Aqueous Chemical Reactions

Transcription

1 Name: Date: Lab Partners: Lab section: Aqueous Chemical Reactions The purpose of this lab is to introduce you to three major categories of reactions that occur in aqueous solutions: precipitation reactions, acid-base reactions, and oxidation-reduction or redox reactions. Precipitation reactions are those in which water soluble compounds react to form new compounds, one of which is insoluble in water. The precipitation reactions in this experiment are all examples of double displacement reactions in which there is an exchange of parts between two compounds. For example, when an aqueous solution of lead nitrate is added to an aqueous solution of potassium iodide, a yellow precipitate of lead iodide is formed: Pb (NO 3 ) 2(aq) + 2KI (aq) PbI 2(s) + 2KNO 3(aq) Acid-base reactions involve the transfer of a proton (H + ) from an acid to a base. In aqueous solutions acids increase the hydrogen ion (H + ) concentration and bases increase the hydroxide ion (OH ) concentration. When an acid and a base react in an aqueous solution the H + and OH ions combine to form water, and the cation of the base and anion of the acid form a salt. For example, the reaction of hydrochloric acid and sodium hydroxide produces common table salt (NaCl) and water as in the following reaction: HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) In reactions with equal molar amounts of the acid and base, acid-base neutralization occurs and only a salt and water remain in solution. One special type of acid-base reaction that is common in aqueous chemistry involves carbonates, carbonic acid and carbon dioxide. In this case, an acid mixes with a carbonate (which is a base) and the result of the reaction is a salt, water and carbon dioxide gas. This is really just a two step process where the first reaction transfers hydrogen ions to the carbonate to form carbonic acid and the second step is the decomposition of carbonic acid into CO 2 and water. For example: Neutralization reaction: Na 2 CO 3(aq) + 2HCl (aq) 2NaCl (aq) + H 2 CO 3(aq) Decomposition reaction: H 2 CO 3(aq) H 2 O (l) + CO 2(g) Net Reaction: Na 2 CO 3(aq) + 2HCl (aq) 2NaCl (aq) + H 2 O (l) + CO 2(g) This reaction is very common in rivers, lakes and the oceans which owe their ph to the presence of carbonates. Also, the ph of blood is largely controlled by carbonates. It is important to note that carbonic acid is inherently unstable and it cannot be isolated as a pure chemical. CHM 113, Cahill, Aqueous Chemical Reactions 1

2 Oxidation-reduction (redox) reactions involve the transfer of electrons from one substance to another. The atom that is losing electrons is oxidized while the atom gaining electrons is reduced. To keep track of electrons in redox reactions, oxidation numbers are assigned to reactants and products. If the oxidation state of any atom in a reaction changes during the course of the reaction, then the reaction is a redox reaction. It is important to mention that neither precipitation reactions nor acid-base reactions are redox reactions; therefore, no oxidation states change in precipitation or acid-base reactions. Examples of oxidation reactions include: 2Fe (s) + 3/2 O 2(g) Fe 2 O 3(s) CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (g) H 2 SO 3(aq) + H 2 O 2(aq) H 2 SO 4(aq) + H 2 O (l) In each of the reactions above, the first reactant written is oxidized and the second reactant is reduced when the reaction proceeds. Complete and Net Ionic Equations: Molecular equations are written as though all the reacting species exist as whole units in solution, but this doesn t actually describe what happens in the solution. Ionic compounds dissolved in water dissociate into their component cations and anions; therefore, complete ionic equations are written to show the dissolved species as free ions. The dissolved ions that remain in solution aren t actually involved in the chemical reaction and are called spectator ions. Net ionic reactions show only the species that actually take part in the reaction and do not include the spectator ions. For example, let s look at the reaction when an aqueous solution of barium chloride is added to an aqueous solution of sodium sulfate. For this double-decomposition reaction, the products are barium sulfate and sodium chloride. Therefore, the molecular equation is: BaCl 2(aq) + Na 2 SO 4(aq) BaSO 4(s) + NaCl (aq) Using the solubility rules in Table 1, we determine that barium sulfate is insoluble and sodium chloride is soluble in aqueous solution. The complete ionic equation is: Ba 2+ (aq) + 2Cl (aq) + 2Na + (aq) + SO 4 2 (aq) BaSO 4 (s) + 2Cl (aq) + 2Na + (aq) And finally, canceling out the spectator ions on both sides of the equation gives us the following net ionic equation: Ba 2+ (aq) + SO 4 2 (aq) BaSO 4(s) To complete the equation, always check that the charges and number of atoms balance. CHM 113, Cahill, Aqueous Chemical Reactions 2

3 For an example of an acid-base net ionic equation, let s look at the reaction between NaOH and HCl. Because both the acid and the base are strong electrolytes, they are completely ionized in solution. Using the solubility rules in Table 1, the product NaCl is soluble in aqueous solution and the complete ionic equation is: Na + (aq) + OH (aq) + H + (aq) + Cl (aq) Na + (aq) + Cl (aq) + H 2 O (l) Canceling out the spectator ions, the net ionic equation is: H + (aq) + OH (aq) H 2 O (l) The net ionic equation for the vast majority of acid-base reactions simply involves the formation of water from H + and OH - as shown above. Some acid-base reactions will form H 2 S instead of H 2 O. Other exceptions are quite rare. PART A: PRECIPITATION REACTIONS In order to predict if a precipitate will form when two solutions are mixed, you will need to determine the solubility of the products. A substance is soluble if a fair amount of it visibly dissolves when added to water, and is considered either slightly soluble or insoluble if little or none of it dissolves. There are some useful rules in determining the solubility of a compound: all ionic compounds containing alkali metal cations; the ammonium ion; nitrate; bicarbonate; and chlorate ions are always soluble. Table 1 summarizes what ions tend to make soluble compounds. Most other ionic substances that contain ions NOT listed in Table 1 tend to be insoluble. You will be required to observe different reactions and from Table 1 make predictions whether or not a precipitate will form. Table 1. The following ions tend to form soluble compounds. Major exceptions Ion: (These will make precipitates with the ions listed to the left.) Anions: NO 3 No exceptions ClO 3 and ClO 4 No exceptions Acetate (CH 3 COO ) No major exceptions Cl Ag +, Pb 2+ 2+, Hg 2 Br Ag +, Pb 2+ 2+, Hg 2 I Ag +, Pb 2+ 2+, Hg 2 2 SO 4 Sr 2+, Ba 2+, Pb 2+ 2+, Hg 2 Cations: Alkali metals (Li +, Na +, K +, etc.) Ammonium (NH + 4 ) Alkali earth metals (Ca 2+, Sr 2+, Ba 2+ ) No major exceptions No major exceptions CO 2 3, PO 3 4, F, many oxyanions of metals CHM 113, Cahill, Aqueous Chemical Reactions 3

4 PROCEDURE: PART A 1. Using the solubility rules listed in Table 1 determine if a precipitate will form for each of the double displacement reactions in the data table. If no precipitation reaction will occur, write NR on the data sheet. Make the predictions before going to the reagent stations! 2. Next, you will test your predictions at each of the reaction stations. Combine 10 drops of each reagent into a small test tube. Always add reagents to the test tube by dropping from above the tube! Do not insert the dropper into the test tube as this will contaminate the reagents when the dropper is placed back into the bottle. Place a clean stopper in the test tube and invert to mix the contents. 3. Record your observations. Did a precipitate form? What color was it? 4. Write the balanced net ionic equation for each of the chemical reactions. Include the signs for (aq), (s), and (l). 5. Dispose of all the chemical waste in the designated container and thoroughly clean your test tubes with DI water. PART B: ACID-BASE REACTIONS Acid-base reactions can result in the formation of water and a salt, gas formation, or precipitation products. In this part of the lab you will use the solubility guidelines in Table 1, and the gases in Table 2 to determine what products will form in each of the reactions in the data table. Table 2 Gas CO 2 NO 2 H 2 S NH 3 Properties Colorless; odorless Brown, pungent odor (toxic) Colorless; rotten egg odor (toxic) Colorless, pungent odor PROCEDURE PART B 1. For each of the reactions listed in the data table, combine 10 drops of each reagent into a small test tube. Always add reagents to the test tube by dropping from above the tube! Do not insert the dropper into the test tube as this will contaminate the reagents when the dropper is placed back into the bottle. You will be able to see if a gas is formed by the evolution of bubbles. Place a clean stopper in the test tube and invert to mix the contents, if necessary. 2. For the Na 2 CO 3 + HCl station, add a small amount of solid Na 2 CO 3 to the test tube and then add 10 drops of the 0.1M HCl. 3. The HCl + ZnS station is set up in the fume hood. Add a small amount of the solid ZnS to the test tube and then add 2 drops of the 12M HCl. When you have finished testing this reaction, leave your test tube in the hood for this station only. 4. Record your observations. Did a precipitate form? What color was it? Did a gas form? Was there a strong odor produced? 5. Write the balanced net ionic equation for each of the chemical reactions. Include the signs for (aq), (s), and (l). 6. Dispose of all the chemical waste in the designated container and thoroughly clean your test tubes with DI water. CHM 113, Cahill, Aqueous Chemical Reactions 4

5 PART C: OXIDATION-REDUCTION REACTIONS Oxidation-reduction reactions involve the transfer of electrons between atoms. Oxidation reactions involve the loss of an electron from the atom while reduction reactions involve the gain of electrons by an atom. These reactions always come in pairs since the electrons lost by one atom need to be gained by another atom, hence these reactions are often referred to as redox reactions. The surest means to identify a redox reaction is that two or more atoms in the reaction change their oxidation number (also called oxidation state). Oxidation numbers are assigned using the following rules. These rules are in order of precedence so that high number rules take priority over lower rule numbers except for the last rule that must always be true. Oxidation number rules: 1) The oxidation number for pure elements is zero. a. A pure element is defined as how you find it in its native state. 2) The oxidation number for monatomic ions is equal to the charge on the ion. a. For example, all alkali metal ions have an oxidation number of +1, alkaline earth metal ions have an oxidation number of +2, etc. b. Keep in mind that ionic substances may consist of many ions electrostatically held together. The ions in ionic substances need to be treated separately. For example, NaNO 3 is an ionic substance consisting of an Na + ion and an NO 3 ion stuck together. The Na + ion is a mono-atomic ion and has an oxidation state of +1. The NO 3 ion is a poly-atomic ion and is not covered by this rule. 3) Oxygen has an oxidation number of 2 in most compounds except peroxides (e.g. H 2 O 2 ) when it is 1. 4) Hydrogen has an oxidation number of +1 when bonded to non-metal atoms and an oxidation state of 1 when it is associated with metals. 5) Halogens tend to have and oxidation number of 1 in compounds except when bonded to oxygen. a. The halogens in these oxyanions (e.g. ClO 4, ClO, BrO 3 ) can have positive oxidation numbers up +7. 6) The sum of oxidation numbers for all atoms in a molecule must equal the charge on the molecule. a. If the molecule is neutral, then the oxidation numbers must sum to zero. b. If the molecule is an ion (e.g. NO 3 ; SO 4 2 ; PO 4 3 ; etc.) then the sum of the oxidation numbers of all the atoms in the molecule must sum to the charge of the molecule. It is worth noting that precipitation and acid-base reactions are not redox reactions. The ions are re-arranged in these reactions or water is formed, but no atoms change their oxidation number. In the data section of the laboratory report, assign oxidation numbers to the compounds listed. Remember that the oxidation number calculations are conducted on a per atom basis. Unfortunately, the nomenclature of redox reactions is a little messy. Chemicals are often classified as oxidizers or oxidation agents. These chemicals tend to pull electrons away from other atoms. Therefore oxidizer agents oxidize other chemicals, so that is their function. The irony in the nomenclature is that the oxidizing agent is taking electrons from other atoms and hence it is actually reduced during the course of the reaction. Oxidizers are particularly hazardous since they tend to be reactive and promote fire and other combustion processes. Oxidizers are commonly found in explosives, rocket fuels, etc. Their reactive nature also makes them degrade CHM 113, Cahill, Aqueous Chemical Reactions 5

6 other chemicals that they come in contact with; hence, they are also used as cleaning agents and disinfectants. In this lab, you will test several compounds to determine which ones are oxidation agents. In all cases, you will test the compound with iron (II) sulfate (or ferrous sulfate). Iron forms two main states (other than elemental) and they are Fe 2+ and Fe 3+. The Fe 3+ state is the more stable and preferred state. Therefore, the Fe 2+ of the iron (II) sulfate would like to give up another electron to reach this more stable state. Therefore, most oxidizers can quickly convert the Fe 2+ (which is mostly clear) to Fe 3+ which is a rusty color. The Fe 3+ then precipitates out as Fe 2 O 3 which is rust. Therefore, most oxidizers will turn a solution of iron (II) sulfate from mostly clear to a rusty yellow-brown. PROCEDURE: PART C 1. Obtain 8 small test tubes and put them in a test tube rack. 2. Add 1 ml of 0.01 M FeSO 4 solution to each of the test tubes. 3. Label one test tube as a control (nothing else added) and put it aside. You will compare the other reaction vials to this control test tube. 4. Label the other 7 test tubes by the reagent that you will be testing. The reagents you will be testing are: a. Potassium iodide (KI, 0.1 Molar solution) b. Hydrogen peroxide (H 2 O 2, 3% solution) c. Sodium sulfate (Na 2 SO 4, 0.1 Molar solution) d. Hydrochloric acid (HCl, 0.1 Molar solution) e. Bleach (NaClO) f. Sodium chloride (NaCl, 0.1 Molar solution) g. Potassium bromate (KBrO 3, 0.1 Molar solution) 5. Add 2-3 drops of each of the 7 test reagents to a separate test tube to observe the reactions. Remember, drop from above the test tube to avoid contaminating the reagents! Record your observations. 6. Once you have recorded your initial observations, let the samples sit for 20 to 30 minutes while you perform other laboratory tasks. Some of the reactions may proceed slowly. Record your final observations after the time has elapsed. PART D: PREPARE A SOLUTION The accurate preparation of solutions is a common task in chemistry. Most of the time, the solutions are already prepared for you, but for this part of the lab you will prepare your own. The general procedure is fairly straight forward: a) Pick a volume of solution that you will need. Try not to create excess solution that will have to be thrown away later. Most of the time you will be making 10 to 25 ml of solutions. Obtain a volumetric flask of the desired size. b) Determine the desired concentration of the solution measured in molarity. Most of the time people try to pick round numbers. Keep in mind that some chemicals are not very soluble and it is hard to get high concentrations into solution. c) Calculate the mass of chemical you need to add to your flask to achieve the desired concentration. Remember that: CHM 113, Cahill, Aqueous Chemical Reactions 6

7 Therefore: Molarity = M = mol/l and (molar mass) = g/mol mol = (M L) and mol = g/(molar mass) Combining these two equations you get: g/(molar mass) = (M L) g = (M L) (molar mass) Therefore, the grams of material you need to create a solution with the desired concentration and volume is (M L) (molar mass). Note that the volume is in units of liters, so you will need to convert the flask size into liters. d) Weigh out the desired mass of chemical as closely to the target mass as possible. Record the actual mass measurement in your lab notebook. You will use this actual mass value to calculate the actual concentration of the solution. e) Pour the chemical into your volumetric flask and fill it about half full of your solvent. Swirl the flask until the chemical completely dissolves. Then add enough solvent to bring the volume up to the calibration mark, which is the line etched in the glass in the long neck of the flask. If your chemical does not dissolve in a reasonable amount of time, then add more solvent and bring it to about 2/3 full. You want to keep some empty volume in the round part of the flask so you can swirl the mixture. Once you add solvent up into the neck of the flask, then it is hard to mix without inverting it. f) Cap or seal the volumetric flask tightly and invert the flask about 10 times to mix the solution. Be sure to wear gloves since the caps sometimes leak. g) Transfer the solution into a jar or test tube for storage. Volumetric flasks are not good for long term storage of chemicals; they are really hard to clean if a precipitate forms on the interior of the bulbous part of the flask. h) Label your solution with the chemical name, the concentration, the date prepared and your name. PROCEDURE: PART D 1. Using the solution preparation protocol given above, prepare 25 ml of a 0.1 M CuSO 4 solution in DI water. Fill in the data on the worksheet and show your calculations. 2. Once the solution is prepared, pour it into a clean 13x100 test tube and measure the absorbance of the solution with a visible spectrophotometer. Your instructor will use your measured absorbance value to determine the accuracy of your solution prep. 3. Check to make sure that the spectrophotometer is set to absorb at 650 nm. 4. Auto-zero the instrument by filling a clean test tube ¾ full with DI H 2 O. Make sure the test tube isn t scratched on the surface and wipe the outside with a tissue before inserting. 5. Close the cover and press the measure blank or 0 Abs button to zero the instrument. 6. Remove the blank and insert your solution. Again, make sure the tube is clean and scratch free. CHM 113, Cahill, Aqueous Chemical Reactions 7

8 7. Close the cover and record the measured absorbance on the data sheet. You do not need to press a button, the absorbance will automatically display on the screen. You do not actually use this number, but it is used in the grading of the laboratory report since it proves that the solution prepared was (or was not) at the correct concentration. Clean-up your station and glassware: 1) Dispose of the hazardous waste in the appropriate container. 2) Wash and brush out your test tubes in the soapy water bath in the sink. a. The test tubes that had precipitates or rust in them will not get clean with simple rinsing, so you will need to use a test tube brush to clean them out. b. If you labeled the test tubes with a pen, then rub off the label with a scratchy pad or a paper towel. c. Rinse the test tubes, both inside and outside, with DI water (the plastic faucet). Put them upside down in the test tube in your equipment bin so they are ready for the next class. d. The only test tube you do not need to clean is the HCl + ZnS test tube that should still be in the fume hood. 3) Your equipment bin should contain the following (check box to verify) Small test tube rack Small test tubes (25, the 24 you washed + 1 replacement for the HCl/ZnS one) Rubber stopper (size 00) CHM 113, Cahill, Aqueous Chemical Reactions 8

9 PART A: PRECIPITATE REACTIONS REPORT SHEET Write the balanced net ionic equation for each of the chemical reactions. Include the physical state, namely (aq), (s), or (l), for each chemical species in the net ionic equation. Also make sure that the charge is given for all aqueous ions. If there is no reaction, then write No Reaction in the net ionic equation column. Test # Reacting Compounds Prediction Will a precipitate form? YES or NO Example Pb(NO 3 ) 2 + KI Y 1 AgNO 3 + NaCl Observation Cream colored precipitate Net Ionic Equation Pb 2+ (aq) + 2I (aq) PbI 2(s) 2 BaCl 2 + Na 2 SO 4 3 NH 4 Cl + NaOH 4 Na 2 CO 3 + Ba(OH) 2 5 KI + AgNO 3 6 Na 2 CO 3 + Ca(C 2 H 3 O 2 ) 2 7 CuSO 4 + Zn(NO 3 ) 2 8 Na 2 CO 3 + CaCl 2 9 BaCl 2 + Na 3 PO 4 of 4.5 pts CHM 113, Cahill, Aqueous Chemical Reactions 9

10 PART B: ACID-BASE REACTIONS Write the balanced net ionic equation for each of the chemical reactions. Include the physical state, namely (aq), (s), or (l), for each chemical species in the net ionic equation. If nothing visible occurs, then write no visible reaction under the observations. However, even reactions with no visible reaction require a net ionic equation to be completed since some reaction occurs in each of these tests. Reacting Compounds Observation Net Ionic Equation Example NaOH + HNO 3 reaction No visible H + (aq) + OH (aq) = H 2 O (l) 1 H 2 SO 4 + Na 2 CO 3 2 Ba(OH) 2 + H 2 SO 4 3 HCl + KOH 4 (Solid) Na 2 CO 3 + HCl 5 NaOH + CH 3 COOH (acetic acid) 6 HCl + ZnS (solid) of 3 pts CHM 113, Cahill, Aqueous Chemical Reactions 10

11 PART C: OXIDIZER TESTS Reagents control Initial Observation Final observation (after 20 to 30 minutes) Is the reagent an oxidizer? potassium iodide hydrogen peroxide sodium sulfate hydrochloric acid bleach sodium chloride potassium bromate PART C: OXIDATION NUMBERS Assign the oxidation number of each element in the following compounds and ions. In the following examples, all the elements of the same type (e.g. all the oxygen atoms in the compound) have the same oxidation number, so just give the oxidation number for that element and not the sum of their oxidation numbers. Example #1 K 3 PO 4 K = +1 P = +5 O = 2 Example #2 + NH 4 H = +1 N = 3 HNO 3 H 2 SO 4 Na 2 (Cr 2 O 7 ) KClO 3 Li(MnO 4 ) HSO 3 of 6.5 pts CHM 113, Cahill, Aqueous Chemical Reactions 11

12 PART D: PREPARE A SOLUTION Desired concentration of CuSO 4 solution M Desired volume of solution ml Desired mass of CuSO 4 5 H 2 O (molar mass = g/mol) (Show calculation) g Actual mass of CuSO 4 5 H 2 O Actual concentration of solution (Show calculation) g M Measured absorbance of the solution of 3 pts CHM 113, Cahill, Aqueous Chemical Reactions 12

13 POST LAB QUESTIONS 1. What is the difference between a net ionic equation and a molecular equation? 2. Characterize the following compounds as soluble or insoluble in water: (a) Ca 3 (PO 4 ) 2 (b) Mn(OH) 2 (c) AgClO 3 (d) K 2 S 3. What are the characteristics of an acid-base neutralization reaction? 4. Write the balanced net ionic equations for the following: (a) HBr + NaOH (b) Ba(OH) 2 + H 3 PO 4 6. Arrange the following species in order of increasing oxidation number of the sulfur atom (a) H 2 S (b) elemental sulfur (S 8 ) (c) H 2 SO 4 (d) H 2 SO 3 of 5 pts Total for lab of 25 pts CHM 113, Cahill, Aqueous Chemical Reactions 13