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Reaction Types (Hints) I) Precipitation Reactions (double displacement reactions) Simple Rules for the Solubility of Salts in Water (Minor revisions to Zumdahl Table 4.1 are highlighted in bold print.) 1. Nitrate, acetate, chlorate salts are soluble. 2. Salts of alkali metals (Li, Na, K, Rb, Cs) and ammonium are soluble. 3. Chloride, bromide, iodide salts are soluble except those containing Ag+, Pb+2, Hg2+2. 4. Sulfate salts are soluble except BaSO4, PbSO4, Hg2SO4, CaSO4, SrSO4, Ag2SO4. 5. Hydroxides (and oxides) are insoluble except for alkali metals and ammonium. Hydroxides of barium, strontium, calcium are marginally soluble. 6. Sulfides are insoluble except for those with ammonium, alkali metals, alkali earth metals. 7. Carbonates, chromates, phosphates, silicates are insoluble except for those with ammonium and alkali metals. II) Acid-Base Reactions (double displacement reactions) Strong Acids – Some Weak Acids – Strong Bases - HCl, HBr, HI, HClO4, HClO3, HNO3, H2SO4 HC2H3O2, H3PO4, H2CO3, H2S, HF, HNO2 LiOH, NaOH, KOH, RbOH, CsOH Ba(OH)2, Sr(OH)2, Ca(OH)2 (soluble) (marginally soluble) The hydroxide ion (OH-) is such a strong base that it can be assumed to react completely with any weak (and strong) acid. Therefore, the weak acid is not dissociated on the reactant side of the equation, but is dissociated on the product side of the equation. Example: OH- (aq) + HF (aq) H2O (liq) + F- (aq) A similar argument can be made for weak bases in strong acids. The weak base is not dissociated on the reactant side of the equation, but is dissociated on the product side of the equation. Example: H+ (aq) + NH4OH (aq) H2O (liq) + NH4+ (aq) III) Redox Reactions (includes synthesis, decomposition, single displacement and combustion) Rules for Assigning Oxidation Numbers (from Zumdahl table 4.2) 1. Oxidation number of the atom of a free element is zero. 2. Oxidation number of a monoatomic ion equals its charge. 3. In compounds – oxygen has an oxidation number of -2 (except peroxides where it is -1). 4. In compounds – hydrogen has an oxidation number of +1. 5. In compounds – fluorine has an oxidation number of -1. 6. The sum of the oxidation states of an electrically neutral compound must be zero. Activity Series (reactivity decreases from beginning to end) Metals – Li K Ba Sr Na Ca Ni Sn Pb H2 Sb Cu Nonmetals - F Cl Br I Mg Hg Al Ag Zn Pd Cr Pt Fe Au Redox Reaction Prediction Important Oxidizers MnO4- (acid solution) MnO4- (base solution) MnO2 (acid solution) Cr2O7-2 (acid) CrO4-2 HNO3 (conc) HNO3 (dilute) H2SO4 (hot, conc) Metallic ions Free halogens HClO4 Na2O2 H2O2 Formed in Reaction Mn (II) MnO2 Mn (II) Cr (III) Cr (III) NO2 NO SO2 Metallous ions Halide ions ClOHO2 Important Reducers Halide ions Free metals Metallous ions Nitrite ions Sulfite ions Free halogens (dilute, basic soln) Free halogens (conc, basic soln) C2O4-2 Formed in Reaction Halogens Metal ions Metallic ions Nitrate ions Sulfate ions Hypohalite ions Halate ions CO2 Redox reactions involve the transfer of electrons. The oxidation numbers of at least 2 elements must change. Single displacement, synthesis (combination), decomposition and combustion reactions are often redox reactions. To predict the products of a redox reaction, look at the reagents given to see if there is both an oxidizing agent and reducing agent. When a problem mentions an acid or base solution – it is probably a redox reaction. Balancing Redox Reactions LEO the lions says GER Lose Electrons – Oxidation – an increase in oxidation number Gain Electrons – Reduction – a decrease in oxidation number Oxidizing Agent – substance that is reduced (electron acceptor) Reducing Agent – substance that is oxidized (electron donor) Example: Cu (s) + 2 Ag+ (aq) 2 Ag (s) + Cu+2 (aq) Cu (s) is oxidized Ag+ (aq) is reduced Ag+ (aq) is the oxidizing agent Cu (s) is the reducing agent 3 Methods for Balancing Redox Reactions #1 “Condensed Oxidation State” Method (Zumdahl section 4.10) Relies on inspection method. Limited use - does not account for reactions that need to be done under acidic or basic conditions in order to balance them. #2 - “Half Reaction” Method (Zumdahl section 18.1 and Science Geek step by step section) Works for balancing any redox reaction. Time consuming. #3 - “Hybrid” Method (shown below) “Hybrid” Method 1. Write down charge numbers for each element in the equation – identify what is oxidized and what is reduced. 2. Write a separate half reaction for the oxidation reaction and a separate half reaction for the reduction reaction. 3. In each half reaction: Balance the element that is oxidized (or reduced) Add electrons lost (oxidation reaction) or electrons gained (reduction reaction) to the appropriate side of each half reaction Rebalance each half reaction equalizing electrons 4. Return to original equation using coefficients obtained from your half reaction. 5. Balance charge by adding the H+ (acidic conditions) or OH- (basic conditions) to the appropriate side of the equation. 6. Balance O and/or H by adding H2O to the appropriate side of the equation. 7. Recheck entire equation for balance (elements, charge, electron transfer). Class Examples MnO4- + Fe+2 Fe+3 + Mn+2 Cr2O7-2 + C2H5OH Cr+3 + CO2 Decomposition Reactions Following are some common decomposition reactions. Many of these reactions are initiated by heat. The reverse of many of these reactions are synthesis reactions. Ion in compound Products Example Carbonate CO3-2 oxide, CO2 CaCO3 CaO + CO2 Hydroxide OH- oxide, H2O Ca(OH)2 CaO + H2O Sulfate SO4-2 oxide, SO3 CaSO4 CaO + SO3 Sulfite SO3-2 oxide, SO2 CaSO3 CaO + SO2 Clorate ClO3- chloride, O2 2 KClO3 2 KCl + 3 O2 Ammonium NH4+ acid, NH3 NH4Cl HCl + NH3 Bicarbonate HCO3- H2O, CO2 H2CO3 H2O + CO2 How about the decomposition of ammonium carbonate? Expect ammonium to produce NH3 Expect carbonate to produce CO2 Figure out the final product based on what’s left over