Unit 6: Properities of Compounds (Chemical Bonding and Molecular Structure) 10-30-07

1. Unit 6:Properities of Compounds(Chemical Bonding and Molecular Structure)10-30-07 Cartoon courtesy of NearingZero.net

2. Bond Length Diagram

3. Bonding Forces • Electron – electron • repulsive forces • Proton – proton repulsive forces • Electron – proton attractive forces

4. Table of Electronegativities

5. Figure 11.4: The three possible types of bonds: (a) a covalent bond formed between identical atoms; (b) a polar covalent bond, with both ionic and covalent components; and (c) an ionic bond, with no electron sharing.

6. Forces that hold groups of atoms together and make them function as a unit. Bonds IntrAmolecular forces • Ionic bonds – transfer of electrons • Covalent bonds – sharing of electrons (Polar and Nonpolar) • Metallic bonds-a sea of electrons

8. The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

9. Formula Unit is the smallest representative particle ie: NaCl • Polar • High melting point and boiling point • Poor conductors of heat and electricity in the solid state. • Good conductors if electricity in the liquid state (Molten) and in a solution. • Unit Cell is the smallest representative particle of a crystal (c) • EN›1.7 = Ionic Bond • EN<1.7= Covalent bond • The formation of ionic bonds is always exothermic!

10. Metallic bonds • Metallic substances can be compared to a group of nuclei surrounded by a sea of electrons. • Very hard, high melting and boiling point, good conductor of electricity and heat due to the freedom of its electrons.

13. HNFOIClBr HONClFIBr (say HONKLE-fibber) BrINClHOF (say Brinckle-hoff) Have No Fear Of Ice Cold Beverages Never Have Fear Of Ice Cold Beverages I Have No Bright Or Clever Friends ClIF H Bron HOFBrINCl Twins (twins because they exist in pairs) There are seven such elements. The first one is the first element Hydrogen; the rest form a 7 on the periodic table: N, O, F across, then going down Cl, Br, I. Diatomic molecules:I kNow Bro. Flo, Clo and HOI2 N2 Br2 F2 Cl2 H2 O2

14. Diatomic Molecules • HI BrONClF • I kNOw Bro, Flo, Clo and Ho I2 N2 O2 H2 Cl2 F2 Br2

15. Molecule-Ion Attractions This "molecule-ion attraction" is appropriately named, as a molecule (water) is attracted to an ion.*

18. 1. Covalent bonds - between NM-NM, bonding electrons shared (sometimes more sometimes less equally) Covalent compounds gain an octet of electrons through sharing electrons, i.e. through forming chemical bonds It is possible for nonmetals to share more than one pair of electrons (single, double and triple bonds) 4. Most molecular compounds do not dissociate since they are not ionic, i.e. solutions of molecular compounds do not conduct electricity 5. The bonding forces between molecules are weaker than between ions, so as a rule molecular compounds have lower MP/BP than ionic compounds Characteristics of Covalent Bonds and Molecular Compounds

19. Figure 11.4: The three possible types of bonds: (a) a covalent bond formed between identical atoms; (b) a polar covalent bond, with both ionic and covalent components; and (c) an ionic bond, with no electron sharing.

21. Covalent Bonds Polar-Covalent bonds • Electrons are unequally shared Nonpolar-Covalent bonds • Electrons are equally shared

22. 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals. Comments About the Octet Rule

23. A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. Polarity

24. Figure 11.2: Probability representations of the electron sharing in HF. (a) What the probability map would look like if the two electrons in the H–F bond were shared equally. (b) The actual situation, where the shared pair spends more time close to the fluorine atom than to the hydrogen atom.

28. Bond Energy • It is the energy required to break a bond. • As the bond strength INcreases, so the bond length DEcreases. • The carbon = carbon triple bond is stronger than the C=C double bond which is stronger than the C-C single bond

29. Completing a Lewis Structure -CH3Cl • Make carbon the central atom • Add up available valence electrons: • C = 4, H = (3)(1), Cl = 7 Total = 14 • Join peripheral atoms • to the central atom • with electron pairs. H .. .. .. H .. • Complete octets on • atoms other than • hydrogen with remaining • electrons C .. Cl .. .. H

30. Polarity Demostration

34. Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

38. The structure around a given atom is determined principally by minimizing electron pair repulsions. Predicting a VSEPR Structure Draw Lewis structure. Put pairs as far apart as possible. Determine positions of atoms from the way electron pairs are shared. Determine the name of molecular structure from positions of the atoms. VSEPR Model (Valence Shell Electron Pair Repulsion)

41. Multiple Covalent Bonds:Double bonds Two pairs of shared electrons

42. Multiple Covalent Bonds:Triple bonds Three pairs of shared electrons

43. Occurs when more than one valid Lewis structure can be written for a particular molecule. Resonance • These are resonance structures. • The actual structure is an average of • the resonance structures.

46. Molecular Shapes from VSEPR theory

49. VSEPR and the water molecule (Bent / polar)

50. Ozone and CFC Reaction