Concepts of Chemical Bonding

1. Concepts of Chemical Bonding Brown, LeMay Ch 8 AP Chemistry

2. 8.1: Types of “Inter-Atomic” Bonds • Ionic: electrostatic attraction between oppositely charged ions • Covalent: sharing of e- between two atoms (typically between nonmetals) • Metallic: “sea of e-”; bonding e- are relatively free to move throughout the 3D structure IncreasingDiff. of EN

3. Lewis symbols Valence e-: • e- in highest energy level and involved in bonding; all elements within a group on P.T. have same # of valence e- Lewis symbol (or electron-dot symbol): • Shows a dot only for valence e- of an atom or ion. • Place dots at top, bottom, right, and left sides and in pairs only when necessary (Hund’s rule). • Primarily used for representative elements only (Groups 1A – 8A) Ex: Draw the Lewis symbols of C and N. Gilbert N. Lewis(1875 – 1946) • •C • • • : N • •

4. The Octet Rule • Atoms tend to gain, lose, or share e- until they are surrounded by 8 valence e- (have filled s and p subshells) and are thus energetically stable. • Exceptions do occur (and will be discussed later.)

5. [ ]1- [ ]1+ ••: F : •• Na 8.2: Ionic Bonding • Results as atoms lose or gain e- to achieve a noble gas e- configuration; is typically exothermic. • The bonded state is lower in energy (and therefore more stable). • Electrostatic attraction results from the opposite charges. • Occurs when diff. of EN of atoms is > 1.7 (maximum is 3.3: CsF) • Can lead to interesting crystal structures (Ch. 11). • Use brackets when writing Lewis symbols of ions. Ex: Draw the Lewis symbol of sodium fluoride.

6. Lattice Energy • Measurement of the energy of stabilization present in ionic solids DHlattice = energy required to completely separate 1 mole of solid ionic compound into its gaseous ions • Electrostatic attraction (and thus lattice energy) increases as ionic charges increase and as ionic radii decrease. Ex: Which has a greater lattice energy? NaCl or KCl NaCl or MgS

7. Transition metals typically form +1, +2, and +3 ions. • It is observed that transition metal atoms first lose both “s” e-, even though it is a higher energy subshell. • Most lose e- to end up with a filled or a half-filled subshell.

8. 8.4 - 8.5: Covalent Bonding • Atoms share e- to achieve noble gas configuration that is lower in energy (and therefore more stable). • Occurs when diff. of EN of atoms is ≤ 1.7 • Polar covalent: 0.3 < diff. of EN ≤ 1.7 (e- pulled closer to more EN atom) • Nonpolar covalent: 0 ≤diff. of EN ≤ 0.3 (e- shared equally)

9. 8.6: Drawing Lewis Diagrams • Add up valence e- from all atoms in formula. • If there is a charge, add e- (if an anion) or subtract e- (if a cation). • Draw the “molecular skeleton”: • Place the least EN atom(s) in the center. • Array the remaining elements around the center and connect them with a single bond. (When in doubt, put the element written first in the formula in the center of the molecule.) • Complete the octets of the outer (more EN) atoms first. • Place leftover e- on the central atom, even if it violates the octet rule. • If the central atom does not have an octet, create multiple bonds by sharing e- with the outer atoms.

10. Ex: Draw the Lewis structure, and name the molecule. SO42- HCN H2O2 CNS1-

11. 8.8: Exceptions to the Octet Rule • Odd-electron molecules:Ex: NO or NO2 (involved in breaking down ozone in the upper atmosphere) • Incomplete octet: H2 He BeF2 BF3 NH3 + BF3 → NH3BF3 (Lewis acid/base rxn)

12. Expanded octet: occurs in molecules when the central atom is in or beyond the third period, because the empty 3d subshell is used in hybridization (Ch. 9) PCl5 SF6

13. 8.6: Formal Charge • For each atom, the numerical difference between # of valence e- in the isolated atom and # of e- assigned to that atom in the Lewis structure. To calculate formal charge: • Assign unshared e- (usually in pairs) to the atom on which they are found. • Assign one e- from each bonding pair to each atom in the bond. (Split the electrons in a bond.) • Then, subtract the e- assigned from the original number of valence e-. #VALENCE e- in free atom – #NON-BONDING e- – ½(#BONDING e-) FC

14. Used to select most stable (and therefore most likely structure) when more than one structure are reasonable according to “the rules”. • The most stable: • Has FC on all atoms closest to zero • Has all negative FC on most EN atoms. • FC does not represent real charges; it is simply a useful tool for selecting the most stable Lewis structure.

15. Examples: Draw at least 2 Lewis structures for each, then calculate the FC of each atom. SCN1- N2O BF3

16. 8.7: Resonance Structures • Equivalent Lewis structures that describe a molecule with more than one likely arrangement of e- • Notation: use double-headed arrow between all resonance structures. Ex: O3 • Note: one structure is not “better” than the others. In fact, all resonance structures are wrong, because none truly represent the e- structure of the molecule. The “real” e- structure is an “average” of all resonance structures.

17. : : O=O : : Bond Order • An indication of bond strength and bond length • Single bond: 1 pair of e- shared Ex: F2 Longest, weakest •• •• :F-F: •• •• • Double bond: 2 pairs of e- shared Ex: O2 • Triple bond: 3 pairs of e- shared Ex: N2 Shortest, strongest :N ≡ N:

18. Bond Order & Resonance Structures • To determine bond orderwith resonance structures: • Pick one bond and add up the integer bond order in one resonance structure to the same bond position in all other resonance structures. • Divide the sum by the number of resonance structures to find bond order.

19. Examples SO3C6H6

20. 8.9: Bond enthalpy: • DH/mol to break a particular bond of substance (g) Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ? • 1 C-H & 1 Cl-Cl bond are broken (per mole) • 1 C-Cl & 1 H-Cl bond are formed (per mole) Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Note: this is the “opposite” of Hess’ Law where Hrxn = DHproducts – DHreactants

21. Ex: CH4 (g) + Cl2 (g) → CH3Cl (g) + HCl (g) DHrxn = ? BondAve DH/molBondAve DH/mol C-H 413 Cl-Cl 242 H-Cl 431 C-Cl 328 C-C 348 C=C 614 Hrxn ≈  (Hbonds broken) -  (Hbonds formed) Hrxn ≈[(1(413) + 1(242)] – [1(328) + 1(431)] Hrxn ≈-104 kJ/mol Hrxn = -99.8 kJ/mol (actual) Note: 2 C-C ≠ 1 C=C 2(348) = 696 kJ ≠ 614 kJ

22. Ex: CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g) DHrxn=? *CH3(g) + H(g) + 2 Cl(g) Absorb E, break 1 C-H and 1 Cl-Cl bond Release E, form 1 C-Cl and 1 H-Cl bond H CH4(g) + Cl2(g) CH3Cl (g) + HCl(g) DHrxn Hrxn =  (Hbonds broken) +  (- Hbonds formed) Hrxn =  (Hbonds broken) -  (Hbonds formed)

23. 23.5: Metallic bonding • Metallic elements have low I.E.; this means valence e- are held “loosely”. • A metallic bond forms between metal atoms because of the movement of valence e- from atom to atom to atom in a “sea of electrons”. The metal thus consists of cations held together by negatively-charged e- "glue.“ • This results in excellent thermal & electrical conductivity, ductility, and malleability. • A combination of 2 metals is called an alloy.

24. Free e- move rapidly in response to electric fields, thus metals are excellent conductors of electricity. http://www.uwgb.edu/dutchs/EarthSC202Notes/minerals.htm Free e- transmit kinetic energy rapidly, thus metals are excellent conductors of heat. Layers of metal atoms are difficult to pull apart because of the movement of valence e-, so metals are durable. However, individual atoms are held loosely to other atoms, so atoms slip easily past one another, so metals are ductile.