1. Chemical Bonding Chapter 6 Sections 1, 2, and 5

2. Chemical Bonds A chemical bond is the mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together Noble gases tend not to do this because of their filled s and p orbitals. They have a stable octet: outer s and p orbitals are completely filled with e-’s ( 8 total)

3. Chemical Bonds Atoms that don’t have a stable octet are more reactive because their potential energy is higher. They become more stable by decreasing their potential energy. Octet Rule: chemical compounds tend to form so that each atom has an octet of e-’s in its highest occupied energy level How to do this? Gain, lose or share electrons between atoms

4. Chemical Bonds By forming a chemical bonds, atoms gain stability! Chemical changes always involve energy What type of bonds can be formed? Ionic bond Covalent bond Nonpolar covalent Polar covalent

5. Chemical Bonds Ionic bonding: bonds that result from electrical attractions between cations and anions Covalent bonding: sharing of electron pairs between 2 or more atoms *** In reality, bonding is often somewhere between the two extremes***

6. Two types of Covalent Bonds Nonpolar –covalent: equal sharing of electron pairs Polar-covalent: unequal attraction for the shared electrons

7. How can we determine the type of bond? Knowing how strong an atom’s ability is to attract electrons (aka electronegativity), helps us determine if it will form a ionic or covalent bond with another atom. A large difference in E.N. between atom’s will result in an ionic bond A small difference between atom’s will result in a form of covalent bonding

8. Nonpolar Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e - Increasing difference in electronegativity What type of Bond is it? Electronegativity Difference Bond Type 0 to 0.3Nonpolar Covalent 0.4 to 1.7Polar Covalent  1.7 Ionic

9. Do you see any trends? A metal and nonmetal tend to form ionic compounds Nonmetal and nonmetal tend to form polar-covalent or nonpolar- covalent compounds

10. H F F H Polar covalent bond or polar bond : covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - riche - poor ++ --

12. Classify the following bonds as ionic, polar covalent,or covalent: Cs – 0.7Cl – 3.03.0 – 0.7 = 2.3Ionic H – 2.1S – 2.52.5 – 2.1 = 0.4Polar Covalent N – 3.0 3.0 – 3.0 = 0 Nonpolar Covalent CsCl H2SH2S N2N2

13. Properties of Molecular Covalent Compounds Not very soluble in water Do not conduct electricity Low melting points Low boiling points Can be solids, liquids and gases at room temperature

14. Comparison of Ionic and Covalent Compounds

15. Types of Crystals

16. Lewis Structures

17. Review: What are valence electrons? Lewis Dot Diagrams − an electron-configuration notation with only the valence electrons of an element are shown, indicated by dots placed around the element’s symbol. − the inner core electrons are not shown.

19. Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE s2p6s2p6 This is called the octet rule The Octet Rule Bond formation follows 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 valence energy level.

20. Lewis Structures for Compounds The pair of dots between two symbols represents the shared pair of a covalent bond. Each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds. An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

21. Lewis Structures The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash.

22. covalent bond : is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond

23. Multiple Covalent Bonds double covalent bond or double bond : covalent bond in which two pairs of electrons are shared between two atoms shown by two side-by-side pairs of dots or by two parallel dashes

24. Multiple Covalent Bonds triple covalent bond or triple bond : covalent bond in which three pairs of electrons are shared between two atoms.

25. 8e - H H O ++ O HH O HHor 2e - Single Bond – two atoms share one pair of electrons Double Bond – two atoms share two pairs of electrons single covalent bonds O C O or O C O 8e - double bonds Triple Bond – two atoms share three pairs of electrons N N 8e - N N triple bond or

26. Bond Type Bond Length (pm) C-CC-C 154 CCCC 133 CCCC 120 C-NC-N 143 CNCN 138 CNCN 116 Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond

27. 1.Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. 2.Count total number of valence e -. Add 1 for each negative charge. Subtract 1 for each positive charge. 3.Complete an octet for all atoms except hydrogen 4.If structure contains too many electrons, form double and triple bonds on central atom as needed. Writing Lewis Structures

28. Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 1 – N is less electronegative than F, put N in center FNF F Step 2 – Count valence electrons N - 5 (2s 2 2p 3 ) and F - 7 (2s 2 2p 5 ) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

29. Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is less electronegative than O, put C in center OCO O Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e - 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e - 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24

30. resonance structure: one of two or more Lewis structures for a single molecule can be drawn to represent a molecule OOO + - OOO + - OCO O -- OCO O - - OCO O - - What are the resonance structures of the carbonate (CO 3 2 -) ion?

31. Exceptions to the Octet Rule The Incomplete Octet HHBe Be – 2e - 2H – 2x1e - 4e - BeH 2 BF 3 B – 3e - 3F – 3x7e - 24e - FBF F 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24

32. Exceptions to the Octet Rule Odd-Electron Molecules N – 5e - O – 6e - 11e - NO N O The Expanded Octet (central atom with principal quantum number n > 2) SF 6 S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48

33. Molecular Geometry Chapter 6.5

34. VSEPR THEORY Lewis Dot Diagrams are 2D but we live in a 3D world. How are molecules actually arranged?? Follows the Valance Shell Electron Pair Repulsion Theory or VSEPR

35. AB 2 – Linear Number of Surround Atoms Number of Lone PairsBond Angle 20180˚ Cl Be

36. AB 3 – Trigonal Planar Number of Surround Atoms Number of Lone PairsBond Angle 30120˚

37. AB 2 E 1 – Bent Number of Surround Atoms Number of Lone PairsBond Angle 21<120˚

38. AB 4 – Tetrahedral Number of Surround Atoms Number of Lone PairsBond Angle 40109.5˚

39. AB 3 E 1 – Trigonal Pyramidal Number of Surround Atoms Number of Lone PairsBond Angle 31107˚

40. AB 2 E 2 – Bent Number of Surround Atoms Number of Lone PairsBond Angle 22104.5˚

41. Predicting Molecular Geometry 1.Draw Lewis structure for molecule. 2.Count number of lone pairs on the central atom and number of atoms bonded to the central atom. 3.Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO 2 and SF 4 ? SO O AB 2 E bent S F F F F AB 4 E distorted tetrahedron

42. H F electron rich region electron poor region   The electronegativity of an atom will create a dipole, or polar molecule.

43. Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H dipole moment polar molecule S O O CO O no dipole moment nonpolar molecule dipole moment polar molecule C H H HH no dipole moment nonpolar molecule

44. Intermolecular forces: attractive forces between molecules. Intramolecular forces: hold atoms together, attractive forces within a molecule. Generally, intermolecular forces are much weaker than intramolecular forces.

45. Properties of Ionic Compounds Combination of ions (cation/anion) Hard and Brittle Tightly packed solids in a crystal lattice Usually soluble in water Conducts electricity when dissolved High melting points

46. Breaking Ionic Bonds Ionic Bonds are very tightly bound positive and negative attraction A LOT of energy needs to be put in to break an ionic bond How does this affect the melting point?