Chemical Bonding

1. Chemical Bonding

2. Chemical Bonds • Compound are formed from chemically bound atoms or ions • Bonding only involves the valence electrons

3. Chemical Bonds • Defn – force holding two atoms together • How are they formed? Atoms gain, lose, or share valence electrons • Why does bonding occur? Stability – achieve octet rule

4. X 2s 2p Lewis Structures • Electron Dot Diagrams • show valence e- as dots • distribute dots like arrows in an orbital diagram • 4 sides = 1 s-orbital, 3 p-orbitals • EX: oxygen O

5. Ne Lewis Structures • Octet Rule • Most atoms form bonds in order to obtain 8 valence e- • Full energy level stability ~ Noble Gases

6. Electron Dot Structure • Shows valence electrons around atomic symbol hydrogen (group 1) H • • • • N • nitrogen (group 5) • • • • (group 7) • Cl chlorine • • •

7. Types of Chemical Bonds • 3 Types • ionic bond • covalent bond • metallic bond

8. Ionic Bonding

9. Ionic Bonding Vocabulary Ionic compounds are referred to as Formula Units. Compounds are composed to two or more elements. Binary Compound – 2 elements - NaCl Ternary Compound – 3 or more elements – NaHCO3 Ion – A charged atom Monatomic Ion – 1 atom Na1+ Polyatomic Ion - 2 or more atoms NO31-

10. Ionic Bond • Defn – force holding cations and anions together • A • • B A+ B- • Ionic bond Cations – positively charged ions Anions – negatively charged ions

11. Ionic Bond • Where are these bonds found? InIonic Compounds

12. Ionic Bonding Properties

13. Ionic Bonding giving/taking of valence electrons • What’s going on? • If I gave you a compound, how can you tell if it is ionic or not? combo of metal + nonmetal

15. NaCl Formation of Ionic Bonds 1- • • • • • • Na Cl Na1+ • • • + + Cl • • • • • • • 2s22p63s1 3s23p5 2s22p6 3s23p6 8 v.e.

17. CaBr2 Formation of Ionic Bonds 1- • • • • • • • Ca Br • • • + Br • • • • • • • Ca2+ + • • • • Br 1- • • • • • • • Br • • • •

18. Ionic Formulas • Don’t show charges in the final formula. • Overall charge must equal zero. • If charges cancel, just write symbols. • If not, use subscripts to balance charges. • Use parentheses to show more than one polyatomic ion. • Stock System - Roman numerals indicate the ion’s charge.

19. Ionic Formula Names • Write the names of both ions, cation first. • Change ending of monatomic anions to -ide. • Polyatomic ions have special names. • Stock System - Use Romannumerals to show the ion’s charge if more than one is possible. Overall charge must equal zero.

20. Common Ion Charges(Oxidation States) 1+ 0 2+ 3+ 4+/4- 3- 2- 1-

21. Ionic Formulas • potassium chloride • magnesium nitrate • copper(II) chloride  KCl • K+ Cl- • Mg2+ NO3-  Mg(NO3)2  CuCl2 • Cu2+ Cl-

22. Ionic Formula Names • NaBr • Na2CO3 • FeCl3 • sodium bromide • sodium carbonate • iron(III) chloride

23. Metallic Bonding

24. Metallic Bonding • Defn – attraction of metallic cations Occurs only in metals

25. Metallic Bonding Properties

26. Metallic Bonding • Defn – bond formed from attraction between positive nuclei and delocalizedelectrons • holds metals together • Delocalized Electrons – electrons detached from parent atom • “lost electron away from home”

27. Electron Sea Model • Defn – electrons move freely within other molecular orbitals

28. Properties of Metals • Electron sea model gives metals certain physical properties • Shiny – due to photoelectric effect • Conduct electricity and heat – electrons move easily from one place to another • Malleable (pound into sheets) • Ductile (put into wires)

29. Why malleable and ductile? atoms can also move from one place to another and still remain in contact with and bonded to the other atoms and electrons around them Shape #1 Shape #2 shifted atoms

30. Covalent Bonding

31. Covalent Bond • Defn – two atoms share one pair of electrons Electrons shared • A • • B A B •

32. Covalent Bonds • Where are these bonds found? - molecules (molecular compounds) - polyatomic ions

33. Covalent Bonding • What’s going on? • Molecule – formed when 2 or more atoms bond covalently Sharing of electrons

34. Covalent Bonding Properties

35. Two Types of Covalent Bonds i) nonpolar covalent – equal sharing of e- ii) polar covalent – UNequal sharing of e-

36. Nonpolar vs. Polar NONPOLAR POLAR

37. Nonpolar vs. Polar

38. Nonpolar vs. Polar

39. Single Bond • Defn – one pair (2) of e- shared • Lewis Structures – represents how atoms in molecules are arranged • atoms MUST obey octet rule (except hydrogen)

40. Lewis Structures • bonded electrons – occur between bonded atoms • A B or A B • single bond

41. Lewis Structures • Unshared or Lone Pairs – electron pairs NOT involved in bonding • • • • • A B A B • • • • • lone pairs

42. Lewis Structures Examples • H2O (8 valence e- or 4 pairs) • • • • • • • • • • H O H H O • • • • • • H • • • H O • H

43. Lewis Structures Examples • NHF2 (20 v.e.or 10 pairs) • • • • • • • • • • • • • • F F N F • • • • • • • • • • • • • • • • • • N H H • • • • • F • • • • • • • • • • • F N F • • • • • • H

44. Multiple Covalent Bonds • Double Bond – two pairs (4) e- shared • • A B A B • • O2 (12 v.e. = 6 pairs) • • • • • • • • • • • • • • • O • • • • • • • O • O O O O • • • • • • • • • • • • • • • • • • • O O • •

45. Multiple Covalent Bond • Triple Bond – three pairs (6) e- shared • • • A B A B • • • N2 (10 v.e. = 5 pairs) • • • • • • • • • • • • • • • • • • • • N • • N • N N N N • • • • • • • • • • • • • • N N • • •

46. Comparing single, double, and triple bonds • Bond Strength • Bond Length Triple > Double > Single Single > Double > Triple The shorter the bond, the stronger it is

47. Polyatomic Ions • Defn – CHARGED group of atoms covalently bonded - ex: SO42-, NH41+, NO31-

48. Polyatomic Ions SO42- (32 v.e. = 16 pairs) 2- • • 2- • • • • O • • • • O • • • • • • • • • • • • • • • • O S O • • • • • • O S O • • • • • • • • • • • • • • O • • • • O • • • • • •

49. Polyatomic Ions NH41+ (8 v.e. = 4 pairs) 1+ 1+ H H • • • • H N H • • H N H • • H H

50. Using electronegativity to determine bond type • Recall electronegativity: how much an atom wants electrons • Each atom is assigned a number between 0-4.0 to determine electronegativity strength