1. Chapter 8 General Bonding Concepts

2. 8.1: I. Types of Chemical Bonds A. Determines behavior/properties of compounds -ex. Carbon can form graphite or diamonds based on configuration of bonds B. Bond dissociation energy: energy required to break a bond, based on length of bond

3. II. Ionic Bonding A. Typically happen between a metal and non-metal B. Metals lose electrons easily and non-metals have a high exothermic affinity for electrons C. Closely packed ions have strong electrostatic attractions for each other D. Bonds exist because this is the lowest energy state for both ions (most stable)

4. III. Coulomb’s Law A. Attractive energy between a pair of ions E = 2.31x10 -19 Jnm (Q 1 Q 2 /r) B. r = distance between ion centers in nm C. Q 1, Q 2 are ion charges ***Know 10 Ǻ (angstroms) = 1 nanometer***

5. IV. Coulomb Calculation A. Na +, Cl - distance is 2.76 Å (0.276 nm) -ex. E = 2.31x10 -19 Jnm ((+1)(-1)/0.276nm) E = - 8.37x10 -19 J B. A negative bond energy means that energy of compound is lower (more stable) than energy of separated ions C. Can also be used to calculate repulsive energy of same charged ions (positive value)

6. V. Why Do Atoms Bond? A. Ex. in H 2 you have 3 forces occurring: 1.Proton – Proton repulsion 2.Electron – Electron repulsion 3.Proton – Electron attraction B. Need to determine where atoms need to be to have minimum energy when combining these forces C. Bond length: distance of minimum energy

7. VI. For H 2 A. At infinite distances, combination of repulsive and attractive forces is zero B. At very close range, proton repulsion is too high leading to positive energy (repulsion) C. At ~ 0.74 Å there is the lowest negative (attractive) energy which is the most stable D. E - occupy most of their time between protons E. Shared e - form a “Covalent Bond”

8. VII. Polar Covalent Bond A. In between electrons being transferred and shared equally, they can also be shared unequally B. In polar covalent bonds, the atom that shares more of the electrons gets a slight negative charge (  - ) and the other atom gets a slight positive charge (  + )

9. 8.2: I. Electronegativity A. Ability of an atom in a molecule to attract shared electrons to itself B. When atoms have same or nearly same electronegativity = covalent bond C. When there is a large difference in electronegativity between atoms = ionic bond D. When the electronegativity difference is in between a covalent or ionic = polar covalent bond

10. II. Electronegativity Trend A. Increases left to right and bottom to top

11. 8.3: I. Bond Polarity and Dipoles A. Dipolar: having an area of slight positive charge and an area of slight negative charge B. Shown by arrow pointing at negative side C. Depending on the way atoms are arranged in a molecule, the entire molecule could have a dipole called a “molecular dipole” CH 2 – F 2  +  -

12. 8.4: I. Ions: E - Configurations and Sizes A. To achieve the most stability, most atoms form attachments to achieve Noble gas configurations (full valence orbitals) B. Non-metallic elements share e - with other non-metals or take e - away from metals to have full valence

13. II. Ions A. We consider ionic compounds as typically solid because in an aqueous form ions are mostly separated and as a gas ions are usually very far apart due to stability B. As solids, ions form complex crystal structures which have all cations and anions arranged so that repulsive forces minimized and attractive forces maximized

14. III. Forming Ions A. When becoming an ion, an atom seeks the most stable form, by giving away or taking electrons it can achieve a Noble gas configuration B. When that ion comes in contact with an ion of opposite charge they electrostatically attract and form a neutral compound C. Since members of the same group on the periodic table have the same valence electrons, each group has a recognized charge it will typically take

15. IV. Sizes of Ions A. Cations smaller than parent atoms due to loss of e - (collapsing of outer orbital) and stronger nuclear attraction of remaining e - B. Anions larger than parent atoms because more e -, less nuclear attraction per e - (smaller Z eff ) C. Ion size increases down group like atomic radius because of more shells

16. V. Isoelectronic Ions A. Contain same number of electrons as each other B. Isoelectronic ions become smaller as you go left to right because of increasing nuclear charge pulling on same number of electrons

17. 8.5: I. Forming Binary Ionic Compounds A. Lattice energy: the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid B. Lattice energy is negative since it is exothermic

18. 8.6: I. Covalent Bonds A. All covalent bonds between two different atoms involve some part of an ionic character B. Ionic character increases as electronegativity differences increase

19. II. How do we define Ionic Compounds? A. Polyatomic ions are held together with covalent bonds so they are not completely ionic B. We define an ionic compound as one that can be dissociated in an aqueous solution and conducts electricity C. Ionic compounds are generally called “salts”

20. 8.7: I. A Model Of Covalent Bonds A. Bonding Summary 1.Bonds are forces causing atoms to behave as a unit 2.Bonds result from tendency of a system to seek its lowest possible energy B. Chemical bonds: we take the overall energy of stabilization of a molecule and divide it by the parts to determine the energy of each part C. Energy of stabilization is amount of energy to break apart a molecule

21. 8.8: I. Covalent Bond Energies A. Bond energy is average amount of energy to break a bond, similar to stabilization energy B. Single bond: when one pair of e - shared C. Double, triple bond: when 2 or 3 pairs of e - shared D. As # of bonds increases, the bond energy increases and bond length decreases

22. II. Bond Energy and Enthalpy A. Bond energies can be used to determine enthalpy (∆H), heat change at constant pressure B. ∆H = Sum of bond energies (broken, reactants) – Sum of bond energies (formed, products)

23. 8.9: I. Localized Electron Bonding Model A. Assumes molecule has atoms bound by sharing pairs of e - using atomic orbitals of all atoms B. E - on particular atom or between atoms C. Lone pair: e - pairs around atom D. Bonding pairs: e - pairs between bonded atoms

24. II. L.E. Model Requirements 1.Lewis dot structures show valence e - arrangement (8.10) 2.Predict molecular geometry w/ VSEPR theory (8.13) 3.Describe atomic/hybrid/molecular orbital types used by atoms to share e - or hold lone pairs (Ch.9)

25. 8.10: I. Lewis Dot Structures A. Shows how valence e - are arranged in molecule B. Based on assumption that most stable form of atom is Noble gas configuration C. Ionic Ex. Na-Cl D. Covalent Ex. H 2

26. II. Duet, Octet Rule A. Duet rule: Hydrogen needs two e - to have a Noble gas configuration (He) B. Octet rule: filling s and p valence orbitals (holds 8 e - )

27. III. Lewis Dot Rules 1.Add all valence e - from atoms involved together 2.Use pair of e - per bond, start with single bonds 3.Arrange remaining e - around atoms to satisfy duet and octet rules 4.If molecule not stable, try double or triple bonds

28. 8.11: I. Exceptions to Octet Rule A. Some atoms tend to have fewer than an octet B. Ex. Boron often only gets 6 e - around it as in BF 3 C. Some atoms can exceed octet rule like Sulfur

29. II. Exception Comments A. 2 nd row elements will follow octet rule except for Be (4), and B (6, 8) which sometimes have less B. 2 nd row elements cannot exceed octet rule C. 3 rd row and higher can exceed octet (like Phosphorous (8, 10), Sulfur (8,10,12) due to presence of “D” orbitals which can hold extra e - D. When writing Lewis dot structures, satisfy octet rule for atoms first, if e - left over place on elements which have available D orbitals

30. 8.12: I. Resonance A. When more than one Lewis dot structure is possible for a given molecule B. E - can “resonate” between these multiple states, Ex. Nitrate (NO 3 - ), 24 valence e - C. The correct Lewis dot structure consists of these three structures happening simultaneously

31. II. Delocalized Electrons A. Unlike what is stated in the localization model of e -, e - are not in set locations B. E - move around constantly, so they can provide equivalent bonding to molecules with resonance C. Nitrate doesn’t have one double bond with two single bonds, it has three partly double bonds D. We still use “Localized E - Model” because it is convenient for Lewis dot structures

32. III. Multiple “Stable” States A. When molecules have extra # of electrons there are multiple ways to assign the extra electrons B. To determine who gets them we assign formal charge to atoms (***not a real charge***) C. Formal charge = (valence e - on free atom) – (e - on atom in molecule) D. To assign e - on atom in molecule assume: lone pair e - belong entirely to atom, shared e - divided equally

33. IV. Formal Charge Examples A. P in phosphate (PO 4 3- ) can have 8 or 10 e - so it can have two possible structures B. We need to calculate formal charge on both to determine which one more likely C. In first each “O” has 6 atomic valence – 7 molecular valence = formal charge of -1, Phosphorous has 5 atomic – 4 molecular valence = +1 D. In second single bonded “O” has -1 charge, double bonded has 0 charge, Phosphorous has 0 charge

34. V. What Does Formal Charge Mean? A. We go with the structure that has formal charges that make more sense 1.Sum of formal charges on all atoms must equal overall charge of the molecule 2.Formal charges of zero or with negatives on more electronegative atoms are more favored B. Because of this we go with the second structure which has 4 resonance structures

35. 8.13: I. VSEPR Model A. We can determine 3-D structure of molecules by making an arrangement that minimizes e - pair repulsions B. Valence Shell Electron Pair Repulsion: e - pairs whether bonded or lone pairs separate to minimize repulsive forces

36. C. BeCl 2 takes a linear structure because it only has two bonds and no lone electron pairs on central atom Linear D. BF 3 takes a flat structure with three bonds and no lone pairs on central atom Trigonal Planar E. CH 4 takes a 3-D structure with 4 bonds and no lone pairs on central atom Tetrahedral

37. G. H 2 O has a flat shape like linear but since there are two lone pairs on central atom, the shape bends more like a tetrahedral or pyramidal shape F. NH 3 has a 3-D shape similar to tetrahedral with 3 bonds and one lone pair on central atom Pyramidal Bent

38. H. PH 5 has five bonded atoms with no lone pair e - on central atoms Trigonal Bipyramidal I. SH 6 has 6 bonded atoms with no lone pair e - on central atom Octahedral

39. II. Other VSEPR Structures A. Depending on lone pairs found on central atoms, there can be other 3-D or planar structures B. Some molecules have multiple possible structures, largest separation of e - pairs is favored structure C. Multiple bonds count as one effective e - pair, Ex. CO 2 containing two double bonds between C and O is linear

40. III. Molecules Without Central Atom A. When there is no middle atom, you determine VSEPR structures for any atoms with surrounding atoms and then combine multiple structures