1. Chapter 9: Chemical Bonds Types of Bonds Ionic –Metal and nonmetal –Electron transfer –Infinite lattice Covalent –Nonmetal and nonmetal –Shared electrons –Individual molecules Metallic –Metal and metal –“electron sea” –Infinite lattice

2. Ionic Bonds: Octet Rule In forming ionic compounds, atoms tend to gain or lose electrons in order to achieve a stable valence shell electron configuration of 8 electrons. Group I metals --> +1 cations (Li +, Na +, etc) Group II metals --> +2 cations (Mg 2+, Ca 2+, etc) Al (group III) --> Al 3+ Group VII (17) --> -1 anions (F –, Cl –, Br –, etc) Group VI (16) --> -2 anions (O 2–, S 2–, etc) Group V (15) --> -3 anions (N 3–, P 3– ) e.g. Na 2s 2 2p 6 3s 1 --> Na + 2s 2 2p 6 {~Ne} Cl 3s 2 3p 5 --> Cl – 3s 2 3p 6 {~Ar} Remember Ch. 2 Figure 2.14 (p62)

3. Ionic Bonds: Lewis Symbols simple notation for showing number of valence electrons ClO group VII (7 valence e – )group VI (6 valence e – ) 3s 2 3p 5 2s 2 2p 4 O Cl e.g. use Lewis symbols to illustrate the formation of a compound of sodium and sulfur Na 2 S Na + combined with S 2–

4. Lattice Energy Ionic Bonding; electrostatic attraction of positive (cation) and negative (anion) ions Neutral atoms cation + anion Ionic compound Lattice energyenergy released when gaseous ions combine to form crystalline solid (an ionic compound) e.g. lattice energy of NaCl is 787 kJ/mol: Na + (g) + Cl – (g) --> NaCl (s)  H = –787 kJ Bigger ions  smaller lattice energy Higher charge  larger lattice energy (IE + EA) e – transfer (lattice energy)

5. Born-Haber Cycles Lattice energy can be calculated using a Born-Haber cycle; a hypothetical series of steps describing the formation of an ionic compound from the elements. Here,  H° lattice =  H° f – (  H° step1 +  H° step2 +  H° step3 +  H° step4 )

6. Covalent Bonding 1. Covalent Bond Formation –results from sharing of one or more pairs of electrons between two atoms –Examples: 2. Octet Rule -- for covalent bonding –In forming covalent bonds, atoms tend to share sufficient electrons so as to achieve a stable outer shell of 8 electrons around both atoms in the bond.

7. Examples

8. Multiple Bonds “double” and “triple” bonds double bondsharing of 2 pairs of electrons between two atoms triple bondsharing of 3 pairs of electrons between two atoms Type of Bondsingle doubletriple Bond order 1 2 3 Examples: O 2 {O=O double bond} N 2 {N ≡ N triple bond} CO 2 {two C=O double bonds} bond energy/bond strength bond distance

9. Electronegativity and Bond Polarity electronegativitytendency of an atom in a molecule to attract electrons to itself e.g. Cl is more electronegative than H, so there is partial charge separation in the H-Cl bond: The H-Cl bond is described as “polar” and is said to have a “dipole” The entire HCl molecule is also polar as a result More complex molecules can be polar or nopolar, depending on their 3-D shape (Later) Electronegativity Increases Periodic Table

10. Lewis Electron Dot Structures General Procedure -- stepwise process --Write the skeletal structure (which atoms are bonded?) --H atoms terminal --More electronegative atoms terminal (e.g. halogens) --Count all valence electrons (in pairs) and charges --Place 2 electrons (1 pair) in each bond --Complete the octets of the terminal atoms --Use multiple bonds if needed to complete the octet of the central atom, or --Put any remaining electron pairs on the central atom --Show formal charges Apply the OCTET RULE as follows: --H never has more than 2 electrons (I.e. one bond) --2nd row elements (e.g. C, N, O) almost always have an octet and never have more than 8 electrons (sometimes boron has only 6) --3rd row and higher elements can have more than 8 electrons but only after the octets of any 2nd row elements are completed

11. Formal Charge The “apparent” charge on an atom in a covalent bond = (# of valence e – in the isolated atom) - (# of bonds to the atom) - (# of unshared electrons on the atom) minimize formal charges whenever possible (but the octet rule takes priority!) NH 4 + CO All nonzero formal charges must be shown in a Lewis structure. Put a circle around the formal charge.

12. Lewis Electron Dot Structures General Procedure -- stepwise process --Write the skeletal structure (which atoms are bonded?) --H atoms terminal --More electronegative atoms terminal (e.g. halogens) --Count all valence electrons (in pairs) and charges --Place 2 electrons (1 pair) in each bond --Complete the octets of the terminal atoms --Use multiple bonds if needed to complete the octet of the central atom, or --Put any remaining electron pairs on the central atom --Show formal chargesand resonance forms as needed Apply the OCTET RULE as follows: --H never has more than 2 electrons (I.e. one bond) --2nd row elements (e.g. C, N, O) almost always have an octet and never have more than 8 electrons (sometimes boron has only 6) --3rd row and higher elements can have more than 8 electrons but only after the octets of any 2nd row elements are completed

13. Resonance When multiple bonds are present, a single Lewis structure may not adequately describe the compound or ion -- occurs whenever there is a “choice” of where to put a multiple bond. e.g. the HCO 2 – ion is a “resonance hybrid” of two “contributing resonance structures” The C-O bond order is about 1.5 (average of single and double bonds) All reasonable resonance structures must be shown in a Lewis structure.

14. Lewis Dot Structure Checklist Correct total # of valence electrons Correct connectivity NO atoms with > octet (or duet) in group 2 or below All atoms have octets (or duets), if possible Lone electrons clearly shown Nonzero formal charges included (and circled) –Note it should be “2+” not “+2” –If the charge is 1+ or 1- do not write the “1”! Important resonance structures included Ions have brackets and overall charge (not circled)

15. Exceptions to the Octet Rule Odd-Electron Species –“radicals” or “free radicals” Incomplete Octets –Group 3 elements often have 6 electrons total, e.g. AlCl 3 –Often form bonds to complete the octets Expanded octets –Only 3 rd row or higher

16. Sample Problems Write Lewis Electron Dot Structures (including formal charges and/or resonance as needed) for the following compounds and ions. PF 3 HCNSF 5 – NO 2 – SOCl 2 O 3 HNO 3 H 2 CON 3 –

17. Sample Problems Write Lewis Electron Dot Structures (including formal charges and/or resonance as needed) for the following compounds and ions. PF 3 HCNSF 5 – NO 2 – SOCl 2 O 3 HNO 3 H 2 CON 3 –

18. Bond Energies and Estimating  H Bond Energy—energy required to break 1 mole of the bond in the gas phase, e.g. HCl (g)  H (g) + Cl (g)  H = 431 kJ/mol CH 4(g)  CH 3(g) + H (g)  H = 438 kJ/mol Endothermic! Estimating  H rxn using bond energies;  H rxn = ∑(  H bonds broken) – ∑(  H bonds formed)

19. Sample Problem Ethanol is a possible fuel. Use average bond energies to calculate  H rxn for the combustion of gaseous ethanol.

20. Sample Problem Ethanol is a possible fuel. Use average bond energies to calculate  H rxn for the combustion of gaseous ethanol. CH 3 CH 2 OH (g) + 3 O 2(g)  2 CO 2(g) + 3 H 2 O (g)  H rxn = -1245 kJ/mol