Experiment 5. Chemical Reactions A + X AX AX A + X A + BX AX + B AZ + BX AX + BZ

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1 Experiment 5 Chemical Reactions OBJECTIVES 1. To observe the various criteria that are used to indicate that a chemical reaction has occurred. 2. To convert word equations into balanced inorganic chemical reactions. 3. To study four of the major types of inorganic chemical reactions. INTRODUCTION Although there are a large number or inorganic chemical reactions, most of these can be classified as one of four types. The first of these is the formation of a single compound from two or more pure substances (elements or compounds). This is classified as a combination reaction and can be represented as the following chemical equation. A + X AX The second type of reaction can be viewed as the reverse of the first type and is classified as a decomposition reaction. In this reaction, represented below, a single compound decomposes into two or more pure substances. This usually requires the application of heat. AX A + X The third type of reaction is termed a single replacement reaction. In this reaction, represented below, an element replaces one of the elements in a compound. This reaction will occur if the element that is replaced is lower in the activity series. A + BX AX + B The fourth type of reaction is termed a double replacement reaction. In this reaction, represented below, two compounds in solution switch partners. The cation of one compound exchanges with the cation of the other compound. Often when this type of reaction occurs a precipitate forms because one of two products is insoluble in water. AZ + BX AX + BZ In this experiment you will observe examples of each of these four types of chemical reactions. In each reaction there will be evidence that a reaction has occurred. You are to observed and record the evidence. The evidence may include any of the following: (1) a gas is produced; (2) a precipitate is formed; (3) a color change is observed; (4) a temperature change is noted. In chemical equations, the substances written to the left of the arrow are called the reactants and the substances written to the right of the arrow are called the products. Several symbols are used in chemical equations to describe the reaction conditions. Some of these are listed in Table 5-1.

2 Table 5-1 Symbols in Chemical Equations Symbol Meaning produces, yields (separates reactants from products) + added to, reacts with (separates two or more reactants or products) heat applied to the reaction (written above the ) NR no reaction (written after the ) (s) solid or precipitate (l) liquid (g) gas (aq) aqueous solution In order to write a chemical equation it is necessary to predict the products that are formed. Initially, this can be a difficult task. To assist you in writing chemical equations, the equations in word form will be given for each reaction. You will then convert the word equations into balanced chemical equations. Balancing the equation means putting appropriate coefficients in front of the reactants and products to ensure that the same number of atoms of each element are in the reactants and the products. Examples for each of the four types of reaction are shown below. Example 5-1 Combination Reaction calcium (s) + oxygen (g) calcium oxide (s) Ca (s) + O 2(g) CaO (s) (unbalanced) 2 Ca (s) + O 2(g) 2 CaO (s) (balanced) Example 5-2 Decomposition Reaction calcium chloride dihydrate (s) calcium chloride (s) + water (g) CaCl 2. 2H 2 O (s) CaCl 2(s) + H 2 O (unbalanced) CaCl 2. 2H 2 O (s) CaCl 2(s) + 2 H 2 O (balanced) Example 5-3 Single Replacement Reaction zinc (s) + silver nitrate (aq) zinc nitrate (aq) + silver (s) Zn (s) + AgNO 3(aq) Zn(NO 3 ) 2(aq) + Ag (s) (unbalanced) Zn (s) + 2AgNO 3(aq) Zn(NO 3 ) 2(aq) + 2Ag (s) (balanced) Example 5-4 Double Replacement iron(iii) nitrate (aq) + sodium hydroxide (aq) Fe(NO 3 ) 3(aq) + NaOH (aq) Fe(NO 3 ) 3(aq) + 3NaOH (aq) iron(iii) hydroxide (s) + sodium nitrate (aq) Fe(OH) 3(s) +NaNO3(aq) (unbalanced) Fe(OH) 3(s) +3NaNO3(aq) (balanced)

3 PROCEDURE For each of the following procedures, record your observations in the Data Table. On the page following the observations, word equations are written for each reaction. Write the corresponding balanced chemical equation. A. Combination Reactions (For safety reason these two reactions will be performed by the laboratory instructor.) 1. Hold a 2 cm strip of magnesium ribbon with crucible tongs and ignite the metal in a hot burner flame. 2. Mix together 2 g of powdered zinc and 1 g of powdered sulfur in a crucible. Place the crucible on a wire gauze in a fume hood. Use the flame from a Bunsen burner to ignite the mixture. B. Decomposition Reactions 1. Add a few crystals of copper(ii) sulfate pentahydrate to a dry test tube. Holding the test tube with a test tube holder, heat the part of the test tube containing the solid strongly. Note the change in color and texture and observe the inside wall of the upper region of the test tube. 2. Take two dry 250-mL Erlenmeyer flasks. To one of the flasks, add sodium hydrogen carbonate (baking soda) so as to sparsely cover the bottom of the flask. (a) Take a lighted wooden splint and plunge it into the Erlenmeyer flask that doesn t contain the baking soda. Observe the time required for the flame to go out. (b) Place the flask containing the baking soda on a wire gauze on a ring stand. Use a Bunsen burner to heat the flask strongly and note the inside wall in the upper region of the flask. When you see evidence of moisture, plunge a lighted wooden splint into the flask and observe how long it burns. Compare this time with the value obtained for the empty flask in part (a). C. Single Replacement Reactions 1. Put about 2 ml (30 drops) of silver nitrate into a test tube and add a small piece of copper wire. Place the tube in a rack and leave it for at least 10 minutes. Record your observations. 2. Place a small piece of magnesium ribbon into a test tube containing about 2 ml (30 drops) of dilute hydrochloric acid. 3. Place a small piece of calcium metal into a test tube containing about 2 ml (30 drops) of water. D. Double Replacement Reactions 1. Put 2 ml (30 drops) of silver nitrate, lead(ii) nitrate, and aluminum nitrate into separate test tubes. Add about 2 ml of potassium iodide to each test tube and check for evidence of reaction. 2. Put 2 ml (30 drops) of silver nitrate, lead(ii) nitrate, and aluminum nitrate into separate test tubes. Add about 2 ml of sodium phosphate to each test tube and check for evidence of reaction.

4 EXPERIMENT 5 DATE NAME SECTION Procedure Evidence of Reaction A. Combination Reactions 1. Mg (s) + O 2(g) 2. Zn (s) + S (s) B. Decomposition Reactions 1. CuSO 4. 5H 2 O (s) 2. NaHCO 3(s) C. Single Replacement Reactions 1. Cu (s) + AgNO 3(aq) 2. Mg (s) + HCl (aq) 3. Ca (s) + H 2 O (aq) D. Double Replacement Reactions 1. AgNO 3(aq) + KI (aq) 2. Pb(NO 3 ) 2(aq) + KI (aq) 3. Al(NO 3 ) 3(aq) + KI (aq) 4. AgNO 3(aq) + Na 3 PO 4(aq) 5. Pb(NO 3 ) 2(aq) + Na 3 PO 4(aq) 6. Al(NO 3 ) 3(aq) + Na 3 PO 4(aq)

5 Convert Each Word Equation into a Balanced Chemical Equation A. Combination Reactions 1. magnesium metal (s) + oxygen (g) magesium oxide (s) 2. zinc metal (s) + sulfur (s) zinc sulfide (s) B. Decomposition Reactions 1. copper(ii) sulfate pentahydrate (s) copper(ii) sulfate (s) + water (l) 2. sodium hydrogen carbonate (s) sodium carbonate (s) + water (l) + carbon dioxide (g) C. Single Replacement Reactions 1. copper metal(s) +silver nitrate(aq) copper(ii) nitrate(aq) +silver metal(s) 2. magnesium metal(s) +hydrochloric acid(aq) magnesium chloride(aq) +hydrogen(g) 3. calcium metal (s) +water (l) calcium hydroxide (s) + hydrogen (g)

6 D. Double Replacement Reactions 1. silver nitrate (aq) +potassium iodide (aq) silver iodide (s) +potassium nitrate (aq) 2. lead(ii) nitrate (aq) +potassium iodide (aq) lead(ii) iodide (s) +potassium nitrate (aq) 3. aluminum nitrate (aq) +potassium iodide (aq) no reaction 4. silver nitrate (aq) +sodium phosphate (aq) silver phosphate (s) +sodium nitrate (aq) 5. lead(ii) nitrate (aq) +sodium phosphate (aq) lead(ii) phosphate (s) +sodium nitrate (aq) 6. aluminum nitrate (aq) +sodium phosphate (aq) aluminum phosphate (s) +sodium nitrate (aq)

7 POSTLABORATORY ASSIGNMENT NAME 1. Give the chemical formula for each of the following substances. Each is one of the products formed in the reaction indicated. Refer to your observations recorded in the Data Table along with the corresponding chemical equations. (a) white smoke (Procedure A.1) (b) colorless liquid (Procedure B.1) (c) flame-extinguishing gas (Procedure B.2) (d) gray solid (Procedure C.1) (e) colorless gas (Procedure C.2) (f) yellow precipitate (Procedure D.2) (g) yellow precipitate (Procedure D.3) (h) white precipitate (Procedure D.6) 2. Convert each of the following word equations into a balanced chemical equation (a) copper metal (s) + oxygen (g) copper(i) oxide (s) (b) mercury(ii) nitrate (aq) + potassium iodide (aq) mercury(ii) iodide (s) + potassium nitrate (aq) (c) aluminum metal (s) + hydrochloric acid (aq) aluminum chloride (aq) + hydrogen (g) (d) sulfuric acid (aq) + sodium hydroxide (aq) sodium sulfate (aq) + water (l) (e) iron(iii) carbonate (s) iron(iii) oxide (s) + carbon dioxide (g)

8 3. As you observed in the reactions in Part D, the combination of aqueous solutions of ionic compounds can often result in the formation of a precipitate. This is due to the fact that compound formed by the cation of one compound with the anion of the other compound is insoluble in water. Whether or not a precipitate will form can frequently be predicted using the following set of general solubility rules. General Solubility Rules 1. Nearly all nitrates and acetates are soluble. 2. All chlorides are soluble except AgCl, Hg2Cl2, and PbCl2. 3. All sulfates are soluble except BaSO4, SrSO4, and PbSO4. 4. Most of the alkali metal (Li, Na, K, etc.) and ammonium salts are soluble. 5. All oxides and hydroxides are insoluble except those of the alkali metals and certain alkaline earth metals (Ca, Sr, Ba). 6. All sulfides are insoluble except those of the alkali metals and ammonium sulfide. 7. All phosphates and carbonates are insoluble except those of the alkali metals and ammonium salts. Use these rules to determine which of the following compounds will be insoluble in water. Circle the ones that are insoluble. (a) AgBr (b) Zn(OH) 2 (c) PbCO 3 (d) MgSO 4 (e) Cd(NO 3 ) 2 (f) Fe3(PO 4 ) 2 (g) NiS (h) KOH (i) Cr 2 O 3 (j) (NH 4 ) 2 CO 3 4. Whether or not a single replacement reaction is possible under standard conditions can be predicted using what is termed the activity or electromotive series. In this series (shown below) an element can replace the ion of any element below it. For example as you observed in Procedure C.1, the element copper replaced the silver ion in silver nitrate to form silver metal and the copper(ii) ion. Use this Series to determine which of the following single replacement reactions will occur. For a reaction that occurs, complete and balance the equation. If the reaction will not occur, write N.R. for no reaction. Activity (Electromotive) Series (a) Mg (s) + NiCl 2(aq) Li K Ba Ca (b) Pb (s) + CaSO 4(aq) Na Mg Al Zn Fe (c) Zn (s) + HCl (aq) Ni Pb (H) Cu Ag Au

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