EXPERIMENT 5: CHEMICAL REACTIONS AND EQUATIONS

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1 PURPOSE EXPERIMENT 5: CHEMICAL REACTIONS AND EQUATIONS To perform and observe simple chemical reactions. To identify the products of chemical reactions and write balanced equations for those reactions. To name reactants and products. BACKGROUND EVIDENCE OF A CHEMICAL REACTION A chemical reaction has taken place when a substance has changed its chemical composition. The evidence that a chemical reaction has taken place can be one or more of the following: i. the appearance of a gas (bubbles) ii. the appearance of a solid (precipitate) iii. a change of color iv. a change of temperature. REACTION TYPES 1. Combination Reactions 2. Decomposition Reactions 3. Neutralization Reactions 4. Precipitation Reactions 5. Gas Forming Reactions Some reactions can be classified as more than one of the reaction types listed above, and also may be called single-replacement or double-replacement reactions.

2 Combination Reactions Two or more substances combine together to form one new substance. When charcoal burns, carbon reacts (burns) with oxygen to form carbon dioxide. Carbon(solid) + oxygen(gas) carbon dioxide(gas) written as a chemical equation: C(s) + O 2 (g) CO 2(g) Decomposition Reactions A single substance breaks apart to give two or more new substances. is: Hydrogen peroxide decomposes to form liquid water and oxygen gas. The hydrogen peroxide is initially dissolved in water; this is called an aqueous solution. Hydrogen peroxide(aqueous) water(liquid) + oxygen(gas) H 2 O 2 (aq) H 2 O(l) + O 2 (g) However, the above equation is not balanced. The balanced equation 2H 2 O 2 (aq) 2H 2 O(l) + O 2 (g) Neutralization Reactions A reaction between an acid and a base to produce a salt and water is called a neutralization reaction. In this experiment the only bases used will be those containing the hydroxide ion OH. The reaction between hydrochloric acid and sodium hydroxide to produce sodium chloride and water. Since sodium chloride is soluble, its formation cannot be seen, but there is evidence that a reaction has taken place because the reaction mixture gets warmer. NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) It is possible to test whether a solution is acidic or basic by using an indicator. An indicator is a compound that changes color depending upon the acidity of the solution to which it is added. One indicator is litmus, and the most convenient form is as litmus paper. Litmus paper colors: In an acidic solution, litmus paper is red. In a neutral solution, litmus paper is lavender In a basic solution, litmus paper is blue.

3 Precipitation Reactions A reaction in which an insoluble substance is produced when two aqueous salt solutions are mixed together is called a precipitation reaction. A salt is a compound that consists of a positive ion, called a cation, and a negative ion, called an anion. When a salt dissolves in water, it separates into its constituent cations and anions. This is an aqueous solution and will be denoted with (aq). A list of the names and formulae of the commonly occurring ions follows this section for your convenience. Also, following this section, is a set of solubility rules which will enable you to identify any insoluble solid (precipitate) formed during a reaction. A white precipitate is formed when solutions of silver nitrate and sodium chloride are mixed. 1. Determine the formulae of the reactants: silver nitrate is AgNO 3 sodium chloride is NaCl 2. Consider the ions that are formed when the salts dissolve: Ag + + NO 3 and Na + + Cl 3. Switch the partners of the original salts (Be sure each cation is paired with a new anion): Ag + pairs with Cl and Na + pairs with NO 3 4. Write the formulae and names for the possible new compounds: AgCl and NaNO 3 silver chloride sodium nitrate 5. Determine if any of the new compounds are insoluble using the solubility rules: The solubility rules say that all nitrates are soluble so the sodium nitrate is not the precipitate. The silver chloride must therefore be the precipitate, and this checks with the rule that says that all chlorides are soluble except Ag +, Hg 2 2+ and Pb 2+. The equation for the reaction can now be written and balanced: AgNO 3 (aq) + NaCl(aq) AgCl(s) + NaNO 3 (aq) The (s) denotes a solid. This is the precipitate.

4 A white precipitate is formed when aqueous solutions of sodium phosphate and calcium chloride are mixed. The final form of the equation would be: 2Na 3 PO 4 (aq) + 3CaCl 2 (aq) Ca 3 (PO 4 ) 2 (s) + 6NaCl(aq) This reaction may also be classified as a double-replacement reaction! Gas-forming Reactions A reaction in which a gas is produced, bubbles form in the solution, is called a gas forming reaction. Such reactions can be classified into other reaction types as well. When Zn metal is reacted with aqueous hydrochloric acid, the metal dissolves and a gas is evolved. Zn(s) + 2HCl(aq) ZnCl 2 (aq) + H 2 (g) This reaction may also be classified as an single-replacement reaction! Bubbles of gas are formed when an aqueous solution of sodium carbonate reacts with an aqueous solution of hydrochloric acid. Na 2 CO 3 (aq) + HCl(aq) CO 2 (g) + NaCl(aq) + H 2 O(l) This reaction is an example of a general rule which states that: carbon dioxide is produced when any carbonate or bicarbonate reacts with any acid. This reaction could also be classified as a neutralization reaction.

5 NAMES AND FORMULAE FOR COMMONLY OCCURRING IONS CATIONS ANIONS Hydrogen H + Magnesium Mg 2+ Hydride H Oxide O 2 Lithium Li + Calcium Ca 2+ Hydroxide OH 2 Peroxide O 2 Sodium Na + Strontium Sr 2+ Fluoride F Sulfide S 2 Potassium K + Barium Ba 2+ Chloride Cl 2 Sulfite SO 3 Silver Ag + 2+ Mercury(I) Hg 2 Bromide Br 2 Sulfate SO 4 Copper(I) Cu + Mercury(II) Hg 2+ Iodide I 2 Thiosulfate S 2 O 3 + Ammonium NH 4 Copper(II) Cu 2+ Hypochlorite ClO 2 Carbonate CO 3 Zinc Zn 2+ Chlorite ClO 2 2 Chromate CrO 4 Iron(II) Fe 2+ Chlorate ClO 3 2 Dichromate Cr 2 O 7 Tin(II) Sn 2+ Perchlorate ClO 4 2- Oxalate C 2 O 4 Lead(II) Pb 2+ Cyanide CN Cobalt(II) Co 2+ Cyanate OCN Nickel(II) Ni 2+ Thiocyanate SCN Manganese(II) Mn 2+ Nitrate NO 3 Nitrite NO 2 Permanganate MnO 4 Hydrogen Carbonate HCO 3 Acetate C 2 H 3 O Aluminum Al 3+ Nitride N 3 Chromium(III) Cr 3+ Phosphide P 3 Iron(III) Fe 3+ 3 Phosphate PO 4 Commonly used Acids Hydrochloric acid HCl Sulfuric acid H 2 SO 4 Nitric acid HNO 3 Acetic acid HC 2 H 3 O 2 Phosphoric acid H 3 PO 4

6 SOLUBILITY RULES 1. Salts containing Na +, K + and NH 4 + are always soluble. 2. Salts containing nitrate (NO 3 ), chlorate (ClO3 ), perchlorate (ClO4 ) and acetate (C 2 H 3 O 2 ) are always soluble. 3. All chlorides (Cl ), bromides (Br ), and iodides (I ) are soluble except for those of Ag +, Pb 2+ and Hg 2 2+, which are insoluble. 4. All sulfates (SO 4 2 ) are soluble except for those of Sr 2+, Ba 2+, Hg2 2+, Hg 2+ and Pb 2+, which are insoluble. The sulfate salts of Ca 2+ and Ag + are moderately soluble. 5. All hydroxides (OH ) are insoluble except for those of the alkali metals, which are soluble, and the hydroxides of Ca 2+, Ba 2+ and Sr 2+, which are moderately soluble. 6. All sulfites (SO 3 2 ), carbonates (CO3 2 ), chromates (CrO4 2 ) and phosphates (PO 4 3 ) are insoluble except for those of NH4 + and the alkali metals. 7. All sulfides (S 2 ) are insoluble except for those of NH 4 +, the alkali metals and the alkaline earth metals, which are soluble.

7 MATERIALS Bunsen burner 0.1 M Solutions of: Ring stand and ring Dilute sulfuric acid Clay triangle Hydrogen peroxide Crucible Sodium bicarbonate Crucible tongs Silver nitrate 6 test tubes Sodium hydroxide Test tube rack Sodium sulfate Evaporating dish Copper(II) sulfate Magnesium ribbon Sodium chloride Barium chloride Potassium bromide PROCEDURE Perform all reactions at your workspace. Do not perform reactions at the main dispensing hood. COMBINATION REACTION 1. Reaction of magnesium with oxygen and nitrogen: you may have performed this procedure in a previous experiment ask your instructor if you may therefore skip to #2. Set up a Bunsen burner, ring stand and ring, and a clay triangle. Support a dry crucible (dry is more important than clean!) on the triangle. Place about 1 inch of loosely coiled magnesium metal in the crucible, and heat the bottom of the crucible until the metal burns. DO NOT LOOK DIRECTLY AT THE BURNING MAGNESIUM Once the magnesium has completely burned, allow the crucible to cool to room temperature. Gently stir the ash with a glass rod and describe the appearance of the ash in your laboratory notebook. DECOMPOSITION REACTION 2. Decomposition of hydrogen peroxide This reaction is described in the background. Without a catalyst this reaction proceeds extremely slowly. A catalyst is a substance that speeds up a reaction without being consumed by the reaction. You will use a catalyst to speed up the decomposition of hydrogen peroxide so that the bubbles of oxygen being produced can be seen.

8 Place about 2mL of 3% hydrogen peroxide in a test tube. Add one drop of iron(iii) chloride solution. Stir with a glass rod. Note the rate of evolution of bubbles of oxygen. Write careful observations into your laboratory notebook. NEUTRALIZATION REACTION Check each reactant with litmus paper before mixing, and check each mixture with litmus paper after the reaction has taken place. The proper technique is to dip a clean stirring rod into the solution and then touch the stirring rod to the litmus paper. Write the results of the tests with litmus paper into your laboratory notebook. 3. Sodium hydroxide and dilute sulfuric acid Obtain 2mL of aqueous sodium hydroxide in a small test tube. Obtain 2mL of dilute sulfuric acid in another test tube. Do not measure these volumes, use 40 drops or approximately 2mL as shown by your instructor. Note the temperature of each solution by touching the outside of the tube. Carefully add the 2mL of dilute sulfuric acid to the 2mL of the sodium hydroxide. Stir with a glass rod. Note any temperature changes by touching the test tube. PRECIPITATION REACTIONS Set up the small test tube rack containing six clean test tubes. Number each tube either with a pencil on its label spot, or with a grease pencil. As you record the results of each reaction in your laboratory notebook, be sure to include the test tube number. Add about 2mL of each of the two reagents. It is not necessary to measure the volumes of the reagents. Stir each mixture with a clean stirring rod. 4. Silver nitrate and potassium bromide 5. Sodium sulfate and barium chloride 6. Copper(II) sulfate and sodium hydroxide 7. Sodium bicarbonate and barium chloride 8. Copper(II) sulfate and sodium chloride After you have run each reaction, write down all observations including colors of original reagent solutions, amount of precipitate formed and colors of precipitates. GAS FORMING REACTION 9. Sodium bicarbonate and dilute sulfuric acid Pour about 2mL of aqueous sodium bicarbonate into an evaporating dish, not a test tube. Add about 2mL of dilute sulfuric acid. Write all observations into your laboratory notebook.

9 Name: Date: PRE-LAB QUESTIONS Write a balanced chemical equation from the word description given. Be sure to use the solubility rules to determine the physical states of the products. 1. Sodium phosphate solution will react with calcium chloride solution to form a white precipitate. For this question, check your equation with the one shown in the background. 2 A yellow precipitate forms when solutions of silver nitrate and sodium phosphate are mixed. 3. Bubbles of gas are produced when a solution of potassium hydrogen carbonate (potassium bicarbonate) reacts with a sulfuric acid solution. 4. A white precipitate appears and the solution gets warm when solutions of barium hydroxide and phosphoric acid are mixed.

10 Name: Date: CHEMICAL REACTIONS AND EQUATIONS Your data should show all of your observations for each reaction. Write down the reactants, colors of compounds, etc

11 RESULTS AND POST-LAB QUESTIONS Following the examples from the background and pre-lab questions, write a balanced equation for each of the 9 reactions that you performed. Also write the names of all reactants and products present in the balanced equations

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