1. Bonding GPS 8

2. Why do atoms bond together? Octet Rule – an atom that has a full outer-most energy level is unreactive (usually it is full with 8 electrons, but there are some exceptions) all atoms want to satisfy the octet rule atoms react with each other in order to satisfy the octet rule

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 When electrons are transferred, ions are formed. Ions are attracted to each other and form ionic bonds When electrons are shared, covalent bonds are formed

5. Chemical Properties of Groups Oxidation Numbers

6. Ionic Bonding – occurs between a metal ions and a nonmetal ions – involves a transfer of electrons – ions are separate, but are attracted to one another because of opposite charges – ions form a compound

7. Ionic Bonds: Lewis Dot Structure For sodium and phosphorus:

8. Ionic Bonding

9. Covalent Bonding Covalent bonding – occurs between two nonmetal atoms – involves a sharing of electrons – combine to form a molecule

10. Covalent Bonding Hydrogen and chlorine combine to form hydrochloric acid. H + Cl → H Cl

11. Covalent Bonding: Lewis Dot Diagrams

12. Metallic Bonding – Occurs between atoms of a metal – Valence electrons are shared between the cations

13. Physical Properties of Ionic and Covalent Compounds Property:Ionic Compound Covalent Compound Phase at room temperature Usually solidUsually gas DensityMore denseLess dense Melting pointHigher Temps.Lower Temps. Boiling pointHigher Temps.Lower Temps. ConductivityConductsNonconductive

14. 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 each element’s symbol 2. Find the number of valence electrons for each element (recall: count across the periodic table) 3. Draw a dot beside the symbol for each element to represent valence electrons as you position atoms so that they all fulfill the octet rule. **Helpful tip: when you have more than two atoms bonded, put the element with the least number of atoms in the center**

16. Lewis Dot Diagrams: Covalent Bonding Examples: HF Br 2 CCl 4 H 2 O

17. Lewis Dot Diagrams: Covalent Bonding Some molecules can also form double bonds or triple bonds – Example of a double bond: CO 2 – Example of a triple bond: N 2

18. Exceptions to the octet rule: Hydrogen (H)2 Beryllium (Be)4 Boron (B)6 Gallium (Ga)6

19. Lewis Dot Diagrams: Ionic Bonding Examples: NaCl MgO

20. Summary: Ionic Bonding -Metal and nonmetal -Electrons transferred -Smallest unit: ions Covalent Bonding -Two nonmetals -Electrons shared -Smallest unit: molecule

21. 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.

22. Molecular Shapes Linear Bent Trigonal planar TetrahedralTrigonal pyramidal

23. Molecular Shape Steps to finding molecular shape: 1. Draw the Lewis Dot structure for the molecule 2. Find the number of unshared pairs of electrons on the central atom 3. Determine the shape 4. Draw a line structure, showing unshared pairs of electrons on the central atom as a “balloon”

24. Molecular Shape Determine the molecular shape for the following GaF 3 HCl CF 4 BeBr 2 PCl 3 H 2 O bent linear tetrahedral linear trigonal pyramidal trigonal planar

25. Exceptions to the octet rule: Hydrogen (H)2 Beryllium (Be)4 Boron (B)6 Gallium (Ga)6 Note: When Be, B, and Ga are sharing electrons in a molecule, they are behaving like nonmetals, and should be treated as covalent molecules

26. Polarity Polar covalent bond – occurs in molecules in which atoms have a difference in electronegativity – Polar molecules are called Dipoles Polarity is shown using: – 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

27. Polarity Example: The water molecule

28. Polarity Example: Carbon monoxide

29. Polarity Example: Hydrogen gas Nonpolar

30. Polarity Example Methane (CH 4 ) Nonpolar

31. Polarity Decide if the following molecules are polar or nonpolar. GaF 3 HCl CF 4 BeBr 2 PCl 3 H 2 O polar nonpolar polar nonpolar polar nonpolar

32. Bond Type using Electronegativity values Bond TypeElectronegativity difference nonpolar covalentLess than 0.5 polar covalentbetween 0.5 and 2.0 ionicGreater than 2.0

33. Bond Type using Electronegativity values Examples: For the following, compute the electronegativity difference. Then, tell what type of bond will form. SrO NaF O 2 AlAs 2.5ionic 3.0ionic 0.0nonpolar covalent 0.6polar covalent

34. Bond Type using Electronegativity values Example: * Of the following, which has the greatest ionic character? * Which has the greatest covalent character? SrO NaF O 2 AlAs NaF has the greatest ionic character O 2 has the greatest covalent character

35. Electronegativity Values

36. Hybrid Orbitals Hybrid orbitals – when atoms bond, their orbitals become disturbed and therefore combine, creating hybrid orbitals. – the type of hybrid orbital depends on the number of electron domains (a bonded or nonbonded pair of electrons) 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

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

38. 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