Chemical Bonding

1. Chemical Bonding

2. Introduction to Bonding

3. Neutral or “free” atoms are rarely found in nature • Most elements exist as part of a compound • Compounds are chemically bonded atoms

4. Chemical bond: strong attractive force that exists between atoms or ions in a compound.

5. So, why do atoms form bonds?

6. Atoms form bonds to become more stable.

7. Bonding involves only valenceelectrons

8. Remember… • The noble gases are particularly stable because their outer shell is full of electrons (usually 8)

9. Octet Rule: • Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons(noble gas configuration)

10. Types of Chemical Bonds • Ionic • Valence electrons are transferredfrom one atom to another. • Usually formed between a metal and a non-metal.

11. This creates + ions and - ions which are then electrostatically attracted.

12. 2. Covalent • Pairs of electrons areshared between atoms. • Usually formed between two non-metals.

13. Two Types of Covalent Bonds A. Polar Covalent • Electrons are sharedunequally. • Usually between two different nonmetals.

14. B. Non-Polar Covalent • Electrons are sharedequally. • Usually between atoms of the same element.

15. Metallic • attraction between a metallic cation and delocalized electrons. • Delocalized electrons – valence electrons not held by any specific atom, but free to move from one atom to another. • “Electron Sea” • Usually formed between metals.

16. Bond Properties ionic vs.covalent • formula unit (NaCl) molecule (CO2) • hard & brittle pliable • soluble in H2O PC: soluble in H2O NPC: insoluble in H2O

17. Bond Properties ionic vs. covalent • high melting & low melting & boiling points boiling points • electrolytes non-electrolytes • electron Lewis structures transfer diagrams

18. Formula unit • ionic compounds are not found as single molecules. • the simplest ratio of cations to anions • Example: ZnCl2represents the simplest combination of zinc and chlorine: one Zn 2+ ion and two Cl- ions

19. WATER • Contains polar covalentbonds. • Called the “universal solvent” “Like dissolves like”.

20. Properties of Metals • Luster • High density • Good heat conductor • Good electrical conductor • High melting/boiling points • Malleable & Ductile

22. Calculating Bond Type

23. Very few bonds are purely one type. • The degree to which bonds are ionic or covalent can be estimated by comparing their electronegativities.

24. Remember: Electronegativity: the tendency of an atom to attract electrons to itself when chemically bonded to other atoms.

25. Remember: Electronegativity • Increases as you move from left to right across PT • Decreasesas you move from top to bottom on periodic table

26. To Calculate Bond Type: • Look up values on Table of Electronegativities. • Subtract. • Locate difference on Bond Type Chart.

27. Bond Type Chart % ionic character 100% 50% 0.5% 0.0% 4.0 1.7 0.3 0.0 ionic polar nonpolar covalent covalent

28. Examples • Ca and Br

29. O and O

30. H and S

32. IONIC BONDS

33. Ionization: formation of an ion by the loss or gain of one or more valence electrons.

34. Cations: positive ions formed by the loss of one or more valence electrons. • Metals tend to form cations.

35. 2) Anions: negative ions formed by the gainof one or more valence electrons • Non-metals tend to form anions.

36. Remember: Oxidation Numbers tell you how many valence electrons an atom will lose or gain to become stable.

37. We use Electron Transfer Diagramsto represent ionic bonds

38. To write Electron Transfer Diagrams: Use equation format: • Left of arrow: show electron dot diagrams • Right of arrow: show ions formed and coefficients to give proper ion ratio

39. Stop and do the Examples

40. COVALENT BONDS

41. Remember: • Covalent bonds are usually formed between non-metals

42. We use Lewis Structuresto represent covalentbonds

43. Understanding Lewis Structures • Element symbols represent the nuclei and core electrons.

44. Element symbols represent nuclei and core electrons • Dashes between symbols represent shared pairs of electrons • Dot pairs around the outside of a symbol represent unshared electrons

45. Single bond = 1 dash • Double bond = 2 dashes • Triple bond = 3 dashes

46. Central atom is the least electronegative atom.

47. Multiple bonds are most common with • Carbon • Nitrogen • Oxygen

48. To Write Lewis Structures • Count the total number of valence electrons.

49. Choose central atom (least electronegative)

50. Connect remaining atoms to central atom with single dash.