Santa Monica College Chemistry 11

Transcription

1 Types of Reactions Objectives The objectives of this laboratory are as follows: To perform several types of simple chemical reactions, To become familiar with some common observable signs of chemical reactions, To identify the products formed and write balanced equations for the reactions studied, To devise a partial activity series using the results of the single displacement reactions, and To practice writing complete ionic equations and net ionic equations for the double displacement reactions (reactions in aqueous solution) Background Matter undergoes three kinds of change: physical, chemical, and nuclear. Physical changes do not alter the composition of matter; examples include freezing and evaporation. However, chemical changes, or chemical reactions, result in the formation of new substances as chemical bonds are formed and/or broken and thus the composition of matter is changed. Some relatively simple but common types of chemical reactions will be studied in this experiment. These reactions are often accompanied by observable changes as they occur. Examples include, but are not limited to: Color change Formation of a precipitate noted as the formation of a cloudy solution, formation of a gel, or an obvious solid Evolution of a gas noted as bubbling in the solution Appearance or disappearance of a distinct separation between two or more liquids Evolution of heat noted as a temperature increase Absorption of heat noted as a temperature decrease Plating out of one metal on another Decomposition, pitting, or the disappearance of a solid metal Reaction Types Combination Reactions occur when two or more substances, elements or compounds, combine to form one new substance: A + B AB. For example, if a mixture of hydrogen and oxygen gases is sparked, they combine explosively to produce water: 2 H 2 (g) + O 2 (g) 2 H 2 O (l) Decomposition Reactions occur when a compound breaks apart to yield two or more new substances: AB A + B. These reactions often require the input of heat, light or a catalyst to occur. For example, when heated to high temperatures potassium chlorate decomposes to yield potassium chloride and oxygen gas: 2 KClO 3 (s) 2 KCl (s) + 3 O 2 (g) Types of Reactions Page 1 of 5

2 Single Displacement Reactions involve the displacement of one element in a compound by another element: A + BC AC + B. Several examples of these reactions are given below. Displacement of one metal by another: Pb (s) + Cu(NO 3 ) 2 (aq) Pb(NO 3 ) 2 (aq) + Cu (s) Displacement of hydrogen gas from an acid by a metal: Mg (s) + 2 HCl (aq) MgCl 2 (aq) + H 2 (g) Displacement of hydrogen gas from water by a metal: 2 K (s) + 2 H 2 O (l) 2 KOH (aq) + H 2 (g) Displacement of one halogen by another: Cl 2 (aq) + 2 NaBr (aq) 2 NaCl (aq) + Br 2 (aq) Single displacement reactions may also be described as Redox reactions. Consider again the reaction between lead and copper(ii) nitrate (first example above). Elemental lead is oxidized to lead(ii) and copper is reduced from copper(ii) to elemental copper. Two electrons are transferred from lead to copper in this process: Pb (s) Pb 2+ (aq) + 2 e - oxidation of lead Cu 2+ (aq) + 2e - Cu (s) reduction of copper The ability of one metal to displace another depends on their relative ease of oxidation. A more active metal (one that is more easily oxidized) displaces a less active metal. In the reaction above, lead is more active than copper. The relative activities of metals can be tabulated in an Activity Series, which ranks metals by relative ease of oxidation. A metal that displaces hydrogen gas from an acid is more active than hydrogen. A metal that displaces hydrogen gas from an acid, but not from water, is less active than one that can displace hydrogen from both acids and water. Double Displacement (Exchange) Reactions occur when two compounds that form ions in aqueous solution react by switching ion partners: AB + CD AD + CB. One of three conditions must be met for these reactions to occur: (1) the formation of an insoluble ionic compound, observed as a precipitate, (2) the production of water, or (3) the formation of a gas. Precipitation Reactions occur when the ions in two aqueous solutions become exchanged and combine to form a compound that is insoluble in water. A Solubility Chart is used to determine if a product is insoluble; one is provided on the next page. This insoluble product is called the precipitate. For example, when solutions of sodium phosphate and calcium chloride are mixed, all the calcium ions will be precipitated out as solid calcium phosphate: 3 CaCl 2 (aq) + 2 Na 3 PO 4 (aq) Ca 3 (PO 4 ) 2 (s) + 6 NaCl (aq) Neutralization Reactions occur between an Arrhenius acid (HX) and an Arrhenius base (MOH). Water is produced as hydrogen ions (H + ) from the acid combine with hydroxide ions (OH - ) from the base. For example, when solutions of hydrochloric acid and sodium hydroxide are combined, water and sodium chloride are formed: HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) Types of Reactions Page 2 of 5

3 Gas Forming Reactions result in the formation of a gaseous product which leaves the reaction mixture as bubbles. Gases produced in these reactions include hydrogen sulfide, sulfur dioxide, carbon dioxide and ammonia. Hydrogen sulfide gas, H 2 S, is formed directly in the exchange reaction between an acid and a sulfide: Na 2 S (aq) + 2 HCl (aq) H 2 S (g) + 2 NaCl (aq) Sulfur dioxide gas, SO 2, is formed by the decomposition of sulfurous acid, which is itself initially formed in an exchange reaction between an acid and a sulfite (or a bisulfite): Na 2 SO 3 (s) + 2 HCl (aq) H 2 SO 3 (aq) + 2 NaCl (aq) H 2 SO 3 (aq) H 2 O (l) + SO 2 (g) Na 2 SO 3 (s) + 2 HCl (aq) H 2 O (l) + SO 2 (g) + 2 NaCl (aq) Carbon dioxide gas, CO 2, is formed by the decomposition of carbonic acid, which is itself initially formed in an exchange reaction between an acid and a carbonate (or a bicarbonate): Na 2 CO 3 (s) + 2 HCl (aq) H 2 CO 3 (aq) + 2 NaCl (aq) H 2 CO 3 (aq) H 2 O (l) + CO 2 (g) Na 2 CO 3 (s) + 2 HCl (aq) H 2 O (l) + CO 2 (g) + 2 NaCl (aq) Ammonia gas, NH 3, is formed by the decomposition of ammonium hydroxide, which is itself initially formed in an exchange reaction between an Arrhenius base and an ammonium salt: NH 4 Cl (aq) + NaOH (aq) NH 4 OH (aq) + NaCl (aq) NH 4 OH (aq) H 2 O (l) + NH 3 (g) NH 4 Cl (aq) + NaOH (aq) H 2 O (l) + NH 3 (g) + NaCl (aq) Solubility Chart of Common Ionic Compounds General Rule All compounds containing Group IA metals and Ammonium are soluble. All Nitrates and Acetates are soluble. Most Halides Chlorides, Bromides and Iodides are soluble. Most Sulfates are soluble. Most Phosphates, Carbonates and Chromates are insoluble. Most Sulfides are insoluble. Most Hydroxides are insoluble. Important Exceptions None None Halides of silver, lead(ii), and mercury(i) are insoluble. Sulfates of lead(ii), strontium, barium, and calcium are insoluble. Phosphates, Carbonates and Chromates of Group 1A metals and ammonium are soluble. Sulfides of Group 1A metals and ammonium are soluble. Group 2A metals sulfides are somewhat soluble but eventually decompose in aqueous solution. Hydroxides of Group 1A metals and barium are soluble. Solubility of Group 2A metals increase going down the group (opposite trend of sulfates). Types of Reactions Page 3 of 5

4 Procedure Chemicals and Equipment Solids: Mg, CuSO 4 5H 2 O, Ca, Cu, Zn, NaHCO 3 Solutions: 6M HCl, 6M NaOH, 6M H 2 SO 4, 1M NH 4 NO 3, and 0.1M solutions of CuSO 4, ZnSO 4, AgNO 3, NaCl, Ni(NO 3 ) 2, Pb(NO 3 ) 2 and K 2 CrO 4 Equipment: Crucible tongs, test tubes (2 large, 1 medium and 10 small), test tube holder, test tube rack, 100-mL beaker, 10-mL graduated cylinder, stirring rod, red litmus paper, wash bottle with distilled water, Bunsen burner Supplies for Instructor Demos: saturated CaO solution, sucrose, 18M H 2 SO 4, Na, means of cutting and removing Na from its storage vessel, 3-4 beakers, straw, stirring rod Safety 1 Do not stare directly at the magnesium when it burns as the light can hurt your eyes. 2 Do not touch the metals Ca or Mg with your hands. 3 Skin discoloration will result from contact with silver nitrate. 4 Be extremely cautious with all acids and bases as they are corrosive and can burn your skin. Chemical Waste/Clean-Up At the end of this lab, all reaction products/residues/waste/leftovers must be placed in the waste container(s) provided. In addition, use a wash bottle to rinse any chemicals stuck on your glassware directly into the waste container. Do not leave any metal pieces in the sink. Part A: Combination Reactions 1. Instructor Demonstration. Pour about 30 ml of a clear saturated solution of calcium oxide (note that CaO actually turns into Ca(OH) 2 in water, but the relevant product is the same as for CaO) into a 100-mL beaker. Then use a straw to blow bubbles into the solution for several seconds. Record your observations. 2. Use your crucible tongs to place a strip of magnesium metal in the hottest portion of a Bunsen burner flame. As soon as it begins to ignite move the metal from the flame into the open air. Record your observations. Do not touch the reaction residue! Part B: Decomposition Reactions 1. Instructor Demonstration. Perform this reaction in the fume hood. Fill a 100-mL beaker about ⅓ full of granulated sugar (sucrose, C 12 H 22 O 11 ). Add about 20 ml of concentrated 18M sulfuric acid and stir until mixed well. Continue stirring as the reaction progresses and the mixture darkens. Do not touch the reaction products or the beaker with your hands; use a stirring rod to guide the solid product that forms. Record your observations. 2. Place a small scoop of solid copper(ii) sulfate pentahydrate in a medium dry test tube. Use a test tube holder to hold the tube at about a 45º angle, and heat the sample gently by wafting it back and forth in a Bunsen burner flame for a few minutes. Record your observations. Types of Reactions Page 4 of 5

5 Part C: Single Displacement Reactions Combine the sets of chemicals listed below. Be sure to place the metals in the test tubes first before adding the solutions (or water). Only 1-2 pieces of metal and about 1 ml of solution are needed for each reaction. You can estimate 1 ml by either counting drops from the reagent bottle dispenser, or, by measuring out 1 ml once in your graduated cylinder, then transferring it to a test tube to see how far it fills the test tube. Use a large test tube for Reaction 2 and small test tubes for Reactions 3-8 (Reaction 1 is a demo). The test tubes should be cleaned and then rinsed with distilled water before use. Record your observations on your lab report. If no discernable changes occur upon immediately mixing reagents, let the reaction mixture stand for five to ten minutes before observing again. Not all of the chemical combinations will yield observable results. 1. Instructor Demonstration. Perform this reaction in the fume hood with the sash lowered. Cautiously add a small piece of sodium metal to a large beaker filled half-way with water. 2. Calcium metal and water (15 ml) in a large test tube 3. Zinc metal and water 4. Copper metal and 6M hydrochloric acid 5. Zinc metal and 6M hydrochloric acid 6. Zinc metal and 0.1M copper(ii) sulfate 7. Copper metal and 0.1M zinc sulfate 8. Copper metal and 0.1M silver nitrate Part D: Double Displacement Reactions Combine the sets of chemicals listed below. About 1 ml of each solution is required for each reaction, unless noted otherwise. Please read the instructions for Part C on how to estimate 1 ml. Use small test tubes for Reactions 1-4, a 100-mL beaker for Reaction 5, and a large test tube for Reaction 6. The test tubes and beaker should be cleaned and then rinsed with distilled water before use. Record your observations on your lab report M silver nitrate and 0.1M sodium chloride M nickel(ii) nitrate and three drops of 6M sodium hydroxide M lead(ii) nitrate and 0.1M potassium chromate 4. 1M ammonium nitrate and 6M sodium hydroxide. Warm the mixture gently by wafting the test tube back and forth through a Bunsen burner flame. While doing this, use the crucible tongs to hold a strip of moistened red litmus paper at the mouth of the test tube (do not let it touch the sides of the tube). Note any color changes to the litmus paper. Then remove the test tube from the flame and cautiously (and quickly) note any odor that emanates from it. 5. Pour about 5 ml of 6M hydrochloric acid into a 100-mL beaker, then carefully add several spatulas of solid sodium bicarbonate. 6. Combine about 5 ml each of 6M sodium hydroxide and 6M sulfuric acid in a large test tube. Mix with a stirring rod. Cautiously feel the outside of the test tube. Be extra careful with the sulfuric acid it is very corrosive! Types of Reactions Page 5 of 5