Chemical Bonding Fundamentals: Ionic, Covalent, and Electronegativity Concepts

1. Chapter 8 Bonding: General Concepts • Forces that hold groups of atoms together and make them function as a unit. How Often Does The Topic Appear On AP Exam? 12% MC Questions Every Year On FR

2. 8.1 Types of Chemical Bonds • Bond Energy: The energy required to break a bond. • It gives us information about the strength of a bonding interaction. • When a bond forms, the system achieves the lowest possible energy

3. Ionic Bonds • Formed from electrostatic attractions of closely packed, oppositely charged ions. • Formed when an atom that easily loseselectronsreacts with one that has a high electron affinity (Metal + nonmetal) . • System achieves the lowest possible energy by behaving this way.

4. Ionic Bonds Coulomb’s Law – used to calculate the energy of interaction between a pair of ions. • Q1 and Q2 = numerical ion charges • r = distance between ion centers (in nm) • A negative sign indicates an attractive force between opposite charged ions and the ion pair has a lower energy than the separated ions. Ex. Na+ Cl-

5. Bond Length • The distance where the system energy is at a minimum.

6. Interaction of Two Hydrogen Atoms

7. Covalent Bonding • Type of bonding in H2 molecule • Electrons are shared by nuclei • Polar Covalent Bond: Intermediate cases in which atoms are not so different that e- are not transferred like ionic bonds but different enough where e- are not shared equally.

8. Bond Polarity 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. The lowercase  (delta) symbol is used to indicate a fractional charge. The arrow indicates the direction of the polarity.

9. 8.2 Electronegativity • The ability of an atom in a molecule to attract shared electrons (covalent bond) to itself. • = (H  X)actual (H  X)expected • This is Linus Pauling’s equation for determining the relative electronegativities of H and X atoms

10. Electronegativity & Bond Type

11. Relative Bond Polarity • Order the following according to polarity: • H-H, O-H, Cl-H, S-H, F-H • Use electronegativity values in fig 8.3 p. 334 to calculate polarity values 0 0.4 1.9 0.9 1.4 Covalent Bond Polar Covalent Bond 11

12. 8.3 Bond Polarity & Dipole Moments A molecule of HF has a center of positive charge and a center of negative charge and is referred to as having a dipole moment. The arrow always points to the negative end.

13. Electrostatic Potential Diagram 13

14. Bond Polarity and Dipole Moment • For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: • HCl, Cl2, SO3, CH4, H2S

15. 8.4 Ions: Electron Configurations & Sizes • Electron Configuration of Compounds • Two nonmetalsreact: They share electrons to achieve Noble Gas Electron Configurations (NGEC). • A nonmetal and a representative group metalreact to form a binary ionic compound. The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal attracts electrons to achieve NGEC.

16. Predicting Formulas of Ionic Compounds • Rules: • Solid Ionic Compounds are most stable • Noble Gas Electron Configurations are generally the most stable. • Chemical compounds are electrically neutral Using Electron Configurations, you can predict the bond formation: Ca[Ar]4s2 O[He]2s22p4 You can see that Ca will give up its valence electrons to O. Ca+2 + O-2

17. Ion Size • Isoelectronic Ions: Ions containing the the same number of electrons • O2> F > Na+ > Mg2+ > Al3+ • largest smallest • Note that they all have the same electron configuration but the size differs, general for all isolectronic ions the size get smaller as you increase the atomic number because of the increase in number of protons (nuclear attraction).

18. Relative Ion Size • Choose the largest ion in each of the following groups. • Li+, Na+, K+, Rb+, Cs+ • Ba+2, Cs+, I-, Te-2

19. 8.5 Energy Effects in Binary Ionic Lattice Energy is the change in energy when separated gaseous ionsare packed together to form an ionic solid. Lattice energy is negative (exothermic) from the point of view of the system. Negative values indicate energy is given off which is what happens when you form bonds.

20. Formation of an Ionic Solid • 1. Sublimation of the solid metal • M(s)  M(g) [endothermic] • 2. Ionization of the metal atoms • M(g)  M+(g) + e [endothermic] • 3. Dissociation of the nonmetal • 1/2X2(g)  X(g) [endothermic]

21. Formation of an Ionic Solid(continued) • 4. Formation of X ions in the gas phase: • X(g) + e X(g) [exothermic] • 5. Formation of the solid MX • M+(g) + X(g)  MX(s) [quite exothermic]

22. Lattice Energy of CsF

23. Lattice energy can be represented by a modified Coulomb’s Law • k = a proportionality constant that depends on structure of the solid & electron config. • Q1, Q2 = charges on the ions • r = shortest distance between centers of the cations and anions • Lattice energy will have a negative sign because your joining cations and anions.

24. Relationship in Electronegativity Difference & Percent Ionic Character

25. Covalent Bond: A Model • Chemical bonds are forces that cause a group of atoms to behave as a unit. • Bonds result from the tendency of a system to seek its lowest possible energy. • Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

26. Fundamental Properties of Models • Fundamental Properties of Models p350 • 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.

27. 8.8 Covalent Bond Energies & Reactions

28. Covalent Bond Energies • Bond breaking requires energy (endothermic). • Bond formation releases energy (exothermic). • ∆H = enthalpy, a measure of the total energy of a system. Unit is Joules • D = Sum of bond energies(D) per mole, always has a positive sign.

29. H from Bond Energies • Using bond energy values (Table 8.4) to estimate H for the following reaction in the gas phase. • H2(g) + Cl2(g)  2HCl

30. Calc. ∆H from Bond Energies • CH4 + 2Br2 + 2F2 → CF2Br2 + 2HBr + 2HF 30

31. Calc. ∆H from Bond Energies • The total ∆H is -1032 kJ of energy for the reaction. It is an exothermic reaction and this is the amount of heat released in the formation of the the 3 products. Do #53b on p. 384 31

32. 8.9 Localized Electron Model • Assumes a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. • Electron pairs are: • Localized on atom = Lone Pair • Shared between atoms = Bonding Pair

33. Localized Electron Model – Three Parts • 1. Description of valence electron arrangement (Lewis Structures 8.10). • 2. Prediction of geometry (VSEPR Model 8.13). • 3. Description of atomic orbital types used to share electrons or hold lone pairs. (Ch 9)

34. Lewis Structure • Shows how valence electrons are arranged among atoms in a molecule. • Reflects central idea that stability of a compound relates to noble gas electron configuration.

35. Steps for Writing Lewis Structures • Sum the valence electrons from all atoms. Add e- for neg. polyatomic ions Subtract e- for pos. polyatomic ions • Draw skeletal structure using pair e- to form bonds, least electronegative atom is usually central position • Add pairs e- to surrounding atoms • Add remaining e- to central atom

36. Comments About the Octet Rule • 2nd row elements C, N, O, F observe the octet rule. • 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 octetsfirst,then place electrons around elements having available d orbitals.

37. 8.12 Resonance • Occurs when more than one valid Lewis structure can be written for a particular molecule. • These are resonance structures. The actual structure is an average of the resonance structures.

38. Formal Charge • Used to help determine the correct Lewis structure when more than one structure is possible. • The difference between the number of valence electrons (VE) on the free atom and the number assigned to the atom in the molecule. • We need to know: • 1. # VE on free neutral atom • 2. # VE “belonging” to the atom in the • molecule • #VEbelonging = (# Lone Pair) + (1/2 # shared E)

39. Formal Charge • Not as good Better

40. 8.13 VSEPR Model • The structure around a given atom is determinedprincipallyby minimizing electron pair repulsions.

41. Predicting a VSEPR Structure • 1. Draw Lewis structure. • 2. Put pairs as far apart as possible to minimize repulsion. • 3. Determine positions of atoms from the way electron pairs are shared. • 4. Determine the name of molecular structure from positions of the atoms.