AP Chemistry Mr. Niedenzu, Providence HS 09/16/08

AP Chemistry Mr. Niedenzu, Providence HS 04/04/18 AP Chemistry Name : Notes - Types of reactions Date :

I. Introduction to writing AP equations One section of the AP test is always on writing chemical equations. The AP student will be expected to do three reactions (recommended time allotment during the test is 15 min) in which he/she will have to predict the products and write the equation for the reaction along with balancing and answering a short answer question. Aspects of equation writing you should know are : - in all cases a reaction will take place (no "no reaction" answers) - there are no equilibrium ("double arrow") reactions - equations do need to be balanced! - the symbols (aq), (l), (g), and (s) are not necessary - all equations are to be written as net ionic equations (no spectator ions present) - all aqueous ions are written as aqueous ions (NaCl(aq) is "Na+ + Cl-") - molecular compounds, insoluble compounds, solids,, liquids, and suspended compounds are written in their molecular form (solid NaCl is "NaCl)". - Charges shown on ions must be correct. Each equations is worth five points, one for the reactants, two for the correct products, one for the balancing and one for the short answer. Points are lost for the presence of spectator ions and ionic species written as molecular.

Examples : a. Equal volumes of 0.1 M hydrobromic acid and 0.1 M sodium hydroxide are combined. + - Ans. H + OH  H2O b. Magnesium metal is added to a solution of zinc chloride. Ans. Mg + Zn2+  Mg2+ + Zn

A. Single replacement (or "single displacement") 1. metal displaces hydrogen from either water or an acid  metal + acid  salt and hydrogen gas + 2+  e.g. Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g) (Net : Mg + 2H  Mg + H2)  Group 1A or 2A metal and water  base and hydrogen gas

 e.g. 2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g) (Net : Li + 2H2O  Li+ + OH- + H2) 2. a more reactive metal replaces a less reactive metal ion from solution(see activity series) 2+ 2+  e.g. ZnCl2(aq) + Mg(s)  MgCl2(aq) + Zn(s) ( Net : Zn + Mg  Mg + Zn ) 3. a halogen replaces another halogen - -  e.g. MgBr2(aq) + F2(g)  MgF2(aq) + Br2(l) ( Net : 2 Br + F2  2 F + Br2 ) 4. Hydrides of alkali metals react with water to form hydroxides :

• e.g. LiH(s) + H2O(l)  LiOH + H2(g) (Net : LiH + H2O  Li+ + OH- + H2) B. Double replacement (metathesis) (must form precipitate, gas or molecular compound) 1. formation of an insoluble ionic compound (precipitate) -solubility rules will need to be memorized!! Solubility Rules

a. Most nitrates, chlorates and acetates are soluble. b. Most group 1 salts and ammonium salts are soluble. c. Most chloride, bromide and iodide salts are soluble, except for salts of Ag+, Pb2+ 2+ and Hg2 . d. Most sulfate salts are soluble. Exceptions include PbSO4, CaSO4, HgSO4 and BaSO4. e. Most hydroxide salts are slightly soluble (will form a precipitate). Ba(OH)2, Sr(OH)2 and Ca(OH) 2 are marginally soluble. f. Most sulfide, chromate, carbonate and phosphate salts are only slightly soluble. AP Chemistry Mr. Niedenzu, Providence HS 04/04/18

2+ 2-  e.g. BaCl2(aq) + MgSO4(aq)  BaSO4(s) + MgCl2(aq) ( Net :Ba + SO4  BaSO4(s)) 2. neutralization reaction between an acid and a base (forms water) a. strong acid/strong base : Common e.g. HCl + NaOH  H O + NaCl (net : H+ + OH-  H O) gases :  2 2 H S b. strong acid/weak base 2 CO2  e.g. HCl + NH3  NH4Cl (net : H+ + NH3  NH4+) SO2  e.g. HCl + CH3NH2  CH3NH3Cl (net : H+ + CH3NH2  CH3NH3+) NH3 c. weak acid/strong base

 e.g. HC2H3O2 + NaOH  H2O + C2H3O2- (net : HC2H3O2 + OH-  H2O + C2H3O2-) d. weak acids and weak bases + -  e.g. NH3 + HF  NH4 + F (net) 3. reactions with acids :

a. carbonates or bicarbonates and acids form a salt, water and CO2 2-  e.g. 2HCl + Na2CO3  2 NaCl + H2O + CO2 (net : H+ + CO3  H2O + CO2) b. sulfites and acids form a salt, water and SO2 2-  e.g. 2 HCl + Na2SO3  2 NaCl + H2O + SO2 (net : H+ + SO3  H2O + SO2) c. metallic sulfides and acids form H2S and a salt 2-  e.g. 2 HCl(aq) + Na2S(aq)  2 NaCl(aq) + H2S(g) (net : H+ + S  H2S ) d. metallic hydrides and acids form H2 and a salt

 e.g. HCl(aq) + LiH(aq)  LiCl(aq) + H2(g) (net : H+ + H-  H2) 4. strong acids with salts of weak acids (salts containing anions such as : OCl-, ClO2-, F-, NO2-, CN-, C2H3O2-) a. metallic acetates and acids form acetic acid and a salt -  e.g. HCl(aq) + NaC2H3O2(aq)  NaCl(aq) + HC2H3O2(aq) (net : H+ + C2H3O2  HC2H3O2) b. metallic nitrites and acids form nitrous acid and a salt -  e.g. HCl(aq) + NaNO2(aq)  HNO2(aq) + NaCl(aq) (net : H+ + NO2  HNO2) 5. reactions with bases : a. ammonium salts and soluble bases yield ammonia, water and a salt + -  e.g. NH4Cl + LiOH  NH3 + H2O + LiCl (net : NH4 + OH  NH3 + H2O) 6. Amides with water

 e.g. sodium amide and water produce sodium hydroxide and ammonia

 Net : NaNH2 + H2O  Na+ + OH- + NH3 C. Combination 1. metal and nonmetal form a binary ionic compound

 e.g. 2 Na + F2  2 NaF 2. two nonmetals form a binary covalent compound

 e.g. S + F2  SF2 3. metal or nonmetal combines with oxygen to form binary ionic or covalent compound

 e.g. 2 Mg + O2  2 MgO 4. nonmetallic element combines with a binary covalent compound  e.g. oxygen and nonmetallic oxide (sulfur dioxide + oxygen  sulfur trioxide)  e.g. nonmetallic halide and additional halogen ( chlorine trifluoride and fluorine  chlorine tetrafluoride) 5. Combination of two compounds a. metallic oxide and nonmetallic oxide form an ionic compound with a polyatomic ion

 e.g. Na2O + CO2  Na2CO3 b. metal oxides react with water to form hydroxides

 e.g. CaO + H2O  Ca(OH)2 c. nonmetal oxides react with water to form acids -  e.g. SO2 + H2O  H2SO3 (net : SO2 + H2O  H+ + HSO3 ) d. hydrates result when anhydrous compounds react with water to form hydrates AP Chemistry Mr. Niedenzu, Providence HS 04/04/18

 e.g. CuSO4 + 5 H2O  CuSO4·5H2O e. Lewis acid (e- pr acceptor) and Lewis base (e- pr. donor)

 e.g. boron trifluoride (e- deficient) and ammonia (lone pair on N) : BF3 + NH3  BF3NH3 D. Decomposition 1. thermal decomposition a. hydrogencarbonates (bicarbonates) yield carbonates, water and carbon dioxide (relatively low temps)

 e.g. NaHCO3  Na2CO3 + H2O + CO2 b. carbonates and hydrogen carbonates yield oxides and carbon dioxide (high temp.)

 e.g. CaCO3  CaO + CO2 c. ammonium carbonate decomposes into ammonia, water and carbon dioxide d. sulfites yield oxides and sulfur dioxide

 e.g. FeSO3  FeO + SO2 e. oxides, chlorates and perchlorates yield oxygen

 e.g. 2HgO  Hg + O2

 e.g. 2KClO3  2KCl + 3O2

 e.g. NaClO4  NaCl + 2O2 f. hydroxides, hydrates and some oxyacids release water

 e.g. Ca(OH)2  CaO + H2O

 e.g. CuSO4-5H2O  CuSO4 + 5 H2O

 e.g. H2SO3  H2O + SO2 2. Electrolysis - use of electricity to decompose compounds

 e.g. 2 NaCl  2 Na + Cl2 E. Combustion of hydrocarbons produces water and carbon dioxide

 e.g. 2 C6H14 + 19 O2  14 H2O + 12 CO2

F. Redox reactions - involve a transfer of electrons (something is oxidized and something else is reduced) and a change of oxidation state. Single replacement and combustion reactions are always redox reactions. Combination and decomposition reactions are sometimes redox reactions. If a reaction takes place in an acidic or basic solution it is most likely a redox reaction. Reactions involving elements in their natural

states (e.g. Al, Fe, Cl2, O2, etc.) are redox reactions.

1. Simple redox reactions : a. Hydrogen displacement

 e.g. Ca(s) + H2O  Ca(OH)2(aq) + H2(g) b. Metal displacement 2+ 2+  e.g. Zn(s) + CuCl2(aq)  ZnCl2(aq) + Cu(s) (Net : Zn + Cu  Zn + Cu) c. Halogen displacement

 e.g. Cl2(g) + MgBr2(aq)  MgCl2(aq) + Br2(l) (Net : Cl2 + Br-  Cl- + Br2) d. Combustion reactions (see above) e. Combination and decomposition reactions (see above) 2. Reactions involving transition metals with multiple oxidation states : • e.g. tin(II) ion with Fe(III) ion (net : Sn2+ + Fe3+  Sn4+ + Fe2+) 3. Free halogens in dilute basic solutions form hypohalite ions

• e.g. Cl2(g) + KOH(aq)  KClO(aq) + KCl(aq) + H2O(l) (Net : Cl2 + OH-  ClO- + Cl- +

H2O) 2- 4. Redox reactions involving oxyanions such as Cr2O7  2- 2+ 3+ 3+ e.g. 14H+(aq) + Cr2O7 (aq) + 6Fe (aq)  2Cr (aq) + 7 H2O(l) + 6Fe The only way to be able to predict these types of reactions is to become with familiar (i.e. memorize) common oxidizing and reducing agents. See table below.

Common Reducing Agents Products formed AP Chemistry Mr. Niedenzu, Providence HS 04/04/18 Halide ions Free halogen Free metals Metal ions

Sulfite ions or SO2 Sulfate ions Nitrite ion Nitrate ion Free halogens (dilute basic solution) Hypohalite ions Free halogens (concentrated basic solution) Halate ions Metallous ions (lower oxidation number) Metallic ions (higher oxidation number)

Common Oxidizing Agents Products formed 2+ MnO4- (acidic solution) Mn 2+ MnO2 (acidic solution) Mn

MnO4- (neutral or basic) MnO2(s) 2- 3+ Cr2O7 (acidic solution) Cr

HNO3 (concentrated) NO2(g)

HNO3 (dilute) NO(g)

H2SO4 (hot, concentrated) SO2(g) Metallic ions (high oxidation number) Metallous ions (lower oxidation number) Free halogens Halide ions

Na2O2 NaOH

HClO4 Cl- 2- C2O4 CO2(g)

H2O2 O2(g)

G. Organic reactions 1. Oxidation (combustion reactions) (See E. above) 2. Substitution reactions by halogens

• e.g. butane and bromine : C4H10 + Br2  C4H9Br + HBr 3. Esterification - organic acid and alcohol forms an ester and water

• e.g. acetic acid and ethanol : CH3COOH + C2H5OH  CH3COOC2H5 + H2O 4. Addition : A C=C bond is broken and atoms are attached to the carbon atoms

• e.g. butene and hydrogen gas form butane : C4H8 + H2  C4H10

• e.g. hexene and bromine : C6H12 + Br2  C6H12Br2 H. Complexation reactions (coordination chemistry) : (look for terms "excess" or "concentrated") These reactions usually involve transition metal ions bonding to ligands (Lewis bases - electron 3+ 3+ pair donors) to form complex ions. E.g. Cr(H2O)6 , where the transition metal ion is the Cr ions and the 6 water molecules are the ligands. They are bonded by a coordinate covalent bond which is formed by the sharing of the electron pairs donated by the water molecules.

Common ligands and their names : Ligand Name

H2O aquo

NH3 ammine Cl- chloro F- fluoro Br- bromo I- iodo CN- cyano OH- hydroxo - NO3 nitrato - NO2 nitro O2- oxo AP Chemistry Mr. Niedenzu, Providence HS 04/04/18 2- SO4 sulfato-

CH3NH2 methylamine

(CH3)2NH dimethylamine NO nitrosyl CO carbonyl Examples of cations : 3+ [Cr(H2O)6] - hexaaquochromium(III) 2+ [Cu(NH3) ] - tetraamminecopper + [Ag(NH3) ] - diamminesilver 3+ [Cr(NH3)5(H2O) ] - pentammineaquochromium(III)

Examples of anions : - [Cr(H2O)2Cl4] - diaquotetrachlorochromate(III)* [GaCl3(OH)]- - trichlorohydroxogallate(III) [Al(OH)4]- - tetrahydroxoaluminate

Examples of compounds :

Na[Al(OH)4] sodium tetrahydroxoaluminate

*ligands are arranged alphabetically (aquo before chloro) - ligands are named before the metal ion

The number of ligands attached to a central metal ion is called the coordination number. e.g. 3+ Cr(H2O)6 has a coordination number of 6. Two, four and six are the most common coordination numbers. The coordination number is often twice the oxidation number of the metallic ion. Ligands are joined one at a time and have a formation constant for each reaction that joins a single 3+ ligand. This means that it would take six reactions to form Cr(H2O)6 , each successive reaction joining another water molecule. Likewise, each step would have its own formation constant.

 Al2O3 + OH-  [Al(OH)4]-  2+ 2+ NH3 + Zn  [Zn(NH3)4]  Fe2+ + SCN-  [FeSCN] 2+

Keys to success in equation writing on the AP test :

1. know ion formulas and charges 2. know solubility rules 3. know common oxidizing agents 4. Identify reaction type - remember to add H+ and water to balance redox reactions (watch for term "acidified" -indicates redox rxn) - free elements indicate redox reactions 5. Remember amphoteric substances Zn2+ and Al3+ form hydroxides that dissolve in both acids and bases. 6. Leave out spectator ions

AP Test instructions : "Use appropriate ionic and molecular formulas to show the reactants and the products for the following, each of which occurs in aqueous solution except as indicated. Omit formulas for any ionic or molecular species that do not take part in the reaction. You need to balance. In all cases a reaction occurs."