TYPES OF CHEMICAL REACTIONS Precipitation Reactions Compounds Soluble Ionic Compounds 1. Group 1A cations and NH 4 + 2. Nitrates (NO 3 ) Acetates (CH 3 COO ) Chlorates (ClO 3 ) Perchlorates (ClO 4 ) Solubility Rules for Ionic Compounds in Water Exceptions 3. Chlorides, bromides, iodides (Cl, Br, I ) Ag +, Pb 2+, Hg 2 2+ 4. Sulfates (SO 4 2 ) Pb 2+, Hg 2 2+, Sr 2+, Ca 2+, Ba 2+ Insoluble Ionic Compounds 5. Hydroxides (OH ) Group 1A cations, NH 4+, Ca 2+, Sr 2+, Ba 2+ 6. Phosphates (PO 4 3 ) Sulphides (S ) Carbonates (CO 3 ) Chromates (CrO 4 ) Group 1A cations, NH 4 + 1. What precipitate (if any) will form? (a) AgNO 3 (aq) + Na 2 S (aq) (b) KCl (aq) + Pb(NO 3 ) 2 (aq) 2. For each of the following reactions, write the formula equation, the complete ionic equation, and the net ionic equation. (a) Aqueous nickel(ii) chloride reacts with aqueous sodium hydroxide. (b) K 2 SO 4 (aq) + FeBr 3 (aq) Page 1
Acid-Base Reactions 3. Write the balanced formula, complete ionic, and net ionic equations for each of the following acid base reactions. (a) HCl (aq) + KOH (aq) (b) H 2 SO 4 (aq) + Mg(OH) 2 (aq) Oxidation-Reduction Reactions Rules Rules for Assigning Oxidation States 1. The oxidation number of an atom of a pure element is 0. 2. The oxidation number of a monatomic ion equals its charge. 3. The oxidation number of fluorine in its compounds is always 1. 4. In most cases, the oxidation number of oxygen is 2, except when it is bonded to fluorine (where it may be +1 or +2), and in peroxides (where it is 1). 5. Hydrogen has an oxidation number of +1 unless it is combined with a metal, in which case its oxidation number is 1. 6. The sum of the oxidation states of all the atoms in a compound or ion is equal to the charge on the compound or ion. Examples Oxygen in O 2, sulfur in S 8, iron in metallic Fe, chlorine in Cl 2 Oxidation number of Cu 2+ is +2; that of S 2 is 2. In HF, the oxidation number of F is 1 and of H is +1. In H 2 O, the oxygen atom has an oxidation number of 2. In OF 2, the oxidation number of O is +2. In H 2 O 2, each oxygen atom has an oxidation number of 1. In HCl, the oxidation number of H is +1 and of Cl is 1. In NaH, the oxidation number of H is 1 and of Na is +1. For H 2 O, the sum of the oxidation states of all atoms is zero. For CO 3, it is 2. For NH 4+, it is +1. 4. Assign the oxidation states to each atom in the following compounds. (a) CaF 2 (b) ICl 5 (c) KmnO 4 (d) SO 4 Page 2
Oxidation Reduction Oxidation Agent Reducing Agent Some Definitions The loss of one or more electrons Increase in oxidation number The gain of one or more electrons Decrease in oxidation number The substance that causes oxidation to occur The substance that accepts electrons from another species It is reduced The substance that causes reduction to occur The substance that supplies the electrons transferred to a second species It is oxidized 5. For each reaction, identify the atoms that are oxidized and reduced, and specify the oxidizing and reducing agents. (a) Zn (s) + Cu 2+ (aq) (b) 2AgCl (s) + H 2 (g) Zn 2+ (aq) + Cu(s) 2H + (aq) + 2Ag (s) + 2Cl (aq) Balancing Redox Equations In Acidic Solutions 1. Break the overall unbalanced equation into half-reactions: oxidation half-reaction reduction half-reaction 2. Balance the atoms in each half-reaction: balance all atoms except for O and H balance O by adding H 2 O balance H by adding H + 3. Balance the charges in each half-reaction by adding electrons to the appropriate side. 4. As needed, multiply one or both of the half-reactions by some coefficient so that both half-reactions will have the same number of electrons. 5. Add the half-reactions together and simplify. 6. Check the final result to make sure both atoms and charges are balanced. 6. Balance the following oxidation-reduction reactions that occur in acidic solution using the halfreaction method. (a) Cu (s) + NO 3 (aq) Cu 2+ (aq) + NO (g) (b) Cr 2 O 7 (aq) + Cl (aq) Cr 3+ (aq) + Cl 2 (g) Page 3
Balancing Redox Equations In Basic Solutions 1. Break the overall unbalanced equation into half-reactions: oxidation half-reaction reduction half-reaction 2. Balance the atoms in each half-reaction: balance all atoms except for O and H balance O by adding H 2 O balance H by adding H + 3. Balance the charges in each half-reaction by adding electrons to the appropriate side. 4. As needed, multiply one or both of the half-reactions by some coefficient so that both half-reactions will have the same number of electrons. 5. Add the half-reactions together and simplify. 6. Because H + does not exist at any appreciable concentration in a basic solution, add an appropriate amount of OH to both sides of the equation 7. H + and OH react to form H 2 O, simplify the number of H 2 O appearing on both sides of the equation 8. Check the final result to make sure both atoms and charges are balanced. 7. Balance the following oxidation-reduction reactions that occur in basic solution using the halfreaction method. (a) MnO 4 (aq) + CN (aq) MnO 2 (s) + OCN (aq) (b) CrI 3 (s) + Cl 2 (g) CrO 4 (aq) + IO 4 (aq) + Cl (aq) Concentration of Solutions Molarity (M) = moles solute liter of solution We can prepare solutions of known molarity (standard solutions) from weighed samples of solute. from concentrated solutions (stock solutions). Dilution: M initial x V initial = M final x V final volumetric flask Page 4
8. Calculate the mass of NaCl needed to prepare 175 ml of a 0.500 M NaCl solution. [MM NaCl = 58.44 g/mol] 9. What volume of a 2.48 M NaOH solution contains 31.52 g of the dissolved solid? [MM NaOH = 40.00 g/mol] 10. Calculate the molarity of Cl ion in a solution prepared by dissolving 9.82 g CuCl 2 in enough water to make 600.0 ml of solution [MM CuCl 2 = 134.45 g/mol] 11. What volume of a 12 M hydrochloric acid must be used to prepare 600. ml of 0.30 M HCl solution? Solution Stoichiometry Two methods for converting between quantity and number of moles If the mass of a compound is given or needed, we use the molar mass of the compound. If the volume of solution of a compound is given or needed, we use the molarity of the solution. 12. What mass of Fe(OH) 3 is produced when 35 ml of a 0.250 M Fe(NO 3 ) 3 solution is mixed with 55 ml of a 0.180 M KOH solution? 13. Limestone, CaCO 3, reacts with HCl to produce the salt CaCl 2, carbon dioxide, and water: CaCO 3 (s) + 2HCl (aq) CaCl 2 (aq) + CO 2 (g) + H 2 O (l) How many grams of calcium carbonate will react completely with 10.0 ml of 3.00 M HCl? Chemists study acid-base reactions quantitatively through titration (volumetric analysis). Page 5
Titration End point Indicator Some Definitions A procedure to determine the concentration of one substance by adding a measured amount of a second substance. Point at which an indicator changes color the titration is stopped. A compound that changes color as an acidic solution becomes basic, or vice versa. Equivalence point Point at which stoichiometrically equivalent amounts of the two reactants are present. Standardization of a solution Process by which one determines the concentration of a solution by titrating it against a standard solid. 14. What is the molarity of a hydrochloric acid solution if 36.7 ml of the HCl solution is required to react with 43.2 ml of 0.236 M sodium hydroxide solution? 15. Calculate the molarity of a solution of H 2 SO 4 if 40.0 ml of the solution neutralizes 0.364 g of Na 2 CO 3. 16. To standardize an HCl solution, a chemist weighs 0.210 g of pure Na 2 CO 3 into a flask. She finds that it takes 5.50 ml HCl to react completely with the Na 2 CO 3. Calculate the concentration of the HCl. 2HCl (aq) + Na 2 CO 3 (aq) H 2 O (l) + 2NaCl (aq) + CO 2 (g) Page 6