CHAPTER 8: BASIC CONCEPTS OF CHEMICAL BONDING Bond-an attractive interaction between two or more atoms. Bonding is the "glue" that holds molecules together. Two extreme types: Ionic (transfer) Covalent (sharing) Polar covalent is in between
LEWIS symbols are a method of using dots to represent VALENCE (those past noble gas core) electrons. #valence electrons=group number
Ionic Bond- the electrostatic attraction between a cation (metal) and anion (non-metal) resulting from transfer of an electron. Covalent Bond-an attractive force resulting from sharing of electron pairs (1-single bond, 2-double bond, or 3-triple bond).
Both try to achieve the Lewis octet (isoelectronic to noble gas) if possible. Illustrate Ionic bond in LiF using Lewis Symbols and equations, do the same for the covalent F 2. Look at Energetics of Ionic cpd formation (Born- Haber, Hess Law) MX(s) M + (g) + X - (g) H lattice >0
Lattice E increases as charges on ions increases and as ion radius decreases
Electron Configuration of Ions are predicted by adding electrons (nonmetals to give anions) or removing electrons (metals to give cations) from the electron
configuration of the neutral atom. Isoelectronic - Species which have the same electron configuration. Main Group Ions form to achieve an electron configuration isoelectronic to the noble gases. Transition Metals lose ns electrons before (n-1)d electrons
Periodic Properties -Atomic Size -Ion Size -Ionization Energy -Electron Affinity -Electronegativity Ion Size Trends -Cations are ALWAYS smaller than the metals from which they were formed -Anions are ALWAYS larger than the non-metals from which they were
formed
Covalent Bond-an attractive force resulting from sharing of electron pairs (1-single bond, 2-double bond, or 3-triple bond). Draw Lewis Structures of a) CH 4, NH 3, H 2 O, and HF b) NH 4 +, and OH - c) F 2, O 2, N 2, CO 2, and HCN d) BeCl 2, BCl 3, PF 5, SF 6
Lewis Structures (p314) 1. Count up valence e -. # valence e - =Group Number Be sure to consider charge. 2. Arrange valence e - in pairs around central atom so that H has 1 pair and C, N, O, and F have 4 pairs (Lewis Octet). You may have to use double or triple bonds or nonbonding (lone pair)(s)- a pair of electrons located
on one atom. 3. Groups II and III (Be and B) cpds. may have less than a Lewis octet. 4. Elements beyond the 2 nd period (P, S, Cl) may have more than a Lewis octet. (they have unused "d" orbitals which may participate in bonding) A Lewis structure by itself DOES NOT!!!!! Describe the shape of a molecule. It
simply accounts for the electron distribution. Exceptions to octet rule -odd electron species -less than an octet (Be, B) -more than an octet( 3rd period elements) Free Radicals are paramagnetic odd electron species that are very reactive and cannot satisfy Lewis Octet. Examples are NO, NO 2, and Cl.
Draw Lewis structure of NO 2 - Resonance structures - possible structures of a species that has more than one acceptable Lewis structure (these differ only in location of double or triple bonds). (horse and donkey) Resonance hybrid- a composite structure based
on the contributing resonance structures that approximates the true structure. (mule) Draw the resonance structures of CO 3 2- and its resonance hybrid. Give the formal charge (FC) on each atom. (p316) FC=GN - # lone pair e ½ # bonding e Exceptions to octet rule -odd electron species
-less than an octet (Be, B) -more than an octet( 3rd period elements) Bond Enthalpy- the energy required to break a bond (endothermic) A-B A + B Bond Length: triple < double < single
Bond Strength: triple > double > single
Single bonds are longer and weaker than double or triple bonds Use the data in Table 8.4 to
calculate the H of the following reaction H 2 + Cl 2 2HCl H = sum of bonds brokensum of bonds formed Charge distribution (even or uneven sharing?) Electronegativity - the ability of an atom in a molecule to attract a pair of electrons to itself. Increases left to right and
bottom to top. (F most electronegative) Polar Bond - a bond between atoms of different electronegativities that results in unequal charge distribution.
Dipole Moment, µ, is a measure of the polarity of a molecule. A species is polar if: 1) it has polar bonds and 2) the arrangement of the polar bonds does not result in cancellation