1. Ionic Nomenclature: 1. Binary: made-up of one monoatomic cation ion and one monoatomic anion. Metal (+) bond to Non-Metal (-) 2. For binary acids, the cation is ALWAYS H +. HBr 3. Writing ‘Formula Units’ -Write the symbol of the element. -Assign oxidation numbers. -Criss Cross - Simplify

2. 4. Naming Binary Ionic compounds - anion will end in –ide - first name is the name of the element. -second name is the name of the anion with the ending dropped and –ide added. Calcium bromide - All cations other than Group 1A, 2A and Aluminum will have and Roman numeral to indicate the oxidation number of the that metal. Iron (III) chloride

3. 5. Binary acids; Hydro-anion root-ic acid. Hydrochloric acid. 6. Ternary Compounds: ionic compound in which at least one of the ions is a polyatomic ion. 7. For a ternary acid, the cation ion is ALWAYS H +. H 2 SO 4. -Naming ternary acids follows –ate,-ic, -ite, -ous Sulfuric acid

4. 8. Writing the ‘Formula Unit’ (Same 4 steps) -write the atom or polyatomic ion -assign the oxidation number -criss cross -simplify (more than one poly, put ( ) ) 9. Naming ternary ionic compounds -first name is the name of the element (or Ammonia if NH 3 ). -second name is the name of the polyatomic ion. Cr(HSO 4 ) 3 Chromium bisulfate

5. 10. Bases: Ionic compounds in which the anion is hydroxide, OH - (except NH 3 ) Ba(OH) 2 Barium hydroxide

6. 1. Cations are positive ions. Always written first. 2. Anions are the negative ions. Written after the cation 3. Positive and negative ion pack in a regular pattern that balances the forces called an ionic crystal or crystal lattice. 4. Lattice energy defines the strength of the ionic bond. Large lattice energy = stronger bond 5. Form 3-D compounds 6. High melting and high boiling point. 7. Compound defined by hardness and brittle. 8. Almost always exothermic when bond is broken. Energy-in to form the bond and energy-out when the bond is broken

7. 9. Further separation on the periodic table, the MORE ionic characteristics. 10. Conductive current dissolved and molted state. Greater solubility = more ionic 11. Ionic compounds are electrolytes = they carry an electric current when dissolved. 12. Ionic compounds are referred to as formula units (FU). 13. Net charge of the compound must = zero 14. Greater ∆EN = greater ionic characteristic

8. UEQ What are the unique characteristics of a covalently bonded molecule? LEQ What is a covalent bond?

9. The joining of two or more elements through the sharing of valance electrons to form a molecule Purpose: To form a stable octet between the elements

10. 1. Covalent bonds are molecules (neutral) 2. Do not conduct electricity 3. Low melting and boiling points 4. Are less soluble in water and more soluble in non-aqueous solvents. 5. Can form multiple bonds. 6. Multiple bonds shorten bonding length and increases bonding strength. 7. Weak force of attraction between molecules.

11. 8. Elements close on the periodic tables tend to form covalent bonds. 9. Metallic bonds are a type of covalent bond 10. Hydrates bond covalently.

12.  Atoms the share electrons form molecules.  Molecules are expressed in ‘Molecular Formulas’  Naming follows binary ionic naming trends with prefixes to define the number of each atom in the molecule.  Hydrates follow binary naming with a prefix to identify the number of water molecules. ◦ Molecular formulas are NOT simplified.

13. Prefixes: mono- for one (used w/ the second name usually) di- for 2hexa- for 6 tri- for 3hepta- for 7 tetra- for 4octa- for 8 penta- for 5nano- for 9 deca- for 10

14. N2O5N2O5 Dinitrogen pentoxide CCl 4 Carbon tetrachloride Silicon dioxide SiO 2 Diphosphorous hexaflouride P2F6P2F6

15. CuO. 5H 2 O Copper (II) oxide pentahydrate NiCl 3. 2H 2 O Nickle (III) chloride dihydrate Chromium (II) sulfate heptahydrate CrSO 4. 7 H 2 O Barium sulfide trihydrate BaS. 3H 2 O

16.  What is polarity and how is it applied to molecular bonding?

17.  Non-Polar: equal sharing of the valance electrons.  Polar: unequal sharing of the valance electrons.

18. Polarity refers to the uneven sharing of valance electrons. Electronegativity refers to the pull on another atoms valance electrons. If the EN values are not the same then there is polarity (uneven pulling). If the EN values are equal, then there is no polarity (even pulling) called non-polar The closer the EN values the less polar.

19.  Shape of molecule is symmetrical.  Homonuclear molecules. ◦ Diatomic molecules :  N 2, O 2, F 2, Cl 2, Br 2, I 2, and H 2  Makes the #7

20.  Shape of the molecules is asymmetrical due to unequal sharing of the electrons.  Heteronuclear: one nuclear charge is stronger than another.  Polarity is the function of the change in electronegativity (  EN) ◦ Increase  EN, = more ionic chstc.

21. Rank these in decreasing covalent characteristics: H 2 O, N 2, NO 3 -, NaBr 2, CO 2 Solution: 1. Identify the electronegativity for each element in the molecule. 2. Less covalent > 1.7 > more covalent. N 2 > NO 3 - > CO 2 > H 2 O > NaBr 2

22.  Identify if the following is pure covalent, polar covalent or ionic: Use ∆EN O3O3 N2O5N2O5 RbCl N2N2 CCl 4 NH 3 Cl 2 SO 2 BaBr 2

23.  Pictorial representation of valance electrons.  Dot structure  Stick structure

24. BCl 3 CCl 4 NF 3

25.  Representative elements share electrons to take on a Noble gas electron configuration.  Each element in a molecule will follow the octet rule.

26.  Formula to determine the number of shared electrons:  N – A = S  N = # of electrons needed to form a Noble gas configuration.  A = # of electrons available in the valance.  S = # of electrons shared

27. CS 2 C2H4C2H4 CO 3 2- CHCl 3

28.  Sigma Bonds are the single electron overlap of the s orbital.  Forms single bonds, end-to-end.  Pi Bonds are the overlap of the s and p orbitals.  Forms double and triple bonds w/ s end- to-end and p side-to-side.

29. The joining together of two or more elements by the sharing of the valence electrons. The Lewis Structure: representation of the electrons in the valance of an atom. * Group A elements = the Group number is the number of valance electrons. For Example: N C PCl Show the Lewis structure for the following. PH 3 CCl 4 H 2 SSiH 4

30. Covalent bonds are formed when valance electrons between two elements are share either by an end-to-end overlap relationship or an side-to-side overlap relationship. 1. Sigma bond (σ ): single covalent bond with a end-to-end relationship BeCl 2 CH 4 AlF 3 NH 3 *sigma bonds are single bonds *bonding length is longer (Table 8.1) *bonding strength is less (Table 8.2)

31. pi bonds ( π ) : A side-to-side overlap of valance electrons. Forms multiple bonds. CO 2 O 2 N 2 *one sigma + one pi = double bond *one sigma + two pi = triple bond *double bond shorter bond length *double bond stronger bond *triple bond shortest bond length *triple bond strongest bond

32.  CO 2 N2N2  CS 2  HNO 3  NO 3 1-  Which are polar?  Which are non-polar?  Which are non-polar with polar bonds?

33.  Equally acceptable formulas.  HNO 3  NO 3 1-

34. MMost beryllium compounds MMost Group IIIA elements CCompound which require more than 8e- in the valance. CCompounds containing d or f transitional elements ‘‘S’ with an odd number of electrons

35. CCCl 4 CCO 2 NN 2 O 5 NN 2 O SS 3 O 5 NNF 3

36.  Follows the rules as ionic compounds except prefixes are used to note ‘how many’.  Table 8-3, page 248.

37.  VB Theory : Valance Bond Theory, orbital overlaps  VSEPR Theory: Valance Shell Electron Pair Repulsion Theory ◦ Electrons arrange to max the distance between electrons ◦ Bonding pairs v. Unshared pairs

38.  Linear (2)  Trigonal Planar  Tetrahedral  Trigonal Pyramidal  Angular (Bent)  Trigonal Bipyramidal  Octehedral

39.  Formula AB 2 w/ no unshared pairs  VSEPR: bonding angle of 180 o  VB: sp overlap  Forms a polar bond and a non-polar molecule.

40.  Formula of AB 3 and no unshared pairs.  VSEPR: bonding angle of 120 o  VB: SP 2 overlap  Polar bond w/ non-polar molecule

41.  Formula of AB 4 w/ no unshared pairs.  VSEPR: bonding angle of 109.5 o  VB: sp 3 overlap  Forms polar bonds and non-polar molecule.

42.  Formula AB 3 w/ one unshared pair on A.  General: subtract 2.5 o for each unshared pair.  VSEPR: bonding angle of 107 o  VB: sp 3 overlap w/ a polar bond and a polar molecule

43.  Formula AB 2 w/ 2 unshared pairs on A  VSEPR: bonding angle of 104.5 o  VB: sp 3 overlap w/ polar bonds and polar molecule.

44.  Formula AB w/ 3 unshared pairs  VSEPR: bonding angle of 102o  VB: sp 3 overlap w/ polar bonds and polar molecule.

45.  Formula AB 5 w/ no unshared pairs  VSEPR: bonding angles at 90 o, 120 o and 180 o  VB: sp 3 d overlap w/ polar bonds and polar molecule.

46.  Formula AB 6 w/ no unshared pairs  VSEPR: bonding angles of 90 o, 120 o and 180 o  VB: sp 3 d 2 overlap w/ polar bonds and molecule

47. Molecule VBVSEPR Shape NI 3 PH 3 CH 4 SF 6 H2SH2S PF 5 BeCl 2