I. Types of Chemical Bonds

1. I. Types of Chemical Bonds • Ionic Bonding • Chemical bond made by attraction of oppositely charged atoms after the transfer of electrons between them • 2 Na + Cl2 2 Na+Cl- (ionic compound) • Metals lose electrons to become cations • Na loses 1 e- [Ne]3s1 [Ne] • Nonmetals gain electrons to become anions • Cl gains 1 e- [Ne]3s23p5 [Ar] • The +/- charge attraction in the NaCl molecule makes it • 8.37 x 10-19 J more stable than 2Na and Cl2 apart. (Coulomb’s law) • Ionic compounds stay bound together as solids, but dissolve in water to give separated cations and anions • NaCl(s) + H2O(l) Na+(aq) + Cl-(aq) • 3. Ionic Lewis Structure:

2. B. Formation of Binary Ionic Compounds • 1. Lattice Energy = DE when separated gaseous ions pack to form a solid • a. Exothermic process since attraction of ions is favorable • b. Lattice Energy has a negative sign • c. Calculate by summing steps: Example Li(s) + ½ F2 ----> LiF(s) • 2. The Born-Haber Cycle: • C. Lattices and Coulomb’s Law • Larger charges (+2/-2) attract more • CaO has > 4x the lattice E of NaF 1. Li(s) ----> Li(g) DE = +161kJ/mol 2. Li(g) ----> Li+(g) + e- DE = +520kJ/mol 3. ½ F2(g) ----> F(g) DE = +77kJ/mol 4. F(g) + e- ----> F-(g) DE = -328kJ/mol 5. Li+(g) + F-(g) ----> LiF(s) DE = -1047kJ/mol Li(s) + ½ F2(g) ----> LiF(s) DE = -617kJ/mol

3. Making Mg2+ and O2- ions requires • More energy than making Na+ and F- • Attractive forces of MgO >> NaF • Lattice energy 4x greater, but overall • energy difference is small

4. Predicting Formulas • D. Noble Gas Electron Configurations (NGEC) • All elements react by gaining, losing, or sharing electrons to achieve the electron configuration of a noble gas (filled electron shells) • Two nonmetalsreact: They share electrons to achieve NGEC. • a. H (1s1) + H (1s1) H2 (1s2 = [He]) • b. 2 Cl ([Ne]3s23p5) Cl2 ([Ar]) • A nonmetal and a representative group metalreact (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC. • a. Na loses 1 e- [Ne]3s1 [Ne] • Cl gainis 1 e- [Ne]3s23p5 [Ar] • Ca lose 2 e- [Ar]4s2 [Ar] • O gains 2 e- [He]2s22p4 [Ne] NaCl CaO

5. Duet Rule: H and He always react to have 2 e- in outer (n =1) shell • Octet Rule: All other main group elements react to have 8 e- in outer shell • Transition metals and heavier metals: Harder to predict • E. Use Periodic Table to predict charges and formulas • MgxOy AlxSy CaxCly • F. Ion sizes: More negative ions are larger, more positive are smaller • Larger attraction for fewer electrons contracts cations: Ca > Ca+ > Ca2+ • Less attraction for more electrons enlarge anions: O2- > O- > O • O2> F > Na+ > Mg2+ > Al3+ • Ions get larger down a group because larger shells are involved

6. Sizes of Ions Related to Positions of the Elements in the Periodic Table

7. G. Trends in Lattice Energies • Increased charge leads to larger lattice energies • Smaller ions lead to larger lattice energies • H. Properties of Ionic Compounds • High melting Points—strong attraction of full charges keeps ions together • Conduct Electricity when dissolved in water—ions dissolve into water

8. I. Covalent Bonding • Chemical bond formed by the energetically favorable sharing of electrons between 2 atoms • Like charges repel each other with a similar strength that opposite charges attract one another • +/+ proton/proton and -/- electron/electron interaction are unfavorable between any two bonded atoms • The +/- attraction of an ionic bond overcomes this repulsion • In covalent bonds, the atoms are neutral, so some other way of interaction must help overcome the +/+ and -/- repulsion • Sharing the electrons between them can create 2 favorable +/- interaction • The covalent bond will only form if the result is overall lower energy

9. In a covalent bond, the 2 atoms involved must come together until a minimum energy is found. • Far away, there is no favorable interaction • Closer together, the +/+ and -/- repulsions increase energy greatly • If the atoms can share their electrons at lower energy, they form a covalent bond with a specific Bond Length • Bond Energy = amount of energy required to pull them apart again • H2 (covalent) is favorable, H+H- (ionic) is not

10. J. Polar Covalent Bonds • Unequal sharing of the e- between two covalently bonded atoms • This type of bond is somewhere between Ionic and Covalent • A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

11. Electronegativity = attraction for shared electrons • Electronegativity increases from left to right on periodic table • Electronegativy decreases down a group • The most electronegative atom in a bond • has the electrons closest to it most of the time • Atoms with the same electronegativities • form purely covalent bonds • Atoms with different electronegativities • form polar covalent bonds

12. Bond Polarity and Dipole Moment • Molecules with +/- ends have a dipole moment • d tells us it is not a full charge • Arrow points to negative end • Complex molecule shapes: • Dipole moments can cancel out to give a nonpolar molecule

13. K. Partial Ionic Character of Covalent Bonds • Range of Bond Character • Few bonds are purely Ionic or Covalent • Most bonds have some character of both due to unequal electronegativity • Calculating the Percent Ionic Character • To compare molecules, we can calculate how Ionic they are • Ionic Character increases as electronegativity difference increases • No individual bonds are 100% Ionic (for the gas phase) • Everything > 50% Ionic is considered Ionic in the solid phase • Any melted compound that conducts electricity = ionic