Bonding: General Concepts

1. Bonding: General Concepts Chemical Bonds Electronegativity, Polarity Ionic Bonds Covalent Bonds: Lewis Structures, VSEPR

2. CHEMICAL BONDS • Forces that hold groups of atoms together to form molecules. • The driving force is the lowering of energy due to electrostatic attractions between the positive nuclei and the negative electrons exceeding repulsions between nuclei and between electrons.. • Separated atoms have zero energy and chemically bonded atoms have negative (lower) energy. (Fig 8.1). • The minimum energy or well corresponds to the bond length

3. Figure 8.1 a & b (a) The Interaction of Two Hydrogen Atoms (b) Energy Profile as a Function of the Distance Between the Nuclei of the Hydrogen Atoms

4. CHEMICAL BONDS (2) • This lowering of energy is achieved when atoms achieve a noble gas electron configuration or an octet. • We will see that bonds form in order that each participating atom achieves an octet. • We will also see that there are exceptions.

5. CHEMICAL BONDS (3) • Form between atoms resulting in molecules (covalent bonds, sharing of electrons). • Form between ions resulting in ionic cmps (ionic bonds, electron transfer). • Chemical bonding model assumes molecule consists of individual chemical bonds. • Bond strength varies and is measured by bond energy (kJ/mol) = energy required to break a mole of bonds.

6. ELECTRONEGATIVITY • Defined as the ability of an atom to attract shared electrons in a covalent bond to itself. • EN > 0; Fig 8.3 • EN largest in upper right hand corner of PT. • This unequally sharing leads to unequal charges on the atoms. • Use δ+ and δ- to indicate partial charges on the atoms.

7. Figure 8.3 The Pauling Electronegativity Vaules

8. BOND POLARITY • Polar covalent bond forms when electron pair is not shared equally due to bonded atoms having different EN values. • ΔEN = difference in EN • ~ 0, nonpolar covalent bond. E.g. H2, O2 • < 2, polar covalent bond; e-pair is held more closely by atom with greater EN • > 2, bond is ionic and electron is transferred to form anion and cation (vs Sec 8.6)

9. Figure 8.12 a-c The Three Possible Types of Bonds

10. DIPOLE MOMENT • When there is a separation of electron charge leading to polar bonds, the molecule may have a dipole moment. • All diatomics with polar bonds have a dipole moment. (HCl, NO, CO) • Polyatomics with polar bonds MAY have a dipole moment. (Fig 8.2). H2O, NH3, SO2)

11. Table 8.2 Types of Molecules with Polar Bonds but No Resulting Dipole Moment

12. Figure 8.6 a-c The Structure and Charge Distribution of the Ammonia Molecule

13. IONIC BONDS (8.4) • (Metal) Cation + (Nonmetal) Anion  Ionic Solid held together with ionic bonds. • This solid has a continuous network of cations surrounded by anions and anions surrounded by cations. • The formation of ionic bonds is driven by favorable energy considerations: this is illustrated by the Born-Haber cycle.

14. ATOMIC ION SIZE • Cations shrink and anions expand as electrons are removed or added to the neutral atom. • In an isoelectronic series, the number of electrons stays the same, but Z is constant. • As Z increases, the ion size decreases. • Fig 8.8 • Note that

15. Figure 8.8 Sizes of Ions Related to Positions of the Elements on the Periodic Table

16. Born-Haber Cycle (Fig 8.9, 8.11) • Li(s)  Li(g) Sublimation energy > 0 • Li(g)  Li+(g) + e- IE, T7.6 • ½ F2(g)  F(g) Dissociation energy > 0 • F(g) + e-  F-(g) EA, T7.7 • Li+(g) + F-(g)  LiF(s) Lattice energy • Sum all of these rxns to get energy for • Li(s) + ½ F2(g)  LiF(s) ΔHfo = -617 kJ/mol

17. Figure 8.9 The Energy Changes Involved in the Formation of Lithium Fluoride from Its Elements

18. Lattice Energy, U • KF(s) K+(g) + F-(g) U > 0 • Electrostatic attraction between Cation and Anion. • As charge increases, U increases.

19. COVALENT BONDS (8.7) • Most common type of chemical bond. • Involve electrons shared by two nuclei. • The covalent bond model assumes that a molecule is an arrangement of individual bonds that form between 2 atoms because the molecule is energetically favored (i.e. energy is at a minimum) compared to the separated atoms.

20. DISSOCIATION BOND ENERGY • Chemical bonds can be assigned average (±10%) dissociation bond energies (T8.4) and bond lengths (T8.5) • D > 0 kJ/mol; measure of bond strength. • AB(g)  A(g) + B(g) • Note single vs double vs triple bonds D values. • ΔHrxn ≈ Σ D(bonds in R) – ΣD (bonds in P) because bond breaking is endothermic and bond formation is exothermic.

21. Table 8.4 Average Bond Energies (kj/mol)

22. Table 8.5 Bond Lengths for Selected Bonds

23. COVALENT BONDS (2) • Determine physical and chemical properties of cmps. • Determine the likelihood and products of chemical reactions. • Determine molecular shape (Sec 8.13).

24. LOCALIZED ELECTRON (LE) BONDING MODEL • Valence electrons participate in the formation of chemical bonds. • Electron pairs are localized between (shared or bonding pair) or on (lone pair) atoms such that each atom has an octet or duet of electrons. • VSEPR model predicts molecular geometry based on LE bonding model.

25. LEWIS SYMBOLS and STRUCTURES • Lewis symbol: picture of molecule showing arrangement of its valence electrons around atoms. • Lewis structure: picture of molecule showing bonding electrons as lines and nonbonding electrons as dots or lines. • Especially used for main group elements (p 357)

26. COVALENT BONDS (3) • Form when electron pairs are shared so that each atom achieves an octet (duet). • Coordinate covalent bond forms when one atom provides both bonding electrons. • Multiple covalent bond forms when more than one electron pair is shared between two atoms (double bond, bond order 2 [CO2] and triple bond, bond order 3 [N2]).

27. WRITING LEWIS STRUCTURES • Determine total # of valence electrons. • Write skeletal structure with central atom [lowest EN]; terminal atoms [H, higher EN] • Use electron pairs to form bonds. • Achieve octet rule for terminal atoms • Add the remaining to the central atom. • Form multiple bonds if needed.

28. WRITING LEWIS STRUCTURES (2) • Exceptions to octet rule [odd # of valence electrons (NO), free radicals, incomplete octets (B), more than 8 electrons (expanded valence shell SF6)]. • Resonance structures showing different but equivalent distributions of electrons; note delocalization (vs localization) of electrons. • Be guided by experimental observations.

29. FORMAL CHARGE (FC) • FC = [VE in free atom] - [VE asigned in molecule] • FC is a hypothetical charge for electron loss (+) or gain (-) due to bond formation. • [VE]free = # valence e’s for Group A atoms • [VE] assigned = all lone pair electrons on atom + 1/2 shared electrons

30. FORMAL CHARGE (2) • Best Lewis structure has minimum FC (zero). • Formal Charge method is not perfect and can lead to incorrect “best” Lewis structures. • The best Lewis structure is consistent with exptal evidence (bond lengths, EN data, etc)

31. VSEPR MODEL • VALENCE-SHELL ELECTRON-PAIR REPULSION (VSEPR) Method helps us determine molecular geometry. • Molecular geometry: 3-D shape of the molecule. • This method assumes that the final positions of nuclei are the ones that minimizes electron repulsions because this is the one associated with the lowest energy.

32. VSEPR METHOD (2) • Determine Lewis structure of molecule. • Count electron “pairs” around the central atom where a “pair” may be a single e, lone pair, single bond, double bond, triple bond. • Determine geometry of electron pairs. • Determine molecular group geometry with A = central atom; X = terminal atom; E = lone pair of electrons. T8.6, 8.7, 8.8

33. Table 8.6 Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion

34. MOLECULAR GEOMETRY

35. MOLECULAR GEOMETRY (2) • Electron pair geometry differs from molecular geometry when there are lone electron pairs (E). • Electron-electron repulsions decrease as E-A-E> E-A-X> X-A-X; X = bonded pair • Resonance structures • Note bond angles