1. Bonding GPS 8

2. Why do atoms bond together? Octet Rule – an atom is most stable when it has a _______________ – Most atoms’ outer energy level are full with 8 electrons, but there are some exceptions such as Hydrogen and Helium that have a full outer energy level with only 2 electrons Atoms will react with one another to obtain a full octet because they are more stable having a full octet

3. Valence Electrons as Lewis Dot Diagrams # of valence electrons: 1 2 3 4 5 6 7 8

4. Octet Rule: Forming Bonds Atoms will gain, lose, or share electrons in order to fulfill the octet rule (to obtain a full outer energy level) When electrons are _______, ____________re formed When electrons are ______________, ions are formed. Ions are attracted to each other and form ____________

5. Ionic Bonding: Oxidation numbers Oxidation Numbers

6. Ionic Bonding – often occurs between metal ions and nonmetal ions – involves a __________ of valence electrons – ions are separate, but are attracted to one another because of opposite charges (one + ion, one - ion) – ions form a compound called a solid crystal lattice Example: Na and Cl react to form sodium chloride (NaCl) Solid NaCl crystal lattice

7. Ionic Bonding: A little funny

8. Ionic Bonding: Electronegativity Difference Ionic bonds form due to a transfer of electrons. Electrons are transferred from one atom to another because of a relatively ________difference in electronegativity between a metal and nonmetal.

9. Covalent Bonding – Usually occurs between two nonmetal atoms – involves a _________of electrons – combine to form a molecule

10. Covalent Bonding: Electronegativity Difference Covalent bonds form due to a sharing of electrons. Electrons are shared between atoms because of a relatively _______ difference in electronegativity between nonmetals.

11. Metallic Bonding – Occurs between atoms of a metal – Valence electrons (delocalized electrons) are shared between the metal cations

12. Physical Properties of Ionic and Covalent Compounds Ionic Compound: NaCl Ionic Compound: CaCl 2 Covalent Compound (Molecule): H 2 O Covalent Compound (Molecule): CO 2

13. Physical Properties of Ionic and Covalent Compounds Property:Ionic Compound Covalent Compound Phase at room temperature often solidsoften gases, some liquids DensityMore denseLess dense Melting pointHigher Temps.Lower Temps. Boiling pointHigher Temps.Lower Temps. ConductivityConductsNonconductive

14. Periodic Trends: Diatomic Molecules Naturally occurring diatomic molecules: Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2 - These occur as gases at room temperature * except Br 2 (liquid at room temp) * except I 2 (solid at room temp) - Naming: Ex: O 2 oxygen gas N 2 nitrogen gas Br 2 liquid bromine I 2 solid iodine

15. Lewis Dot Diagrams Steps: 1.Find the number of valence electrons for each element 2.Add to get the total number of valence electrons for that molecule (if representing more than just a single atom) 3.When you have more than two atoms bonded, the element with the least number of atoms usually goes in the center of the molecule 4.Place two electrons (dots) between atoms to represent a single bond between them (sharing two electrons) 5.Place the rest of the electrons around the atoms so that they are all fulfilling the octet rule (keeping in mind Be, B, and Ga are exceptions). Any extra electrons go on the central atom.

16. Lewis Dot Diagrams: Covalent Bonding Examples: HFI Br 2 Si CCl 4 N H 2 OMg

17. Exceptions to the octet rule: Number of electrons in outer energy level when bonded: Beryllium (Be)4 Boron (B)6 Gallium (Ga)6 Most atoms bond in order to obtain a full octet (a full outer energy level, usually 8 valence electrons). This makes atoms more stable. However, there are a few exceptions:

18. Exceptions to the octet rule: Examples: BeCl 2 BF 3 GaBr 3

19. Violating the Octet Rule: Some atoms can violate the octet rule, meaning they can accommodate MORE than 8 valence electrons. Atoms that have available space in an existing d-sublevel (or can create a d-sublevel) can be stable with more than 8 valence electrons. Only elements on the 3 rd period or periods below the 3rd can violate the octet rule because only these have access to a d-sublevel in their structure.

20. Violating the Octet Rule: Examples: PCl 5 RnCl 2 SF 6

21. Dot Diagrams of Ions Recall that an ion is an atom or molecule with a charge. It is charged because electrons have been lost or gained. Examples: NH 4 + OH - ICl 4 -

22. Lewis Dot Diagrams: Multiple Bonding Some molecules can also form multiply bonds (double bonds or triple bonds) Example of a double bond: CO 2 Example of a triple bond: N 2

23. Connecting Dot Structures to Molecular Particle Views: A molecule of O 2 gas: Dot Diagram Line Structure Particle View A molecule of H 2 gas: A molecule of CO 2 gas:

24. Lewis Dot Diagrams: Ionic Bonding Recall that electrons are transferred to create ionic bonds (rather than being shared like in covalent bonds). An ionic bond is the electrostatic attraction of the oppositely- charged ions that remain after electrons have been transferred between atoms Examples: NaCl MgO MgCl 2

25. Connecting Dot Structures to Ionic Particle Views: Solid Sodium Chloride: Crystal Lattice structure, NaCl

26. Bonding Type Summary: Ionic Bonding -Metal and nonmetal -Electrons transferred -Large difference in electronegativity between metal and nonmetal -Smallest unit: ions Covalent Bonding -Two nonmetals -Electrons shared -Small difference in electronegativity between two nonmetals -Smallest unit: molecule

27. Molecular Shapes Valence-shell electron pair repulsion theory (VSEPR) – states that in a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible. Each “balloon” below represents a bonded atom’s electron cloud. The atoms bond in a way that maximizes the distance between electron clouds while still allowing atoms to bond. 2 atoms 3 atoms 4 atoms 5 atoms 6 atoms

28. Molecular Shapes

29. Molecular Shape Steps to finding molecular shape: 1.Draw the dot structure for the molecule 2.Count the number of atoms in the molecule and the number of unshared pairs of electrons on the central atom 3.Determine the shape by imagining the electron pairs around the central atom as far apart as possible while still bonded to the central atom 4.Draw a line structure depicting the shape, showing unshared pairs of electrons on the central atom as a “balloon”

30. Molecular Shape Determine the molecular shape for the following GaF 3 PCl 5 HCl CF 4 RnCl 2 BeBr 2 PCl 3 SF 6 H 2 O Note: Although Be, B, and Ga are metals, they share electrons with nonmetals (forming covalent bonds) due to the relatively small difference in electronegativity between Be, B, or Ga and the nonmetal.

31. Polar Covalent Bonds Polar covalent bond – occurs between atoms that have a relatively moderate difference in electronegativity Polarity is shown using symbols: δ - means “partially negative” δ + means “partially positive” The atom(s) with higher electronegativity will attract electrons more strongly and have the partially negative sign More electronegative atom

32. Polar Molecules (Dipoles) Molecules are nonpolar if all of the individual bond polarities do cancel out (due to the symmetry of the molecule). Molecules are polar if all of the individual bond polarities do not cancel out. These molecules have a dipole moment.

33. Polarity Decide if the following molecules are polar or nonpolar: GaF 3 PCl 5 HCl CF 4 RnCl 2 BeBr 2 PCl 3 SF 6 H 2 O

34. Polarity Decide if the following molecules are polar or nonpolar: CH 3 FPCl 2 F BeBrCl SF 5 Br

35. Bond Type: Electronegativity Difference Bond TypeElectronegativity difference nonpolar covalentLess than 0.4 polar covalentbetween 0.4 and 2.0 ionicGreater than 2.0

37. Bond Type: Electronegativity Difference For the following, determine the bond type using electronegativity values: SrO NaF O 2 AlAs

38. Bond Type: Electronegativity Difference Of the following, which bond has the greatest ionic character? Which bond has the greatest covalent character? SrO NaF O 2 AlAs

39. Hybrid Orbitals When atoms bond, their orbitals become disturbed and therefore combine, creating hybrid orbitals. sp hybrid orbitals sp 2 hybrid orbitals

40. Example of sp 3 hybrid orbitals: nitrogen in ammonia, NH 3 Hybrid Orbitals

41. Hybrid Orbitals Hybrid Orbitals used by central atom Shape/GeometryNumber of Electron Domains around central atom Example splinear2BeF 2 sp 2 trigonal planar3BCl 3 sp 3 Trigonal pyramidal4NH 3 sp 3 tetrahedral4CH 4 The type of hybrid orbital depends on the number of electron domains (a bonded or nonbonded pair of electrons)

42. Hybrid Orbitals What type of hybrid orbitals would the central atom use in each of the following? CCl 4 BeBr 2 BI 3 PCl 3

43. Molecular Shape and Hybrid Orbitals ShapeAtoms bonded to central atom Lone pairs of electrons on central atom Hybridization Linear20sp Bent21sp 2 Trigonal Planar30sp 2 Tetrahedral40sp 3 Trigonal pyramidal31sp 3 Bent22sp 3

44. Bond Angles Questions: Why is the bond angle smaller in a trigonal pyramidal molecule than in a tetrahedral molecule? Why is the bond angle smaller in a bent molecule than in a trigonal pyramidal molecule? Answer (for both questions above): Nonbonded (unshared) pairs of electrons have a stronger repulsive force than the electron cloud from bonded atoms around the central atom. Since these unshared pairs of electrons have a stronger repulsive force, they push the other bonded atoms closer together.

45. Bond Length: Groups/Families Bond length increases as you move down a group of the periodic table Example: Br 2 has a longer bond length than F 2

46. Bond Length: Single vs Multiple Bond Single bonds are the longest. Triple bonds are the shortest. Examples: N≡N bond is shorter than O=O bond O=O bond is shorter than Br–Br bond Single Bond: Double Bond: Single Double Triple

47. Bond Length Determine whether Cl 2 or I 2 has a longer bond length. Justify your answer.

48. Bond Length Determine whether F 2 or N 2 has a longer bond length. Justify your answer.

49. Correlating Bond Length to Bond Energy Atoms bond at a distance that maximizes attraction (between nucleus and electrons) and minimizes electron repulsion (between electron clouds) At this distance, the least amount of energy is required and the molecule is the most stable

50. Enthalpy (∆H) from Bond Energies ∆H = Energy of Bonds Broken – Energy of Bonds Formed Enthalpy, ∆H, can be used to determine if a reaction is endothermic or exothermic - ∆H means the reaction is ____________ (releases energy) + ∆H means the reaction is ____________(absorbes energy)

51. Enthalpy (∆H) from Bond Energies Example: Use the table of bond energies to calculate ∆H for the formation of HF from its elements. H 2 (g) + F 2 (g) → 2HF(g) Average Bond Energy (kJ/mol) H-H432 H-F565 F-F154

52. Enthalpy (∆H) from Bond Energies Example: Using the bond energies listed, calculate the enthalpy for the reaction of methane with chlorine and fluorine to give Freon-12 Average Bond Energy (kJ/mol) C-H413 Cl-Cl239 F-F154 C-F485 C-Cl339 H-F565 H-Cl427 CH 4 (g) + 2Cl 2 (g) + 2F 2 (g) → CF 2 Cl 2 (g) + 2HF(g) + 2HCl(g)