Unit 7 – Bonding & Molecular Geometry

1. Unit 7 – Bonding & Molecular Geometry

2. Definitions Chemical Bonds • Force that holds atoms together • It’s all about the electrons (e-) • Electrons available for bonding are called valence electrons!

3. Types of Chemical Bonds Ionic Bond • Bond between metal and nonmetal due to “electrostatic interactions” • Attraction between positively and negatively charged ions (cations and anions) • Electrons are completely transferred from metal to nonmetal

4. Ionic bonds Result from a Transfer of Valence Electrons + -

5. Types of Chemical Bonds Covalent Bond • Bonds in which e- are shared • Most common type

6. Shared Electrons Complete Shells F F

7. Hydrogen Molecule—Energy Diagram Note: at .074 nm, attractive forces are balanced with repulsive forces!

8. Definitions • Octet rule (Rule of 8) • Most atoms want 8 e- in the outer shell very stable • H2 and He want a “duet” (2 e-) • Electron configuration for duet = ns2 • Electron configuration for octet = ns2 np6

9. Two other definitions you need to know: • Bonding pairs are electrons involved in bonding. • Lone pairs are electrons NOT involved in bonding. • They are only located on one atom. (a.k.a non-bonding pair)

10. Lewis Dot Diagrams . . . C . Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon. • A Lewis dot diagram depicts an atom as its symbol and its valence electrons. • Ex: Carbon

11. Drawing Lewis Dot Diagrams . . . . . Cl . . • Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons. • Same group # = Same Lewis Dot structure • Ex. F, Cl, Br, I, At • Example: Chlorine (7 valence electrons b/c it is in group 17)

12. Note: In the final structure, the placement of the dots around the element is not crucial: Maximum # of valence electrons = 8

13. Paired and Unpaired Electrons • As we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired. • When it comes to bonding, atoms tend to pair up unpaired electrons. • A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond. • A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

14. Bonding in Ionic Compounds The ionic bond forms from attraction of cations for anions.

15. Review of Ionic Charge and Isoelectronic Ions Isoelectronic: having same # of e- (same e- configuration) • Na  Na+ + e- • Cl + e-  Cl1- What elements are Na+ and Cl-isoelectronic with?

16. Structure of Ionic Compounds Ionic compounds have formula units—these show ratio of ions in the crystal lattice.

17. Writing Lewis Dots Structures for Ions • Uses either 0 or 8 dots, brackets and a superscript charge designate to ionic charge • Ex.) Li+, Be+2, B+3, C+4, N-3, O-2, F-1

18. Writing Lewis Dots Structures(Ionic Compounds) Lewis Dot Diagrams of Ionic Compounds • Ex. 1) NaCl • Ex. 2) Li2O

19. Lewis Representations of Ionic Structures NaCl MgO Li2O

20. Covalent Compounds and Lewis Dot Diagrams • Lewis structures for covalent molecules show sharing of e- : H:H OR H-H • Bonding pair e-(shared e-)are counted as belonging to both atoms. (each atom has octet) • Bonding pair can also be shown as a dash between atoms.

21. Drawing Electron Dot Diagrams for Molecules • Chemists usually denote a shared pair of electrons as a straight line. F F • Sometimes the nonbonding pair of electrons are left off of the electron dot diagram for a molecule

22. Examples H CH4 H C H H H N H NH3 H

23. Types of Covalent Bonds • Single Bond • 2 e- are shared in a bond (1 from each atom) • Double Bond • 2 pairs of e- are shared (4 e- total, 2 from each atom) • Triple Bond • 3 pairs of e- are shared (6 e- total, 3 from each atom)

24. Rules for Drawing Lewis Dot Diagrams • Add up the total number of valence e- for each atom in the molecule. • Each (-) sign counts as 1 e-, each (+) sign subtracts one e- • Write the symbol for the central atom then use one pair of e- to form bonds between the central atom and the remaining atoms. • Count the number of e- remaining and distribute according to octet rule (or the “duet” rule for hydrogen) • If there are not enough pairs, make sure the most electronegative elements are satisfied. Then, start shifting pairs into double and triple bonds to satisfy the octet rule. • If there are extra e-, stick them on the central atom.

25. Hints: • H is NEVER a central atom! • Halogens (Group 17) are usually not central atoms. • If you only have 1 of a certain element, it is usually the central atom.

26. Checking Your Work! • But Remember.... • The Structure MUST Have: the right number of atoms for each element, the right number of electrons, the right overall charge, and 8 electrons around each atom (ideally).

27. Examples: F2 H2O OCl- PO43-

28. Examples: O2 CH4 HF NH3

29. Examples: NH4+ SO32- N2 CH3OH

30. BF3 Exceptions to the Octet Rule Reduced Octets – electron deficient molecules (Be and B) Be: 2 valence e-, doesn’t form octet (BeH2: Be has 4 e-) B: 3 valence e-, doesn’t form octet (BF3: B has 6 e-)

31. Exceptions to the Octet Rule Expanded Octets (Examples: P, S, Cl, As, Se, Br, Kr, Xe) How to recognize: • The central atom in PERIOD 3 or greater is surrounded by > 4 atoms. • You draw the Lewis diagram and the results don’t make sense – the central atom has > 8 e-

32. Expanded Octets(P, S, Cl, As, Se, Br, Kr, Xe) Examples: PF5 XeF4

33. Resonance Structures Definition: When a single Lewis structure does not adequately represent a substance, the true structure is intermediate between two or more structures which are called resonance structures. Resonance Structures are created by moving electrons, NOT atoms.

34. Resonance Structure Example, SO2 Central atom = S This leads to the following structures: These equivalent structures are called RESONANCE STRUCTURES. The true structure is a HYBRID of the two. Arrow means “in resonance with”

35. Resonance Structure Example, NO3- Draw the Lewis diagram for NO3- with all possible resonance structures.

36. Radicals When there is an odd # of total electrons, there will be a single, unpaired electron in the structure! Example: NO Radicals are extremely reactive: they want to have paired electrons!!

37. Linus Pauling, 1901-1994 The only person to receive two unshared Nobel prizes (for Peace and Chemistry). Chemistry areas: bonding, electronegativity, protein structure

38. Electronegativity Definition: A measure of the ability of an atom in a molecule or bond to attract electrons to itself. Scale proposed by Linus Pauling ***Greater E.N. means element more strongly attracts electrons.

39. Electronegativity Trends on periodic table: Highest on upper right (F has highest with e/n = 4.0) Lowest on lower left (Francium = 0.7) Noble gases have ZERO E.N.

40. Electronegativity

41. Bond Polarity • Polar Covalent Bond • Covalent bond in which the electrons are unequally shared • Ex. H2O • Non-polar Covalent Bond • Covalent bond in which the electrons are equally shared • Ex. F2 or CH4 • Predicting Bond Polarity • Use Electronegativity!! (see next slide)

42. Predicting Bond Polarity • Calculate the difference between the Pauling electronegativity values for the 2 elements 0 – 0.4  Non-polar covalent 0.4 – 1.7  Polar covalent (more e/n element has greater pull) 1.7 and up  Ionic (e- are transferred between atoms)

43. Using e/n to predict polarity of individual bonds A polar bond has a partial charge due to unequal sharing of electrons. A polar bond is shown using partial charges either with delta or cross/arrow. Negative delta or arrow next to more E.N. atom.

44. Bond Polarity HCl is POLAR because it has a positive end and a negative end. Cl has a greater share of bonding electrons than H. Cl has slight negative charge (d-) and H has slight positive charge (d+)

45. Bond Polarity What type of bonds are these? O—H O—F E.N. 3.5 - 2.1 3.5 - 4.0  1.4 0.5

46. Molecular Geometry Molecular Geometry describes the 3-D arrangement of atoms in a molecule. We will use VSEPR theory to predict these 3-D shapes!

47. VSEPR: Shapes of Molecules • VSEPR Theory (definition) • “Valence Shell Electron Pair Repulsion” • Based on idea that e- pairs want to be as far apart as possible • The molecule adopts the shape that minimizes the electron pair repulsions. • Based on molecular shape of Lewis diagram

48. We define the electron pair geometry by the positions in 3D space of ALL electron pairs (bonding and non-bonding). The molecular geometry only considers the positions of the bonded electrons.

49. To determine the electron pair geometry: • 1. Draw the Lewis structure. • 2. Count the number of bonded (X) atoms and non-bonded or lone pairs (E) around the central atom. • 3. Based on the total of X + E, assign the electron pair geometry. • 4. Multiple bonds count as one bonded pair!

50. Electron-pair geometry around a central atom Sum of X + EShapes 2 linear 3 trigonal planar 4 tetrahedral 5 trigonal bipyramidal 6 octahedral