Chapter 4; Reactions in Aqueous Solutions I. Electrolytes vs. NonElectrolytes II. Precipitation Reaction a) Solubility Rules III. Reactions of Acids a) Neutralization b) Acid and Carbonate c) Acid and Metal IV. Metal and Water Chapter 4; Reactions in Aqueous Solutions V. Molarity VI. Acid-Base Titrations VII. Dilution of Solutions A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. The solvent is the substance present in the larger amount In aqueous solutions (aq) *solvent is water *solute can be ionic compounds, aqueous acids, bases, or molecular compounds 1
Conduct electricity in solution? Cations (+) and Anions (-) Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution Strong Electrolyte 100% dissociation C 6 H 12 O 6 (s) H 2 O C 6 H 12 O 6 (aq) NaCl (s) H 2 O Na + (aq) + Cl - (aq) Weak Electrolyte not completely dissociated HNO 2 NO 2- (aq) + H + (aq) Strong Weak Electrolyte Nonelectrolyte Electrolyte Strong Acids Weak Acids Molecular Compounds Strong Bases Weak Bases Ionic Compounds Precipitation Reactions Mix two aqueous solutions made by dissolving ionic compounds in water. If a reaction happens, a precipitate (solid) is formed. Predicting Products of Precipitation Reactions 1) Ionic Compounds are Strong Electrolytes Determine charge on all ions of reactants 2) Using Ion Charges; Predict formula of products. ( + ion of one reactant forms compound with ion of other reactant) 3) Balance Equation 4) Determine is product is solid or aqueous solution 2
Solubility Rules for Common Ionic Compounds In water at 25 0 C Soluble Compounds Compounds containing alkali metal ions and NH + 4 NO 3-, HCO 3-, ClO 3 - Exceptions Cl -, Br -, I - Halides of Ag +, Hg 2 2+, Pb 2+ SO 2- Sulfates of Ag +, Ca 2+, Sr 2+, Ba 2+, 4 Hg 2+, Pb 2+ Insoluble Compounds Exceptions CO 3 2-, PO 4 3-, CrO 4 2-, S 2- OH - Compounds containing alkali metal ions and NH 4 + Compounds containing alkali metal ions and Ba 2+ Predicting Products of Precipitation Reactions (Cont) 5) Determine spectator ions (Ions that are still dissolved in water in the product) 6) Write net ionic equation (Only shows ions involved in forming solid) Precipitation Reactions A precipitation reaction and its equation Precipitate insoluble solid that separates from solution precipitate Pb(NO 3 ) 2 (aq) + 2NaI (aq) PbI 2 (s) + 2NaNO 3 (aq) molecular equation Pb 2+ + 2NO 3- + 2Na + + 2I - PbI 2 (s) + 2Na + + 2NO 3 - ionic equation PbI 2 Pb 2+ + 2I - PbI 2 (s) net ionic equation Na + and NO 3- are spectator ions 3
Writing Net Ionic Equations 1. Write the balanced molecular equation. 2. Write the ionic equation showing the strong electrolytes 3. Determine precipitate from solubility rules 4. Cancel the spectator ions on both sides of the ionic equation Write the net ionic equation for the reaction of silver nitrate with sodium chloride. AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) Ag + + NO 3- + Na + + Cl - AgCl (s) + Na + + NO 3 - Ag + + Cl - AgCl (s) Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases. Acids Produce H + (proton) or (H 3 O) + when dissolved in water Proton donor H 2 O HNO 3 (aq) H + (aq) + (NO 3 ) - (aq) HNO 3 (aq) + H 2 O (l) H 3 O + (aq) + (NO 3 ) - (aq) Monoprotic acids; Produce one H + when dissolved in water HNO 3 H + + NO 3 - Strong electrolyte, strong acid Diprotic acids; Produce two H + when dissolved in water H 2 SO 4 2 H + + SO 4-2 Strong electrolyte, strong acid Triprotic acids; Produce three H + when dissolved in water H 3 PO 4 3 H + + PO 4-3 Weak electrolyte, weak acid 4
Bases Produce (OH) - when dissolved in water Proton (H + ) acceptor H Na(OH) (s) -----> 2 O Na + (aq) + (OH) - (aq) Neutralization Reaction Acid + Base Salt + H 2 O F - (aq) + H 2 O (l) HF (aq) + (OH) - (aq) Acid + Carbonate Salt + CO 2 (g) + H 2 O (l) Displacement Reactions Metal Displaces H from acid or water Carbonate; Contains (CO 3 ) -2 or (HCO 3 ) - Chalk; Ca(CO 3 ) Metal + Acid Salt + H 2 (g) Metal + Water Base + H 2 (g) Use Activity Series to Know if a Reaction Will Happen 5
Oxidation Reduction Reactions Activity Series Transfer of electrons LEO says GER Gain of Electrons REDUCTION Loss of Electrons OXIDATION Compound Formation by Electron Transfer Terminology of Redox Reactions LEO says GER e - X transfer or shift of electrons Y X loses electron(s) X is oxidized X is the reducing agent X increases its oxidation number Y gains electron(s) Y is reduced Y is the oxidizing agent Y decreases its oxidation number 6
Rules for Assigning an Oxidation Number (O.N.) General rules 1. For an atom in its elemental form (Na, O 2, Cl 2, etc.): O.N. = 0 2. For a monoatomic ion: O.N. = ion charge 3. The sum of O.N. values for the atoms in a compound equals zero. The sum of O.N. values for the atoms in a polyatomic ion equals the ion s charge. Rules for specific atoms or periodic table groups 1. For Group 1A(1): O.N. = +1 in all compounds 2. For Group 2A(2): O.N. = +2 in all compounds 3. For hydrogen: O.N. = +1 in combination with nonmetals 4. For fluorine: O.N. = -1 in combination with metals and boron 5. For oxygen: O.N. = -1 in peroxides O.N. = -2 in all other compounds(except with F) 6. For Group 7A(17): O.N. = -1 in combination with metals, nonmetals (except O), and other halogens lower in the group Oxidation numbers of the main-group elements Determining the Oxidation Number of an Element Recognizing Oxidizing and Reducing Agents PROBLEM: PLAN: Determine the oxidation number (O.N.) of each element in these compounds: (a) zinc chloride (b) sulfur trioxide (c) nitric acid The O.N.s of the ions in a polyatomic ion add up to the charge of the ion and the O.N.s of the ions in the compound add up to zero. SOLUTION: (a) ZnCl 2. The O.N. for zinc is +2 and that for chloride is -1. (b) SO 3. Each oxygen is an oxide with an O.N. of -2. Therefore the O.N. of sulfur must be +6. (c) HNO 3. H has an O.N. of +1 and each oxygen is -2. Therefore the N must have an O.N. of +5. PLAN: SOLUTION: Identify the oxidizing agent and reducing agent in each of the following: (a) 2Al(s) + 3H 2 SO 4 (aq) Al 2 (SO 4 ) 3 (aq) + 3H 2 (g) (b) PbO(s) + CO(g) (c) 2H 2 (g) + O 2 (g) Pb(s) + CO 2 (g) 2H 2 O(g) Assign an O.N. for each atom and see which gained and which lost electrons in going from reactants to products. An increase in O.N. means the species was oxidized (and is the reducing agent) and a decrease in O.N. means the species was reduced (is the oxidizing agent). 0 +1 +6-2 +3 +6-2 0 (a) 2Al(s) + 3H 2 SO 4 (aq) Al 2 (SO 4 ) 3 (aq) + 3H 2 (g) The O.N. of Al increases; it is oxidized; it is the reducing agent. The O.N. of H decreases; it is reduced; it is the oxidizing agent. 7
Potassium permanganate and sodium oxalate-inspection Combining elements to form an ionic compound Decomposition Displacing one metal with another 8
Activity Series strength as reducing agents Li K Ba Ca Na Mg Al Mn An Cr Fe Cd Co Ni Sn Pb H 2 Cu Hg Ag Au can displace H 2 from water can displace H 2 from steam can displace H 2 from acid cannot displace H 2 from any source PLAN: Classify each of the following redox reactions as a combination, decomposition, or displacement reaction, write a balanced molecular equation for each, as well as total and net ionic equations for part (c), and identify the oxidizing and reducing agents: (a) magnesium(s) + nitrogen(g) magnesium nitride (aq) (b) hydrogen peroxide(l) (c) aluminum(s) + lead(ii) nitrate(aq) water(l) + oxygen gas aluminum nitrate(aq) + lead(s) Combination reactions produce fewer products than reactants. Decomposition reactions produce more products than reactants. Displacement reactions have the same number of products and reactants. Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M = molarity = moles of solute liters of solution What mass of KI is required to make 500. ml of a 2.80 M KI solution? M KI M KI volume KI moles KI grams KI 1 L 2.80 mol KI 166 g KI 500. ml x x x = 232 g KI 1000 ml 1 L soln 1 mol KI 9
Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Fig. 4.17a,b Equivalence point the point at which the reaction is complete Indicator substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color 4.7 What volume of a 1.420 M NaOH solution is Required to titrate 25.00 ml of a 4.50 M H 2 SO 4 solution? Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. WRITE THE CHEMICAL EQUATION! Dilution H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO 4 M rx M volume acid moles acid moles base volume base acid coef. base 4.50 mol H 2 SO 4 2 mol NaOH 1000 ml soln 25.00 ml x x x = 158 ml 1000 ml soln 1 mol H 2 SO 4 1.420 mol NaOH Moles of solute before dilution (i) M i V i Add Solvent = = Moles of solute after dilution (f) M f V f 10
How would you prepare 60.0 ml of 0.2 M HNO 3 from a stock solution of 4.00 M HNO 3? M i V i = M f V f V i = M fv f 0.200 x 0.06 = = 0.003 L = 3 ml M i 4.00 3 ml of acid + 57 ml of water = 60 ml of solution 11