?
    Electronegativity: the relative ability of elements to attract electrons within a chemical bond.

    *Electronegativity INCREASES as you move to the right and DECREASES as you move down a column.
  5. What is polarity?
    When two nonmental atoms of different electronegativities form a covalent bond, the electrons in the bond are not evenly shared & the bond is POLAR.

    *In molecules w/ more than 2 atoms, polar bonds may cancel, forming a nonpolar molecule, or they may sum, forming a polar molecule.
  6. What is the relevance of the polarity of a molecule?
    It influences many of its properties such as whether it will be a solid, liquid, or gas at room temp. and whether it will mix w/ other compounds.

    *Oil and water, for example, do not mix because water is polar while oil is nonpolar.
  7. 10.12 What is, the lewis structure for elements?
    The lewis structure of any element is the symbol for the element with the valence electrons represented as dots drawn around the element.

    The # of valence electrons is equal to the group number of the element (for main-group elements).
  8. 10.13 On Writing Lewis Structures of Ionic Compounds.
    In an ionic lewis structure, the metal loses all of its valence electrons to the nonmental, which attains an octet. The nonmental, w/ its octet, is normally written in brackets w/ the charge in the upper right corner.

    Ex: (for lithium bromide) LiBr

    Answer: Li+ [ Br ]- (Br has 8 v. e- to obtain an octet)
  9. 10.14 Using Lewis Theory to predict the chemical formula of an ionic compound.

    Q: Use Lewis Theory to predict the formula for the compound that forms between potassium and sulfur.
    Solution: The lewis structure of K and S are:

    K (1 v. e-), S (6 v. e-)

    K must lose one e- and sulfur must gain two. Consequently, we need two K atoms to every sulfur atom. The Lewis structure is:

    • K+[S] 2- K+ (S has an octet, 8 e-)
    • (formula = K2S, two K for every sulfur)

    1) To determine the chemical formula of an ionic compound, write the lewis structures of each of the elements.

    2) Choose the correct # of each type of atom so that the metal atom(s) lose all of their v. e- and the nonmetal atom(s) attain an octet.
  10. 10.15 Writing Lewis structures for covalent compounds

    Q: Write the Lewis structure for CS2.
    • Solution: S C S (C=4, S=6), 2 Sulfur atoms
    • 1) Write the correct skeletal structure for the molecule.

    • total e- = S (6 v. e-) x 2 atoms + C (4 v. e-) = 16
    • 2) Caluculate the total # of e- for the lewis structure by summing the valence e- of each atom in the molecule.
    • 2.2) Remember that the * of v e- for any main-group element is equal to its group # in the periodic table. For polyatomic ions, add one e- for each negative charge and subract one e- for each positive charge.

    • S:C:S (4 of 16 e-used)
    • 3) Distribut the e- among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible.
    • 3.3) Begin by placing two e- between each pair of atoms. These are the bonding e-. Then distribute the remaining e-, first to terminal atoms and then to the central atom.

    :S:C:S: (16 of 16 e- used, 4 e- top/bottom of both S to total 16 used)

    :S::C::S: (with 2 v. e- above S's)


    • :S==C==S: (2 double bonds, w/ 2 v. e- above the S's)
    • 4) If any atoms lack an octet, form double or triple bonds as necessary to give them octets.
    • 4.4) Do this by moving lone e- pairs from terminal atoms into the bonding region w/ the central atom.
  11. 10.16 Writing Resonance Structures

    Q: Write resonance structures for SeO2
    S: We can write a lewis structure for SeO2 by following the steps for writing covalent Lewis structures. We find that we can write two equally correct structures, so we draw them both as resonance structures.

    :O (2 v. e- t/b) --Se (1 lone pair)==O: (2 v. e- above O)


    :O (1 lone pair above) == Se (1 lone pair above) -- O: (2 lone pairs, one above/below, O)

    *When you can write two or more equivalent (or nearly equivalent) Lewis structures for a molecule, the true structure is an average between these. Represent this by writing all of the correct structures (calle dresonance structures) with double-headed arrows between them.
  12. 10.17 Predicting the shapes of molecules

    Q: Predict the geometry of SeO2
    S: The lewis structure for SeO2 is composed of the following two resonance structures.

    :O (2 v. e- t/b) --Se (1 lone pair)==O: (2 v. e- above O)and:O (1 lone pair above) == Se (1 lone pair above) -- O: (2 lone pairs, one above/below, O)

    Either of the resonance structures will give the same geometry.

    • Total # of e- groups = 3
    • # of bonding groups = 2
    • # of lone pairs = 1

    • Electron geometry = Trigonal planar
    • Molecular geometry = Bent

    1) Draw the lewis structure for the molecule.

    • 2) Determine the total # of electron groups around the central atom.
    • 2.2) Lone pairs, single, double, and triple bonds each count as one group.
    • 3) Determine the # of bonding groups and the # of lone pairs around the central atom.
    • 3.3) These should sum to the result from step 2. Bonding groups include single, double, and triple bonds.
    • 4) Refer to table 10.1 to determine the electron geometry & molecular geometry.
  13. 10.18 Determing whether a molecule is polar.

    Q: Determine whether SeO2 is polar.
    S: Se & O are nonmetals w/ different electronegativities (2.4 for Se and 3.5 for O). Therefore, the Se-O bonds are polar. The geometry of SeO2 is bent.

    *The polar bonds do not cancel but rather sum to give a net dipole moment. Therefore the molecule is polar.

    • 1) Determine whether the molecule contains polar bonds.
    • 1.1) A bond is polar if the two bonding atoms have different electronegativities. If there are no polar bonds, the molecule is nonpolar.

    2) Determine whether the polar bonds add together to form a net dipole moment. Use VSEPR theory to determine the geometry of the molecule. Then visualize each bond as a rope pulling on the central atom. Is the molecule highly symmetrical? Do the pulls of the ropes cancel? If so, there is no net dipole moment & the molecule is nonpolar. If the molecule is asymmetrical & the pulls of the rope do not cancel, the molecule is polar.
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