Essential Question:

2. What are properties of ionic bonds and ionic compounds? Essential Question:

3. SC3 Students will use the modern atomic theory to explain the characteristics of atoms. • e. Compare and contrast types of chemical bonds (i.e. ionic, covalent). GP Standards:

4. I. Ion Formation octet rule: atoms tend to gain, lose, or share electrons in order to acquire eight valence electrons chemical bond cation anion Ions are formed when atoms gain or losevalenceelectrons to achieve a stableoctet electron configuration. Section 7-1

5. Valence Electrons and Chemical Bonds • A chemical bond is the force that holds two atoms together. • Chemical bonds form by the attraction between the positive nucleus of one atom and the negative electrons of another atom. • Atoms try to form an octet by gaining or losing valence electrons. Section 7-1

6. Positive Ion Formation • A positively charged ion is called a cation. • Metals usually form positive ions. Section 7-1

7. Positive Ion Formation (cont.) • Metals are reactive because they lose valence electrons easily. • Transition metals commonly form 2+ or 3+ ions, but they can form greater than 3+ ions because they have lots of electrons in their outer energy level or close to the outer energy level. • Other relatively stable electron arrangements are referred to as pseudo-noble gas configurations. Section 7-1

8. Negative Ion Formation • An anionis a negatively charged ion. • Nonmetals usually form anions. Section 7-1

9. Negative Ion Formation (cont.) • Nonmetal ions gain the number of electrons required to fill an octet. • Some nonmetals can gain or lose electrons to complete an octet. Section 7-1

10. Group 1 = +1 Group 2 = +2 Group 13 = +3 Group 14 = +4 Groups 3-12 are generally +1 or +2 Group 15 = -3 Group 16 = -2 Group 17 = -1 Group 18 = 0 How do you know what ion will form? (MEMORIZE THESE NUMBERS)

11. A B C D Section 7.1 Assessment Oxygen gains two electrons to form what kind of ion? A.1– anion B.2– anion C.1+ cation D.2+ cation Section 7-1

12. A B C D Section 7.1 Assessment Elements with a full octet have which configuration? A.ionic configuration B.halogen configuration C.noble gas configuration D.transition metal configuration Section 7-1

13. II. Ionic Bonds and Ionic Compounds (cont.) ionic bondionic compound crystal latticeelectrolyte lattice energy Oppositely charged ions attract each other, forming electrically neutral ionic compounds. Section 7-2

14. Formation of an Ionic Bond • The force that holds oppositely charged particles together in an ionic compound is called an ionic bond. • Compounds that contain ionic bonds are called ionic compounds. • Binary ionic compounds contain only 2 different elements—a metallic cation and a nonmetallic anion. Section 7-2

15. Properties of Ionic Compounds • Positive and negative ions exist in a ratio determined by the number of electrons transferred from the metal atom to the non-metal atom. • The repeating pattern of particle packing in an ionic compound is called an ionic crystal. The formula for salt is NaCl—this means that the RATIO of Na atoms to Cl atoms is 1:1 Section 7-2

16. Properties of Ionic Compounds (cont.) • The strong attractions among the positive and negative ions result in the formation of the crystal lattice (3-D geometric arrangement of particles). • varies w/ the size, shape, and # of ions bonded. Section 7-2

17. Properties of Ionic Compounds (cont.) • The high melting & boiling points are an indication of the strength of the bond. • The higher the melting or boiling point, the stronger the bond Section 7-2

18. Hard, rigid, brittle solids due to the strong attractive force holding the ions in place. Properties of Ionic Compounds (cont.) • In a solid, ions are locked into position and electrons cannot flow freely—solid ions are poorconductors of electricity.

19. Properties of Ionic Compounds (cont.) • Liquid ions or ions in aqueous solution have electrons that are free to move, so they conduct electricity easily. • An ion in aqueous solution that conducts electricity is an electrolyte. Salt water is an example. Section 7-2

20. Energy and the Ionic Bond • Reactions that absorb energy are endothermic. • Reactions that release energy are exothermic. • The energy required to separate 1 mol of ions in an ionic compound is referred to as the lattice energy. Section 7-2

21. Energy and the Ionic Bond (cont.) • Lattice energy is directly related to the size of the ions that are bonded. • Force of attraction increases as the distance between charges decreases. So, the smaller the ion, the greater the attraction. • Lattice energy is also affected by the charge on the ions bonding. The higher the charge on the ions, the higher the lattice energy. Section 7-2

22. Energy and the Ionic Bond (cont.) Section 7-2

23. A B C D Section 7.2 Assessment Why are solid ionic compounds poor conductors of electricity? A.They are non-metals. B.They are electrolytes. C.They have electrons that cannot flow freely. D.Solids do not conduct electricity. Section 7-2

24. A B C D Section 7.2 Assessment What is the charge holding two ions together? A.covalent bond B.pseudo-noble gas bond C.crystal lattice bond D.ionic bond Section 7-2

25. III. Metallic Bonds and the Properties of Metals (cont.) electron sea modeldelocalized electron metallic bondalloy Metals form crystal lattices and can be modeled as cations surrounded by a “sea” of freely moving valence electrons. Section 7-4

26. Metallic Bonds and the Properties of Metals • Metals are not ionic but share properties with ionic compounds. • Metals also form lattices in the solid state where 8 to 12 other atoms closely surround each metal atom. • Within the crowded lattice, the outer energy levels of metal atoms overlap. Section 7-4

27. Metallic Bonds and the Properties of Metals • The electron sea model says that all metal atoms in a metallic solid contribute their valence electrons to form a "sea" of electrons. • The electrons are free to move around and are referred to as delocalized electrons, forming a metallic cation. Section 7-4

28. Metallic Bonds and the Properties of Metals • A metallic bondis the attraction of an metallic cation for delocalized electrons. Electrons don’t belong to one specific atom, but roam amongst all the atoms. Section 7-4

29. Metallic Bonds and the Properties of Metals • Boiling points are much more extreme than melting points because of the energy required to separate atoms from the groups of cations and electrons. Section 7-4

30. Metallic Bonds and the Properties of Metals (cont.) • Metals are malleable because they can be hammered into sheets. • Metals are ductile because they can be drawn into wires. • This is because the mobile particles can be pushed or pulled past each other. Section 7-4

31. Metallic Bonds and the Properties of Metals (cont.) • Mobile electrons around cations make metals good conductors of electricity and heat. Why? • As the number of delocalized electrons increases, so does hardness and strength. Section 7-4

32. Metal Alloys • An alloy is a mixture of elements that has metallic properties. • The properties of alloys differ from the elements they contain. Adding another element to iron (to create steel) increases the strength of the mixture. Section 7-4

33. Section 7-4

34. A B C D Section 7.4 Assessment The attraction of a metallic cation and delocalized electrons forms what kind of bond? A.ionic B.covalent C.diatomic D.metallic Section 7-4

35. A B C D Section 7.4 Assessment Which property of metals allows them to be easily drawn into wires? A.malleability B.ductility C.conductivity D.durability Section 7-4

36. IV. The Covalent Bond (cont.) covalent bond molecule Lewis structure Atoms gain stability when they share electrons and form covalent bonds. Section 8-1

37. Why do atoms bond? • Another way atoms can gain stability is to share electrons rather than gaining or losing them. This results in a covalent bondforming amolecule. • Diatomic molecules (i.e. H2 & F2) exist because two-atom molecules are more stable than single atoms. Section 8-1

38. Why do atoms bond? (cont.) • By sharing electrons, both atoms achieve 8 electrons in their outer energy level and have a full octet. Section 8-1

39. Single Covalent Bonds • When only one pair of electrons is shared, the result is a single covalent bond. Section 8-1

40. Multiple Covalent Bonds • Double bonds form when 2 pairs of electrons are shared between 2 atoms. • Triple bonds form when 3 pairs of electrons are shared between 2 atoms. Section 8-1

41. Single Covalent Bonds (cont.) • In a Lewis structure dots are used to represent unshared electrons and a line is used to symbolize a single covalent bond. • The halogens form single covalent bonds with atoms of other non-metals. Section 8-1

42. Single Covalent Bonds • Atoms in group 15 form 3 single covalent bonds, such as in ammonia. • Atoms in group 16 can share 2 electrons and form two covalent bonds. Section 8-1

43. Single Covalent Bonds (cont.) • Atoms of group 14 form 4 single covalent bonds, such as in methane. Section 8-1

44. The Strength of Covalent Bonds • As the distance between 2 nuclei increases, the strength of the bond decreases. • The shorter the bond length, the higher the energy required to break it. Section 8-1

45. Structural Formulas • A structural formulauses letter symbols & bonds to show relative positions of atoms. Section 8-3

46. Drawing Lewis Structures Determine which atom goes in the center. (H never, C always, or lowest electroneg.) Draw the dot diagram for the central atom. Using a line, add the other atoms to the central atom on sides that DON’T already have 2 electrons. Erase the dot on sides of the central atom you added other atoms to. Show the valence electrons on the non-central atoms. Section 8-3

47. Exceptions to the Octet Rule • Some molecules do not obey the octet rule and form stable compounds with less than 8 valence electrons. Section 8-3

48. Exceptions to the Octet Rule (cont.) • Other compounds have central atoms with more than 8 valence electrons, called an expanded octet. • Elements in period 3 or higher have a d-orbital and can form more than four covalent bonds. Section 8-3

49. Draw Lewis Structures for the following: H2O PCl3 HBr CH4 OF2 SiH4 H2S CCl4

50. NaCl • Cu • PCl3 • KI • SrO 6) H2O 7) Mn 8) OF2 9) MgO 10) IBr Categorize the following compounds as metallic, ionic, or covalent (solely based on the location of the elements on the PT)