Welcome Back!!

1. Welcome Back!! • This semester… • January- Bonding and Naming Ionic and Covalent Compounds • Jan/February- The Mole • After Feb. break- Reactions • March/April- Stoichiometry • April- Chemical Solutions • May- Acids and Bases, Final Exam Review

2. Chemical Bondingand Nomenclature

3. Valence electrons • The electrons in the highest energy level of an atom. • s and p electrons* • *d electrons are never valence electrons • Total # of s and p electrons in a “FULL” energy level = 8 (THE OCTET RULE)

5. Lewis Structures • We include only valence electrons Water How many valence electrons?

6. LEWIS DOT STRUCTURES FOR THE FIRST 20 ELEMENTS

7. Bonding • Chemical bond: • A force that holds groups of two or more atoms together and makes them function as a unit • Fundamental types of bonds • Ionic • Covalent • Metallic • Bonds form in order to achieve a full outer shell (8) of electrons to: • Decrease potential energy (PE) • Increase stability

8. Ionic bonding • Described as a TRANSFER OF ELECTRONS • Takes place between metals and nonmetals. Nonmetals “want” electrons and metals “want” to give electrons away. • Atoms become ions by gaining or losing electrons • Oppositely charged ions are attracted to one another: this attraction is an “ionic bond.” • Results in an ionic COMPOUND

9. Metals and Nonmetals Create Ionic Bonds

10. Ions • Cation • Positively charged • Atom has LOST one or more electrons • Metals tend to become cations • Anion • Negatively charged • atom has GAINED one or more electrons • Nonmetals tend to become anions

11. Types of ions • Monatomic ion • Cation or anion that consists of a single type of element • Ex: Li+, Br- • Polyatomic ion • 2 or more elements that are combined, but act as a single ion (or particle) • Ex: CO32- or NH4+

12. Formula units • Ionic compounds are represented by formula units: similar to “molecule” except ionic compounds exist as a 3D network of multiple formula units because of the electrostatic attraction of opposite charges. • The lowest whole number ratio of ions in an ionic compound is a formula unit. NaCl CaO MgCl2 Al2O3

13. Types of ionic compounds COMPOUND more than 2 Elements  Means there is a Polyatomic ion present 2 elements Binary Compound Ternary Compound NaCl NaNO3

14. Rule of Zero Charge • The charges on the metal cations (+) and nonmetal anions (-) add up to zero. • Ex: Na+ and Cl- add up to an overall charge of zero, so the formula unit is NaCl • Ex: Ca2+ and O2- add up to an overall charge of zero, so the formula unit is CaO

15. Writing Binary Ionic Compounds • RULES: ****Cation always listed first**** • Drop Charges • Criss Cross Apple Sauce • Reduce (simplify) subscripts if needed • Na+ and Cl-  NaCl

16. What if the charges are opposite, but not equal??? • Criss-cross the charge values to subscripts • Ex: Al3+ + O2- is represented as Al2O3 • Ex: Mg2+ + Cl- is represented as MgCl2 • But why? • Because the overall charge must be zero • All atoms “want” a full valence

17. Naming Binary ionic compounds CaO Fe2O3 • List name of cation • If metal cation is not a transition element, give it the same name as the element and move on to anion. • If cation can have more than one charge (i.e. most transition metals), use roman numerals in parentheses to specify charge. • For Example: • Iron has 2 common oxidation states (charges): Fe2+ and Fe3+ • Fe2+ would be iron (II) • Fe3+ Would be iron (III)

18. Naming Binary ionic compounds • Name the cation first • Name the anion last; for binary compounds, it always ends in –ide. • Examples: CaO would be calcium oxide Fe2O3 is iron (III) oxide NaCl is sodium chloride AgCl is silver chloride

19. Common anion endings • P- phosphide F-fluoride • O-oxide Br- bromide • S- sulfide I- iodide • Cl-chloride C- carbide • N-nitride

20. Some common polyatomic ions

21. Polyatomic Ions- QUIZ TUESDAY!!

22. Naming polyatomic anions: these just have to be memorized! • With anions that end in oxygen, if you know one , you can figure out the other! # of oxygensnaming convention 1 less (PO33-) ends in –ite(phosphite) 1 more (PO43-) ends in –ate(phosphate) Notice that within one group of polyatomic ions, the charges are the same!

23. Writing formulas for ternary compounds • If more than 1 of a polyatomic ion is present, write parentheses around any polyatomic ion and add subscripts to the outside of the parentheses. • Never change the subscript of a polyatomic ion (it will change the composition of the ion): • Ca(OH)2, Ca3(PO4)2 • Include the subscript inside the parentheses

24. Naming ternary ionic compounds • For a monatomic cation and polyatomic anion, write both names together; do not change them in any way • Ex: Na2SO4 is sodium sulfate • For a polyatomic cation with a monatomic anion, write the polyatomic ion’s name first, then the monatomic root with the –ide ending • Ex: NH4Cl is ammonium chloride

26. Metallic Bonding • Metals form lattices but do not lose, gain, or share e- • Outer e- shells overlap = electron sea model • e- are free to move = delocalized e- making a metallic cation • Metallic bond: attraction of a metallic cation for delocalized electron • Properties of metals: melting point varies greatly, malleable, ductile, good conductors

27. Sea of electrons

28. Ionic Metallic • Electron sea model – delocalized electrons attracted to metal cations • Malleable, ductile, lustrous, colorful • Conducts electricity • Overall highest MP of 3 types • Oppositely charged ions attracted to each other (cations and anions) • Metals with non-metals • Brittle, dull solids • Only conducts electricity as electrolytes • Generally, medium MP of 3 types All bonding involves valence electrons Both involve metals

29. Covalent bonding

30. Covalent bonds • Occur between atoms that are “sharing” electrons • Form covalent compounds

31. QUIZ: 1-5, NAME THE COMPOUND…6-10,WRITE THE FORMULA 1) SrO 2) NaClO3 3) AlF3 4) NH4OH 5) Fe3(PO3)2 6) barium nitrite 7) potassium chloride 8) manganese (II) iodide 9) copper (II) acetate 10) sodium carbonate