Chem 110 McGill SOS

1. Chem 110McGill SOS • Midterm 2

2. Topics on Midterm II • Chapter 10+11+12 • Remember, only 15% of the questions are going to be from midterm 1! Don’t overdo it!

3. Chapter 10 – Bonding I • Covalent bonding • Polarity & dipole moment • Lewis structures • Formalcharge • Resonance • VSEPR theory • Bond Energies

4. 3 types of bonding • Covalent = sharing of electrons • Ionic = transfer of electrons • Metallic = delocalization of electrons

5. Covalent bonding • pure covalent = same element, equal sharing • polar covalent = different elements (non-metals), unequal sharing

6. Chapter 10 – Bonding I • Covalentbonding • Polarity & dipole moment • Lewis structures • Formalcharge • Resonance • VSEPR theory • Bond Energies

7. Polar covalent = think Dipole moment • Arrow always points from the partial positive charge (δ+) to partial negative charge(δ-) • electronegativity differences determine the degree of polarity

8. Electronegativity • a measure of the electron-attracting power of a bonded atom • type of bond determined by difference in EN • remember, 1.7 is arbitrary! • Ionic -greater than 1.7 • Polar Covalent - between 0.5-1.7 • Pure Covalent - less than 0.5

9. Chapter 10 – Bonding I • Covalentbonding • Polarity & dipole moment • Lewis structures • Formalcharge • Resonance • VSEPR theory • Bond Energies

10. Lewis theory • Valence electrons are most important in chemical bonding • Why do atoms bond? To acquire a stable octet. • Remember, eight is yummy!

11. Lewis theory -terminology • Octet Rule: Each atom must have 8 valence shell electrons in a lewis structure • Bonding pair vs. lone pair • Single, double and triple covalent bonds

12. Lewis Symbols • Chemical symbol in the middle • Dots represent valence electrons

13. Lewis Structures - Ionic Compounds

14. Lewis Structures - Covalent Compounds • Strategy • 1. Count total # valence electrons • Add/subtract additional charges • +ve charge, remove e- • -ve charge, add e- • 2. Identify the central and the terminal atoms (any carbon or hydrogen?) • 3. Join atoms through SINGLE covalent bonds. This is your skeletal structure • 4. Achieve octet around each atom • For remaining valence e-, add to central atom. Add double and triple bonds when necessary to complete an octet.

15. Lewis Structures - Tips • Hydrogen is always a terminal atom • Carbon is always a central atom • Central atom (if not carbon) will be the least electronegative element • Go for a compact, symmetrical structure

16. Example: HCN • Step 1: Count total # valence e- • H=1, C=4, N=5 • So…1+4+5=10 e-

17. Example: HCN • Step 2: Identify the central and the terminal atom (any carbon or hydrogen?) • Carbon=central atom • Hydrogen= terminal atom • therefore Nitrogen= also terminal atom

18. Example: HCN • Step 3: Join atoms through SINGLE covalent bonds. This is your skeletal structure.

19. Example: HCN • Step 4: Complete the octet of terminal atoms. • Count the number of valence electrons remaining. If you have any remaining, use them on the central atom. If your central atom does not have a complete octet, start switching single covalent bonds to double (or triple if necessary)

20. Chapter 10 – Bonding I • Covalentbonding • Polarity & dipole moment • Lewis structures • Formal charge • Resonance • VSEPR theory • Bond Energies

21. Formal charges • helps us draw Lewis structures • tells us where the e- are located • F.C. = # valence e- - # lone pair e- - 1/2 # bonding e- • NOTE: sum of all F.C in molecule must be equal to overall charge!

22. Formal Charge - Question Types • 2 types of questions: • 1. what is the formal charge on each atom? • Use FC equation

23. Example: CHO2(-) • Draw structure out, if necessary • Use F.C. = # valence e- - # lone pair e- - 1/2 # bonding e- • F.C. (H) =1-0-0.5(2)=1-0-1=0 • F.C. (C) =4-0-0.5(8)=4-0-4=0 • F.C. (O)-SB =6-6-0.5(2)=6-6-1=-1 • F.C. (O)-DB =6-4-0.5(4)=6-4-2=0 • overall F.C. = -1 = charge on molecule

24. Formal Charge - Question types • II. Use formal charges to see if the lewis structure is plausible or not? • Rules: • Sum of the formal charges must equal 0 • exceptions: the molecule is an ion (like in the last example) • Keep structures as small as possible • Negative formal charges are usually on the most electronegative atoms • Positive formal charges are usually on the least electronegative atoms

25. Example: NH2CN • NH2CN has two possible Lewis structures • How do we know which one is more plausible? • Use formal charges!

26. Chapter 10 – Bonding I • Covalentbonding • Polarity & dipole moment • Lewis structures • Formalcharge • Resonance • VSEPR theory • Bond Energies

27. Resonance forms • Two or more plausible lewis structures can be written for the same molecule. • Note: They will have the same skeletal structure. It’s only the electron distribution that changes! • Each lewis structure contributes to the resonance hybrid. • increasing resonance = more stable

29. Exceptions to the Octet Rule • Odd-electron species • total # of valence electrons is odd • use formal charges to determine where to place the single electron • Incomplete octets • in order to avoid breaking a formal charge rule (i.e. negative formal charges are usually on the most electronegative atoms), you break the octet rule • usually only seen with compounds containing Be, B, Al • Expanded valence shells • seen with elements containing d-obitals like S, P + bonded to a highly electronegative atom like O,F,Cl

30. Chapter 10 – Bonding I • Covalentbonding • Polarity & dipole moment • Lewis structures • Formalcharge • Resonance • VSEPR theory • Bond Energies

31. VSEPR Theory • concerns shapes of molecules • Electron pairs (either lone pair or bond pairs) assume an orientation about the atom that will minimize repulsions • therefore we want to maximize the distance between electron pairs • order of repulsion interactions: • Lone pair-lone pair = most repulsive • lone pair-bond pair • bond pair-bond pair = least repulsive

32. VSEPR Theory - Tips • Molecular shape depends on: • number of electron pairs • type of electron pair (lone pair vs bonding pair) • When determining themolecular geometry, you must first determine the electron group geometry! • Treat multiple bonds (double, triple) as if it were a single bond

33. KNOW THIS: Molecular Geometries • 2 electron groups • 0 lone pairs=linear • 3 electron groups • 0 lone pairs= trigonal-planar • 1 lone pair= bent • 4 electron groups • 0 lone pairs= tetrahedral • 1 lone pair= trigonal-pyramidal • 2 lone pairs=bent • 5 electron groups • 0 lone pairs=trigonal-bipyramidal • 1 lone pair=seesaw • 2 lone pairs=T-shaped • 3 lone pairs=linear • 6 electron groups: • 0 lone pair=octahedral • 1 lone pair=square-pyramidal • 2 lone pair=square-planar

34. VSEPR Theory • Step 1: Count the # of electron groups (the number of bonded atoms+the number of lone pairs) • Step 2: Determine electron group geometry • 2 electron groups: linear • 3 electron groups: trigonal-planar • 4 electron groups: tetrahedral • 5 electron groups: trigonal-bipyramidal • 6 electron groups: octahedral

35. VSEPR Theory • Step 3: Count the number of lone pairs. If there are no lone pairs, the electron group geometry IS the molecular geometry. • Step 4: Determine molecular group geometry. • each additional lone pair changes the molecular geometry

36. Example: Sulfur Tetrafluride • 5 electron groups • Electron geometry is trigonal bipyramidal • However, because 1 of our electron groups is a lone pair, the molecular geometry will be seesaw

37. Chapter 10 – Bonding I • Covalentbonding • Polarity & dipole moment • Lewis structures • Formalcharge • Resonance • VSEPR theory • Bond Energies

38. Bond engergies, order, and bond length • Bond energy is the energy required to break the bond between atoms • single covalent bond= weakest • Triple covalent bond = strongest • Bond order • Single covalent bond=1 • Double covalent bond=2 • Triple covalent bond=3 • Bond length is the distance between the centers of the two atoms joined by a covalent bond • NOTE: As bond order increases, bond energy increases and bond length decreases.

39. Chapter 11- Bonding II • Valence Bond Theory • Hybridization • M.O. Theory • Metallic Bonding

40. Valence Bond Theory • In a molecule, electron probability is highest where the orbitals of the atoms overlap to form covalent bonds • overlap allows electrons greater freedom across both orbitals • thus greater orbital overlap = more stable (think about single vs double vs triple bonds)

41. Valence Bond Theory - types of hybrid bonds • σ bond: most of e- density around bond axis • formed by overlap of s-s, s-p, or p-p • head-to-head overlap • π bond: e- density above and below the bond axis • formed by overlap of p-p • side-to-side overlap

43. Valence Bond Theory - multiple bonds • Single covalent bond= 1 sigma bond • Double covalent bond=1 sigma bond (σ)+ 1 pi bond (π) • Triple covalent bond=1 sigma bond (σ) + 2 pi bonds (π) • NOTE: add π bonds only after σ bonds! π bonds only for double or triple bonds

44. Practice Question The resonance structure of PO4(3-) where the central atom carries a formal charge of zero is surrounded by: • A. 4 sigma bonds • B. 4 sigma and 1 pi bonds • C. 4 sigma and 2 pi bonds • D. 4 sigma and 3 pi bonds • E. 4 sigma and 4 pi bonds

45. Practice question - steps • draw Lewis structure • remember to complete octets and use F.C. to determine best structure • count single and multiple bonds • single bond = 1 sigma • multiple bonds = 1 sigma + additional pi

46. Chapter 11- Bonding II • Valence BondTheory • Hybridization • M.O. Theory • Metallic Bonding

47. Hybridization • Bonded atoms do not have the same orbitals as non-bonded, isolated atoms. They have hybrid orbitals. • VSEPR -> geometry • Valence bond theory -> bonding • Number of hybrid orbitals= Number of electron groups • NOTE: lone pairs are also hybridized! Double and triple bonds are not!

48. Hybrid Orbitals

49. Hybridization • Question: What molecular shapes can a sp^3d^2 hybridized molecule have?

50. Hybridization • Answer: Octahedral, Square Pyramidal and Square Planar