1. Chemical Bonding Chem I: Chapters 4, 5, 9 Chem IH: 7, 8, 9

2. Unit Objectives To be able to: Explain why some elements react (form bonds.) Explain why some elements react (form bonds.) Explain the difference between an ionic & a covalent bond. Explain the difference between an ionic & a covalent bond. Demonstrate e- reconfiguration when a simple compound is formed. Demonstrate e- reconfiguration when a simple compound is formed. Explain how interparticle forces affect the properties of ionic & covalent compounds. Explain how interparticle forces affect the properties of ionic & covalent compounds.

3. What is Happening?  Watch the videos.  What is happening in all of these reactions ? Hydrogen balloon burning 2H 2 + O 2  2H 2 O sodium metal & chlorine Na + Cl  NaCl sodium iodide & mercury (II) chloride 2NaI + HgCl 2  2NaCl + HgI 2

4. The Chemical Bond  The force that holds two atoms together  Occurs using valence electrons

5. Lewis Dot Diagram  Definition: System of showing the valence e-s for an atom  Help us predict bonds

6. Lewis Dot Diagram, cont. 1) Element symbol in middle 2) Valence e-s represented by dots  Place v.e.-s around the element symbol in 4 locations  Above  Below  Left  Right  Don’t pair up any, until 1 in each location.

7. Lewis Dot Diagram, cont.  Practice:  Together: let’s do LDD for Hydrogen, helium,  You practice: lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine, neon.

8. Lewis Dot Diagram of Main Group Elements

9. The Noble Gases  All occur in nature  BUT no compounds containing them are ever found in nature  Common denominator is that they all have 8 valence e-s. (*except He, which has 2)  This is what makes them stable/unreactive  This is called “Noble Gas Configuration” (NGC)

10. All “Other” Elements  All other elements “want” NGC, so they can be stable.

11. Atoms can Get NGC in one of two ways 1. Lose/gain e-s (IONIC bond)  This occurs when the 2 atoms in the bond pull on their e-s w/different strengths  Big EN difference  Typically occurs btwn a _____ & a _____.

12. Atoms can Get NGC in one of two ways (cont.) 2. Share e-s (COVALENT bond)  This occurs when the 2 atoms in the bond pull on their e-s w/the same (or similar) strength  Small EN difference  Typically occurs between 2 ___________.

13. 3 “Classes” of Elements (Review- Do Not Copy)  Metals = pink Conductors? Yes! Hold e-s? Loosely How many V.E’s? 1-3  Nonmetals = green Conductors? No! Hold e-s? Tightly How many V.E.’s? 5-7  Metalloids = blue “In between”

14. Ions & the Octet Rule  Ions are charged particles that have lost or gained e-s to satisfy the octet rule (8 e-s)  They will typically form based on what requires the smallest gain or loss of e-s to complete an octet.  Ex: Na  Na + + e -  Ex: Cl + e -  Cl -

15. Ions & the Octet Rule, cont.  This also means that atoms will typically obtain the NGC of the nearest *noble gas  “Nearest” means the one with the closest atomic number  Ex: Na obtains the NGC of Neon, not Argon b/c Neon’s atomic # is 10, while Argon’s is 18

16. Formation of Ions  The charge on an ion is called its “oxidation number”

17. Formation of Ions, cont.  Metals tend to lose electrons  Have + oxidation number  Called “cations”  Metal ions are given the element name  Ex: Na+ = “sodium ion”

18. Formation of Ions, cont.  Nonmetals tend to gain electrons  Have – oxidation number  Called “anions”  Nonmetal ions are given the element name w/an “-ide” ending  Ex: O 2- = “oxide ion”

19. TO DO: Label Your Periodic Table  Label Groups 1, 2, 3, 16, 17 w/the oxidation number of each group. Group #Oxidation # 1+1 2+2 3+3 16-2 17-1

20. Practice  Together, write the arrow orbital diagram for a sodium atom.  Now draw the arrow orbital diagram for a sodium ion.  Identify the noble gas that has the same notation as a sodium ion.  Write the symbol for a sodium ion

21. Formation of an ionic bond  When e-s are lost by 1 atom they are gained by another. Therefore, cations & anions are formed at the same time.  Ex: Na & Cl Ex: Na & Cl

22. Formation of an ionic bond, cont.  The attraction between + and - ions is what forms bond. Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1 Cl -1 Na +1

23. Practice 1. Draw Lewis Dot Diagrams for sodium and chlorine ATOMS. 2. Show how e-s rearrange to form sodium & chlorine IONS. 3. Write the formula (cation, then anion) 4. Name this compound (cation, then anion) 5. (BONUS: What NG is each ion like?)

24. Formation of an ionic bond, cont.  Note: you will not always have a 1:1 ratio of + to - ions.  Ex: MgI 2 -see next slide

25. Formation of MgI 2........ .Mg. +:I: → Mg. + + :I: - (are they happy?)........Mg. +:I: :I: → Mg 2+ + :I: - :I: -......

26. Formula Unit  (Def): Lowest ratio of ions in an ionic compound

27. Writing a Correct Ionic Formula 1.The symbol for the cation is always listed 1 st. 2. The symbol for the anion is always listed 2 nd. 3. Include subscripts that show how many cations and anions are present in a Formula Unit. 4. Ensure they are in the lowest ratio possible. Our Example: MgI 2

28. 2 Techniques to Determine Correct Ratio of Cations to Anions  1. Use Lewis Dot Diagrams and rearrange electrons until all ions are “happy.” (learned this already)  2. Use oxidation #s of cation & anion to calculate the ratio of ions to create a neutrally charged compound.

29. Technique #2: Using Oxidation #s to Write Formulas for Ionic Compounds Step 1. Identify the oxidation # on the cation & anion. LecturePLUS Timberlake29

30. Using Oxidation #s to Write Formulas for Ionic Compounds Step 2. A compound has NO CHARGE on it, so a formula must have equal numbers of + & - charges. (use the LCM) LecturePLUS Timberlake30

31. Using the LCM to Write Ionic Formulas (Practice)  Ex: Li & F  Ex: Li & O  Ex: Al & O LecturePLUS Timberlake31

32. Polyatomic Ions  (Def): Charged particles containing more than one element!  Ex: SO 3 2- = “sulfite”  See Table 9.3 on p 257 of text (ChemIH)  See Table 5.2 on p 159 text.  COPY THE APPROPRIATE TABLE ONTO YOUR P. TABLE OR INTO YOUR NOTES.

33. Writing Formulas for I.Cpds Containing Polyatomic Ions Step 1. Identify cation & anion Ex: aluminum sulfite  Aluminum = cation  Sulfite = anion LecturePLUS Timberlake33

34. Writing Formulas for I.Cpds Containing Polyatomic Ions, cont. Step 2: Determine the oxidation # on each ion. (found on a table for polyatomics) LecturePLUS Timberlake34 Ex:  aluminum = Al 3+  sulfite = SO 3 2-

35. Writing Formulas for I.Cpds Containing Polyatomic Ions, cont. Step 3: Write a balanced formula  If there is more than 1 of a polyatomic ion, use parentheses, then a subscript Ex: Al 2 (SO 3 ) 3 LecturePLUS Timberlake35

36. Cross-Over Method  “Shortcut” version of using oxidation #s to write formulas Step 1: Take the charge on the cation and use it as the subscript on the anion Step 2. Take the charge on the anion and use it as the subscript on the cation  Step 3: Reduce the subscripts, if necessary LecturePLUS Timberlake36

37. Oxidation #s of Transition Metals (Don’t Copy!)  The d sublevel e-s are so close to the actual valence e-s, they sometimes act like valence e-s!  See Table 5.4 on p 164 of text: Common Ions of Select Transition Metals  Ex: Fe 2+ & Fe 3+  Write the e- configuration LecturePLUS Timberlake37

38. Ionic Compound w/Transition Metals  Transition metals may have multiple oxidation #s!  You must be told the oxidation #  See Table 9.2 on p 255 of text (ChemIH)  See table Table 5.4 on p 164 of text (ChemI)

39. Writing Formulas for Ionic Cpds containing Transition Metals Step 1. Identify the cation & anion Step 1. Identify the cation & anion LecturePLUS Timberlake39

40. Writing Formulas for Ionic Cpds containing Transition Metals, cont. Step 2. Determine the oxidation #s on the cation & anion  The oxidation # of the cation is given in the name ex: lead (IV) = Pb 4+  The Roman numeral is the oxidation# LecturePLUS Timberlake40

41. Writing Formulas for Ionic Cpds containing Transition Metals, cont. Step 3. Write a neutrally charged formula. LecturePLUS Timberlake41

42. “Isoelectronic”  Term used to describe atoms/ions with the same e- configuration  Ex: F- and Ne  Both have 2 e-s in the 1 st energy level  Both have 8 electrons in the 2 nd energy level  Ex: He and H-  Both have 2 electrons in the exact same arrangement

43. Formulas  Tell us:  the elements that make up the compound  the # of atoms of each element in a unit of the compound  The smallest unit of an Ionic Compound is called a “Formula Unit”  The smallest unit of a Covalent compound is called a Formula Unit or “Molecule”

44. Quiz  Ionic bonds are formed by the ______ of electrons.  Covalent bonds are formed by the ________ of electrons.  CO is a/n ionic/covalent compound.  NaF is a/n ionic/covalent compound.