Chapter 8 - Chemical Bonding

• Lewis Symbols and the Octet Rule
• Valence electrons - outer shell electrons
• Lewis dot structures
• dots represent electrons in orbitals
• use for representative elements only
• Dot numbers = group numbers (for representative elements)
• Octet rule
• species tent to have 8 electrons in the outer shell
• lose or gain to form ions
• share in covalent bonds
• H, He have 'octet' with 2 electrons.
• Ionic Bonding
• Transfer of electrons
• Held together by electrostatic attraction
• Tend to have noble gas configuration of electrons
• pseudo-noble gas configuration
• Transition metals
• ions are formed by losing electrons from the outermost shell
• In Cu2+ the two 4s electrons are lost rather than any 3d electrons

• Sizes of ions
• Down a family - size increases
• Isoelectronic species - size increases as charge becomes more negative
• Negative ions larger than neutral atom
• Positive ions smaller than neutral atom

• Covalent bonding
• electrons are shared
• shared electrons count toward octet for both atoms that are sharing
• visualize with Lewis structures
• multiple bonding
• single - share 2 electrons
• double - 4 electrons shared
• triple - 6 electrons shared
• Bond polarity and Electronegativity
• pure covalent - electrons are shared equally
• most covalent bonds are polar
• electrons shared unequally
• determine with electronegativity
• tenancy of an atom in a molecule to attract electrons in a covalent bond.
• trends follow ionization energy trends
• greater the difference in electronegativity the more polar the bond

• Drawing Lewis structures
• N = electrons needed to complete an octet (8*atoms except H) + (2*H atoms)
• S = valence electrons from each atom + (1*negative charge)
• A = N - S
• # of electrons needed for bonding
• bonds = S/2
• method will work for all species which follow the octet rule
• impossible results implies non-octet compound
• Multiple bonds
• must involve O, N, or C
• Ternary Acids - H bonded to O
• Alternate method in textbook
• Resonance structures
• Arbitrary placement of multiple bonds - resonance structures
• 'True' structure intermediate of resonance forms

• Exceptions to the Octet Rule
• Odd number of electrons
• ClO2, NO, NO2
• Less than an octet
• BF3
• More than an octet
• most common
• Draw Lewis structure w/'obvious' bonding, and then 'split' electron pairs to form bonds 2 at a time.
• occurs because of use of d orbitals (Chapter 9)

• Oxidation Number
• electronic bookkeeping
• numbers assigned to atoms
• Assigned by a set of rules
• Element in elemental form = 0
• Monatomic ion = ionic charge
• Fluorine = -1 in all compounds
• Oxygen = -2 except
• OF2, O = +2
• peroxide - O = -1
• superoxide - O = -
• H = +1 except
• hydride (bonded to a metal) = -1
• Halogen = -1 except when bonded to a more electronegative element
• For a molecule - sum of oxidation numbers = 0
• For a polyatomic ion - sum of oxidation numbers = ionic charge

Learning Goals

1. Determine the number of valence electrons for any representative element, and write its Lewis structure.

2. Recognize when the octet rule applies to the arrangement of electrons in the valence shell of an atom.

3. Predict on the basis of the periodic table the probable formation of ionic substances formed between common metals and nonmetals.

4. Describe how the radii of atoms relates to those of ions.

5. Explain the concept of isoelectronic species and the origin of changes in ionic radii of such species.

6. Describe general differences in physical properties between substances with ionic bonds and those with covalent bonds.

7. Describe the basis of Lewis theory, and predict the valence of common nonmetallic elements from their position in the periodic table.

8. Explain the significance of electronegativity, and in a general way relate the electronegativity of an element to its position in the periodic table.

9. Predict the relative polarities of bonds using either the periodic table or electronegativity values.

10. Using the periodic table, write Lewis structures for molecules and ions using covalent bonds.

11. Write resonance forms for molecules or polyatomic ions that are not adequately described by a single Lewis structure.

12. Write the Lewis structure for molecules and ions containing covalent molecules that have a deficiency of electrons or an excess of electrons (also understand the possibility of odd numbers of electrons).

13. Assign oxidation numbers to atoms in molecules and ions.

14. Assign acceptable names to simple inorganic compounds and ions.

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