2. Chapter 8 - Bonding

3. Ionic Bonding  Electrons are transferred  Electronegativity differences are generally greater than 1.7  The formation of ionic bonds is always exothermic!

4. Determination of Ionic Character Compounds are ionic if they conduct electricity in their molten state Electronegativity difference is not the final determination of ionic character

5. Coulomb’s Law “The energy of interaction between a pair of ions is proportional to the product of their charges, divided by the distance between their centers”

6. Table of Ion Sizes

7. Estimate  H f for Sodium Chloride Na(s) + ½ Cl 2 (g)  NaCl(s) Na(s)  Na(g) + 109 kJ Na(g)  Na + (g) + e - + 495 kJ ½ Cl 2 (g)  Cl(g) + ½(239 kJ) Cl(g) + e -  Cl - (g) - 349 kJ Na + (g) + Cl - (g)  NaCl(s) -786 kJ Na(s) + ½ Cl 2 (g)  NaCl(s) -412 kJ/mol

8. Sodium Chloride Crystal Lattice Ionic compounds form solid crystals at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions. All salts are ionic compounds and form crystals.

9. Properties of Ionic Compounds

10. Covalent Bonding Bonding models for methane, CH 4. Models are NOT reality. Each has its own strengths and limitations.

11. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  Electrons are equally shared  Electronegativity difference of 0 to 0.3

12. Covalent Bonding Forces  Electron – electron repulsive forces  Proton – proton repulsive forces  Electron – proton attractive forces

13. Bond Length and Energy Bonds between elements become shorter and stronger as multiplicity increases.

14. Bond Energy and Enthalpy D D = Bond energy per mole of bonds Energy requiredEnergy released Breaking bonds always requires energy Breaking = endothermic Forming bonds always releases energy Forming = exothermic

15. The Octet Rule sharing Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Monatomic chlorineDiatomic chlorine

16. The Octet Rule and Covalent Compounds  Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level.  Covalent compounds involve atoms of nonmetals only.  The term “molecule” is used exclusively for covalent bonding

17. The Octet Rule: The Diatomic Fluorine Molecule F F 1s 2s 2p seven Each has seven valence electrons FF

18. The Octet Rule: The Diatomic Oxygen Molecule O O 1s 2s 2p six Each has six valence electrons O O

19. The Octet Rule: The Diatomic Nitrogen Molecule N N 1s 2s 2p five Each has five valence electrons N N

20.  Lewis structures show how valence electrons are arranged among atoms in a molecule.  Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron configuration.  Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - ) Lewis Structures

21. Comments About the Octet Rule  2nd row elements C, N, O, F observe the octet rule (HONC rule as well).  2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.  3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.  When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

22. Show how valence electrons are arranged among atoms in a molecule. Reflect the central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

23. The HONC Rule HH Hydrogen (and Halogens) form one covalent bond O Oxygen (and sulfur) form two covalent bonds One double bond, or two single bonds N Nitrogen (and phosphorus) form three covalent bonds One triple bond, or three single bonds, or one double bond and a single bond C Carbon (and silicon) form four covalent bonds. Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

24. C H H H Cl.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make the atom wanting the most bonds central..

25. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons Ethene

26. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons Ethyne

27. Acetic Acid Two electrons (one bond) per hydrogen Eight electrons (four bonds) per carbon Eight electrons (two bonds, two unshared pairs) per oxygen

28. Resonance Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule. The actual structure is an average of the resonance structures. Benzene, C 6 H 6 The bond lengths in the ring are identical, and between those of single and double bonds.

29. Resonance Bond Length and Bond Energy Resonance bonds are shorter and stronger than single bonds. Resonance bonds are longer and weaker than double bonds.

30. Resonance in Ozone, O 3 Neither structure is correct. Oxygen bond lengths are identical, and intermediate to single and double bonds

31. Resonance in a carbonate ion: Resonance in an acetate ion: Resonance in Polyatomic Ions

32. The Localized Electron Model Lewis structures are an application of the “Localized Electron Model” L.E.M. says: Electron pairs can be thought of as “belonging” to pairs of atoms when bonding Resonance points out a weakness in the Localized Electron Model.

33. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical

34. Fundamental Properties of Models  A model does not equal reality.  Models are oversimplifications, and are therefore often wrong.  Models become more complicated as they age.  We must understand the underlying assumptions in a model so that we don’t misuse it.