Chapter 2 Sections 9-10 Review :  Nomenclature

1. Acids and Bases (section 2.9)
   1. Acids - molecules that can lose H⁺ (proton donors) in water making H₃O⁺(aq) in water.  Note H⁺(aq) is really just short hand for H₃O⁺(aq).  Also note that (aq) means dissolved and surrounded by water. If there are H⁺(aq) ions called hydronium ions present, then the solution is acidic. 
      1. Know these acids:  HCl, HNO₃ and H₂SO₄.  Hydrochloric, nitric and sulfuric. 
      2. Reactions looks like this:  HCl(g) +H₂O(l) g  H₃O⁺(aq) + Cl^-(aq)
      3. Remember people that take acid are losers - so acids LOSE H⁺
   2. Bases - molecules than can gain H⁺ (proton acceptors) and/or make OH^- in water.  If there are OH^-(aq) ions present called hydroxide ions, then the solution is basic. 
      1. Know these bases:  NaOH, KOH, Ba(OH)₂ which are all solids
      2. Reactions looks like this when put in water:  NaOH(s) g  Na⁺(aq) + OH^-(aq)  
      3. Try Problem 2.14
   3. Why are these dissolving in water in the first place?  Well because the sum of all the ion-dipole forces created by having the ions surrounded by water are greater than the original ionic or covalent bond. Not all acids and bases dissolve in water so easily!  You'll learn a lot more in second semester.           

       

2. Nomenclature (section 2.10)   SUPER IMPORTANT
   1. Ionic Compounds
      1. Why do they form in the first place?  Well remember that each atom's goal is to be s²p⁶ and to get there they must lose or gain electrons (except the Noble Gases) and then they are cations or anions with an oxidation state.  Review it here! IMPORTANT
      2. You MUST be able to put ionic formulas together so that the charges sum to zero.  Practice is below this outline. DO IT!
      3. name of metal cation + name of nonmetal anion + ide
      4. Examples
         1. NH₄Cl - ammonium chloride
         2. CaCO₃ - calcium carbonate
         3. NaOH - sodium hydroxide
         4. MgO - magnesium oxide
         5. Li₃P - lithium phosphide
         6. CaBr₂ calcium bromide
         7. Na₂S - sodium sulfide
         8. Mg₃(PO₄)₂ - magnesium phosphate
         9. Na₂SO₄ - sodium sulfate
   2. Transition metal compounds
      1. name of transition metal, (charge in Roman numerals) + name of nonmetal + ide
      2. Why the Roman numeral?  Well the oxidation state of most transition metals varies - so we have to indicate what it is. 
      3. Examples
         1. Fe₂O₃ - iron(III) oxide
         2. AgCl - silver(I) chloride
         3. CuS - copper(II) sulfide
         4. Co(NO₃)₃ - cobalt(III) nitrate
         5. FeSO₄ - iron(II) sulfate
         6. Zn(OH)₂ - zinc(II) hydroxide
         7. iron(III) flouride - FeF₃
         8. iron(II) phosphate - Fe₃(PO₄)₂
      4. Look at examples 2.9 and 2.10.  Try problems 2.15 - 2.17.
      5. Remember you are supposed to know all these atom's names: H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Cr, Mn, Fe, Co, Ni, Cu, Zn, Br, Kr, Sr, Ag, Sn, I, Ba, Au, Hg, Pb and U  IMPORTANT
      6. Hey - Ag is always +1, Zinc and Cadmium are always +2 so you don't need to use a Roman Numeral with them:  AgCl is silver chloride, Zinc fluoride is ZnF₂, Ag₂S is silver sulfide
   3. Covalent Compounds
      1. prefix name of least electronegative nonmetal + prefix name of most electronegative nonmetal
      2. prefixes are mono, di, tri, tetra, penta, hexa, hepta, octa (mono is ommitted from the first element)
      3. Examples
         1. CO₂ - carbon dioxide
         2. CO - carbon monoxide
         3. PCl₅ - phosphorus pentachloride
         4. SF₄ - sulfur tetraflouride
         5. P₂O₅ - diphosphorus pentoxide
         6. SiO₂ - silicon dioxide
         7. H₂O₂ - dihydrogen dioxide
         8. nitrogen monoxide - NO
         9. carbon tetrachloride - CCl₄
         10. nitrogen trihydride - NH₃
         11. iodine heptaflouride - IF₇
      4. Look at example 2.11.  Try problems 2.18 and 2.19
   4. Ionic compounds with polyatomic ions
      1. treat polyatomic ions as groups that always stay together
      2. never change the name of the polyatomic ion, otherwise name as ionic compounds
      3. Memorize these:  NH₄⁺,  CH₃COO^-,  CN^-,  HCO₃^-,  OH^-,  MnO₄^-,  NO₂^-,  NO₃^-,  CO₃^2-,  SO₄^2-,  SO₃^2-,  PO₄^3-  Find their names in Table 2.3.  IMPORTANT
      4. NaOH is sodium hydroxide, Ca(NO₃)₂ is calcium nitrate, MgSO₄ is magnesium sulfate
      5. ammonium sulfide is (NH₄)₂ S, calcium carbonate is CaCO₃, magnesium phosphate is Mg₃(PO₄)₂
   5. Oxoanions - different numbers of oxygens
      1. Know hypochlorite, chlorite, chlorate, perchlorate
      2. Try example 2.12 and 2.13.  Try problems 2.20, 2.21, 2.22.
   6. Oxoacids
      1. Know hypochlorous, chlorous, chloric, and perchloric acids
      2. Know nitrous and nitric acids
      3. Learn them in Table 2.4.  Try example 2.14 and problem 2.23

Final Notes - At this point in time you should be able to:

1. Given any covalent molecule, you should be able to draw the Lewis dot structure, tell the molecular shape, electron geometry, hybridization and bond angles, name the molecule, tell if the bonds are polar or not, and tell if the entire molecule is polar or not. 
2. Given any group 1 or 2 metal, and a nonmetal, you should be able to predict the formula of the resulting ionic compound, know the oxidation states on the ions, and name the compound. 
3. Be able to name a transition metal compound.

Practice!!!  Ya Hoo!!!

1. Predict the most common oxidation state for the ions the following elements form:  Li, S, Cs, Al, Br, N, Mg, K, P, Bi, F, Ne, Ba, Pb, As, Te.
2. How many electrons total do the noble gases have?
3. What will the following atoms do in order to be like a noble gas?  Ga, F, I, Be, K
4. Predict the ionic compound formulas for the following pairs of elements.  (Assume they can make an ionic compound)  Li and Cl, K and S, Na and P, Mg and Br, Ca and S, Ba and N, Al and F, Ga and O, In and P, Pb and F, Pb and O, Bi and Cl, Bi and S, Po and F
5. What makes the noble gases so stable?
6. What is the difference between an ionic and covalent bond?
7. If an atom has more protons than electrons what is it called?
8. Ionic compounds are held together by ionic bonds which are basically ____________ attractions between ions. 
9. Name all the elements that form covalent bonds with themselves and give their formulas.
10. Why does hydrogen occur naturally as H₂?
11. Name the following ionic compounds:  KI, AlBr₃, MgSO₄, Na₃PO₄, LiCH₃COOH, Fe(CN)₂, Ni(HCO₃)₃, KOH
12. Give the chemical formulas for the following ionic compounds:  calcium nitrate, lithium iodide, sodium carbonate, ammonium nitride
13. Name the following covalent compounds:  CO, PI₅, IF₇, SiF₄
14. Write chemical formulas for the following covalent compounds:  nitrogen trihydride, silicon tetrabromide, sulfur dihydride
15. Name the following compounds:  KNO₃, N₂O₅  NaBr, PH₃, MgS, Ca(OH)₂, CS₂, NH₄I, LiF, (NH₄)₂CO₃, PCl₅
16. Write chemical formulas for the following:  dihydrogen disulfide, lithium nitride, calcium phosphide, nitrogen tribromide, magnesium nitride, sodium iodide, sulfur tetrahydride, potassium phosphate, silver permanganate, zinc nitrite, sodium sulfite

Answers:

1. Li is +1, S is -2, Cs is +1, Al is +3, Br is -1, N is -3, Mg is +2, K is +1, P is -3, Bi is +5, F is -1, Ne does not make ions, Ba is +2, Pb is +4, As is +5 or -3, Te is +6 or -2.
2. 2, 10, 18, 36, 54 and 86. 
3. Ga will lose 3 electrons to become Ga⁺³, F will gain one electron to become F^-, I will gain one electron to become I^-, Be will lose two electrons to become Be⁺², K will lose one electron to become K⁺.
4. Li and Cl make LiCl, K and S make K₂S, Na and P make Na₃P, Mg and Br make MgBr₂, Ca and S make CaS, Ba and N make Ba₃N₂ (Ba is +2 and N is -3 so we need three Ba's to make +6 and we need two N's to make -6 because +6-6 = 0), Al and F make AlF₃, Ga and O make Ga₂O₃, In and P make InP, Pb and F make PbF₄, Pb and O make PbO₂, Bi and Cl make BiCl₅, Bi and S make Bi₂S₅ (S is -2 while Bi is either +5 or -3.  Since S is already negative, Bi must be positive.  We need five S's to make -10 and we need 2 Bi's to make +10), Po and F make PoF₆ (Here Po will be +6 instead of -2 since it is bonding with F which is negative)
5. The have full outer s and p subshells = 8 electrons.  Except helium which just has the full 1s with 2 electrons.
6. An ionic bond involves the transfer of electrons from a metal to a nonmetal.  A covalent bond involves the sharing of electrons between 2 or more nonmetals.
7. cation
8. electrostatic
9. hydrogen occurs naturally as H₂, Nitrogen occurs naturally as N₂, Oxygen as O₂, Fluorine as F₂, Chlorine as Cl_2, Bromine as Br₂, Iodine as I₂, and just for fun Phosphorus as P₄, Sulfur as S₈ and Selenium as Se₈
10. Each hydrogen atom has only one electron so it is not stable.  H wants to have 2 electrons to be like helium.  If two H atoms come together they can share their electrons so that each H then has 2 electrons.  So then each H is stable and like a noble gas.  H : H
11. potassium iodide, aluminum bromide, magnesium sulfate, sodium phosphate, lithium acetate, iron(II) cyanide, nickel(III) bicarbonate, potassium hydroxide
12. Ca(NO₃)₂, LiI, Na₂CO₃, (NH₄)₃N
13. carbon monoxide, phosphorus pentaiodide, iodine heptafluoride, silicon tetrafluoride
14. NH₃, SiBr₄, SH₂
15. potassium nitrate, dinitrogen pentoxide, sodium bromide, phosphorus trihydride, magnesium sulfide, calcium hydroxide, carbon disulfide, ammonium iodide, lithium flouride, ammonium carbonate, phosphorus pentachloride
16. H₂S₂, Li₃N, Ca₃P₂, NBr₃, Mg₃N₂, NaI, SH₄, K₃PO₄, AgMnO₄, Zn(NO₂)₂, Na₂SO₃

We have now finished Chapter 2 completely.  Study Hard.

Recommended problems:  74, 78, 82, 84, 86, 104, 110