CHM-201 General Chemistry and Laboratory I Laboratory 4. Introduction to Chemical Reactions (based in part on Small Scale Chemistry methodology as described in Chemtrek by Stephen Thompson at Colorado State University) October 16, 2018 Purpose: This laboratory will introduce you to a variety of chemical reactions. You will observe precipitation reactions, acid-base reactions, oxidation-reduction reactions and reactions in which complex ions are formed. You will gather information about known compounds to help you deduce the identity of an unknown compound. Introduction: There are several types of chemical reactions. A fundamental understanding of chemical reactions is necessary to consider the other topics in this course. This laboratory is designed to give you experience with chemical reactions and to observe what is meant by a chemical change. By observing some chemical reactions, you may find it easier to understand what they are. Rather than discuss each reaction as it comes, it is easier when we group reactions into categories. We will use three categories to describe reactions in this laboratory: 1. Precipitation Reactions: reactions that result in the formation of an insoluble (solid) compound. 2. Acid-Base Reactions: reactions that result from the transfer of a hydrogen ion (H + ) from one species to another. 3. Oxidation-Reduction Reactions: reactions that result in the transfer of electrons from one species to another. To state it another way, there is a change in oxidation number for at least two elements in the reaction. Now that we have categorized reactions, we shall discuss each type of reaction in more detail and relate that to what we expect to observe in the laboratory session. 1. Precipitation Reactions In precipitation reactions, two substances, one insoluble (a precipitate) and one soluble (that you probably cannot observe) form in what is called an exchange reaction. In an exchange reaction, the cation from one solution combines with the anion in the second solution to form a precipitate. Exchange reactions have the following general form: C 1A 1(aq) + C 2A 2(aq) C1A2(s) + C2A1(aq) Where C 1 is the cation of solution 1, A 1 is the anion of solution 1, C 2 is the cation of solution 2, and A 2 is the anion of solution 2 As an example, we shall use the reaction of silver nitrate with sodium chloride to form silver chloride, a precipitate, and sodium nitrate. The silver chloride will appear as a white precipitate. The chemical reaction is written below: AgNO 3(aq) + NaCl AgCl(s) + (aq) (Note that C 1 = Ag +, A 1 = NO 3-, C 2 = Na +, A 2 = Cl - in this example) Since AgCl is an insoluble material we might suspect that the combination of AgNO 3 and NaCl would give an insoluble material (You will do this in the experiment. Check it out!). Notice that the other possible product, NaNO 3, would also be formed but remains in solution, giving you no way to observe it. The reason that this is a precipitation reaction is that it forms an observable solid, AgCl, from the combination of two solutions.
Being able write an exchange reaction for two ionic compounds does not necessarily mean that a chemical reaction occurs, however. When you combine sodium chloride, NaCl, and copper(ii) sulfate, CuSO 4, in the laboratory, no precipitate forms. The sodium chloride solution is colorless and the copper sulfate solution is blue. When you combine them, you get a solution that is lighter blue than the original copper sulfate solution. Nothing else seems to happen. No chemical reaction is apparent. We can write an exchange reaction as follows: 2NaCl(aq) + CuSO 4(aq) CuCl 2(aq) + Na 2SO 4(aq), but there is no precipitate (no solid formed). What ions are present on the left side of the equation? Answer: 2 sodium ions, 2 chloride ions, 1 copper ion and 1 sulfate ion. And what ions are present on the right hand side of the equation? Answer: 2 sodium ion, 2 chloride ion, 1 copper ion and 1 sulfate ion. If both sides contain the same ions in the same amounts, there is no chemical reaction. If you correctly write the net ionic equation, you will find that there is no net ionic reaction because everything cancels (prove this to yourself). There is no reaction. Sodium chloride and copper sulfate (the reactants) are both soluble, and copper(ii) chloride and sodium sulfate (the products) are also soluble. In this case, since the same ions are present in solution as reactants and as products, there was no net ionic equation. We have simply mixed two soluble ionic compounds. 2. Acid-Base Reactions We will use a very simple definition of acids and bases in order to understand them better. We will define acids as compounds that can donate hydrogen ions (H + ) to other substances. We can define bases as compounds that consume hydrogen ions (note that when we discuss acids and bases in greater detail in CHM202, the definitions of what we call acids and bases will be expanded). If an acid and a base are combined, they will react with one another such that the acid will donate a hydrogen ion and the base will consume it. A common acid used in the laboratory is hydrochloric acid (HCl). HCl is an acid because it can donate hydrogen ions to other substances. A common base is sodium hydroxide (NaOH) which can consume hydrogen ion since hydroxide ion (OH - ) combines with hydrogen ion to form water (H 2O). The chemical reaction we would write for this chemical change would be: The net ionic equation for this reaction is: HCl(aq) + NaOH(aq) H 2O(l) + NaCl(aq) H + (aq) + OH - (aq) H 2O(l) The combination of hydrogen ion (H + ) and hydroxide ion (OH - ) to form water is a very common net ionic equation for acid-base reactions (it accounts for about 80% of such reactions you will see in this course). This type of reaction can be described as, "an acid and base react to form water and a salt." Does the reaction of HCl with NaOH meet this description? Another common acid-base reaction takes place when an acid reacts with carbonate (CO 3 2- ) or hydrogen carbonate, also known as bicarbonate, (HCO 3- ) to give water and carbon dioxide. An example is given below for the reaction of sulfuric acid with potassium carbonate: H 2SO 4(aq) + K 2CO 3(aq) K 2SO 4(aq) + H 2O(l) + CO 2(g) In this case, carbon dioxide gas is formed in addition to a salt and water. The net ionic equation for this reaction is: 2H + (aq) + CO 3 2- (aq) H 2O(l) + CO 2(g) The two net ionic equations described above should apply to almost all of the acid-base reactions you will observe in this course.
It is more difficult to observe an acid/base reaction than a precipitation reaction. Since there may be no immediately observable differences (for example the HCl-NaOH reaction has the combination of two colorless solutions and the product is also a colorless solution), we need a method of finding out if solutions contain a high concentration of hydrogen ions (acidic) or hydroxide ions (basic). We can do this with indicator paper (litmus paper) and/or solutions of indicators that change color depending on the concentration of H + or OH -. If we use litmus paper we can define if solutions are acids (acids turn blue litmus paper red) or bases (bases turn red litmus paper blue). In general, indicators change color depending on the acidity or basicity of a solution. One example is phenolphthalein, which is clear in acidic solution but turns pink in basic solution. The universal indicator used in this laboratory undergoes a wide range of color changes depending on the concentration of H + ions. When acids and bases react, the ph (acidity/basicity) usually changes, and there may be other indications of a chemical reaction (gas evolution for example). It may be difficult to determine when some acid-base reactions have occurred without using indicators, if the ph is different from those of the two starting solutions, it is likely that an acid-base reaction has occurred. For the HCl-NaOH example, the final reaction mixture (NaCl in water) will be neutral (no change in litmus paper) but only if equal amounts of the starting materials are combined. Why is that the case? If HCl is in excess, the resulting solution will remain acidic. If NaOH is in excess, the resulting solution will remain basic. In this laboratory, you will determine whether some solutions before and after reacting are acids or bases. Be careful to observe any other changes such as gas evolution, since that will help you identify your unknown. 3. Oxidation-Reduction Reactions Oxidation-reduction (redox) reactions are the most complex of the three. In order to describe oxidationreduction reactions we need first to understand oxidation numbers. Oxidation numbers represent the electrical charge or equivalent charge of an atom in a compound or element. The oxidation number of a pure element is always zero. You will need to know how to assign oxidation numbers in order to determine whether a reaction is an oxidation-reduction reaction. We can look at a simple case where the oxidation number is the charge of an ion (remember that a pure element has an oxidation number of zero). If we react sodium metal (an element) with chlorine gas (an element) we will obtain sodium chloride. In this reaction, the charge is equal to the oxidation number. The reaction is as follows with the oxidation number of each material below it: Oxidation # 0 0 +1-1 2Na + Cl 2 2NaCl Any reaction in which there is a change in oxidation numbers is an oxidation-reduction reaction. One atom has to increase its oxidation number (be oxidized), and one atom has to decrease its oxidation number (be reduced). It is possible for the two atoms involve to be of the same element, but we will not see that in this laboratory. In the example above, sodium increases its oxidation number from 0 to +1 whereas chlorine decreases its oxidation number from 0 to -1. Sodium is oxidized and chlorine is reduced. When one atom undergoes an increase in oxidation number, another atom always undergoes a decrease in oxidation number. Oxidation cannot take place without a corresponding reduction and vice versa. Oxidation numbers give us a way to account for the number of electrons present. Sodium atoms and sodium ions do not have the same number of electrons. Neither do chlorine atoms and chloride ions. Oxidation-reduction reactions are reactions accompanied by a change in the number of electrons present in the materials that react. There are electrons transferred from one compound to another. In the example above, sodium transfers electrons to chlorine in order to form sodium chloride.
In the laboratory, you may observe oxidation-reduction reactions because of some observable change, and you may be able to reverse these reactions by supplying electrons from an electrical source such as a 9V battery. If electrons are transferred when you react some materials together, you can reverse this process if you transfer the electrons back. Watch for instances of this in the laboratory. You can also connect a battery to an aqueous solution and convert the water to its constituent elements, hydrogen and oxygen. This is also an oxidation-reduction reaction. 4. Introduction to Complex Ions Finally, you will observe a few cases in which complex ions are formed in solution. Complex ions are ions in which a cation binds to one or more anions or with a molecule having lone pairs of electrons (the topic of lone pairs is covered much later in this course). For this brief introduction to the formation of complex ions in solution, you will observe the ammonia complex ions of silver and copper. The reaction of silver chloride, insoluble in water, with ammonia, results in the silver chloride dissolving. This happens because silver chloride is insoluble but silver ion reacts with ammonia to form a more soluble silver ion complex with ammonia: AgCl(s) + 2NH 3(aq) Ag(NH 3) 2+ (aq) + Cl - (aq) Similarly, blue copper ion in solution turns a much darker blue when combined with ammonia. This is due to a complex formed between copper ion and four ammonia molecules: Cu 2+ (aq) + 4NH 3(aq) Cu(NH 3) 4 2+ (aq) For this course, you do not have to predict when complex ions of this type will form, but you should be aware that this might occur. Procedure: General: Be sure to write all pertinent observations, including those you don't expect. Be sure to keep track of all precipitates, gas formations and information from litmus and indicators. Be careful not to contaminate your droppers or solutions. Poor techniques will lead to poor results, and you may not be able to identify your unknown. NOTE: You should have a graphical chart indicating the most effective order of procedures as given below. It should help clarify the procedures. 1. Precipitation reactions 1.1 Using the grid provided place one drop of AgNO 3 in each of the nine grid positions directly below it. 1.2 Repeat procedure 1.1 for each of the following: Na 2CO 3, Ba(NO 3) 2, Pb(NO 3) 2, Zn(NO 3) 2 and unknown. 1.3 Add one drop of the first compound in the left column (HCl) to the drop of silver nitrate (AgNO 3) already there. Be especially careful to keep the dropper from touching the other chemical. You may use a clean toothpick to mix the chemicals, if necessary. 1.4 Write your observations (precipitation, gas formation, color change, etc.) for any changes you observe. If there was no apparent change, write NC for no change. Use the similar chart included in your lab report. 1.5 Repeat 1.3 and 1.4 with hydrochloric acid across the top row for all the compounds on the left side under Precipitation Reactions. 1.6 Place one drop of hydrochloric acid to each of the remaining four columns under Acid/Base Reactions. 1.7 Repeat 1.3 1.6 for each of the compounds in the left column (H 2SO 4, NH 4Cl, NH 3, NaOH, Na 2CO 3, Na 2HPO 4, Na 2SO 4 and NaCl, but not NH 3) one at a time, recording your observations
as you complete each addition. If you make a mistake, dab up the incorrect solutions and reapply. You only need one drop of each compound to make an observation. NOTE: Ammonia is a volatile solution and its fumes will affect some of the other reactions. DO NOT place any NH 3 on the sheet until all the other reactions have been observed, including Acid-base reactions. You will come back to 1.7 then. 1.8 Observe the pattern from the precipitation reactions and the acid-base reactions. Is there one pattern that is identical to that of the unknown? Use the information you have gathered to determine the identity of the unknown. (Do at home not in the lab.) 1.9 Write the number of the unknown and its identity in the space provided. The unknown is one of the five compounds across the top of the chart on the Precipitation Reactions side of the chart. (Do at home not in the lab.) 1.10 Write down your reasoning in the space provided. Your explanation need not be overly detailed or even complete, but you must articulate it. No reasoning no credit! (Do at home not in the lab.) 2. Acid-base reactions. 2.1 There should now be a drop of each solution from the left column on each space under each column on the Acid/Base Reactions side of the sheet. 2.2 Test each of the solutions with both red and blue litmus paper in the vertical columns provided. If there is a color change, record the new color (P = Pink, B = Blue). If the litmus paper does not change color, mark NC for no change. 2.3 Add one drop of phenolphthalein to each solution in the column so labeled and record the result as P for pink, or NC for no change. 2.4 Add one drop of the universal indicator to each solution in the column so labeled and record the color change (if any). NC for no change. 2.5 Now add one drop of NH 3 solution across the bottom row as in Procedure 1.7 and record your observations. 2.6 Repeat 2.2 2.4, testing the newly added ammonia solution with the indicators. 3. Complex ion formation 3.1 Place one drop of copper sulfate solution (CuSO 4) in each of the unused ovals (marked CuSO 4 and CuSO 4 + NH 3). Add a few drops of ammonia to the appropriate oval. What happens? Does the color change. Record your observations. 3.2 Rinse your acetate sheet with water over a large beaker and transfer the waste to the waste collection vessel at the instructor's desk. Wipe down your acetate sheet with a paper towel and discard the entire sheet. Be sure you have recorded all your results before you clean your sheet! 4. Oxidation-Reduction Reactions 4.1 Using the spot plate wells, add 10 drops of copper sulfate solution to the first well and 10 drops of zinc nitrate to the second well. 4.2 Place a small piece of zinc metal into the copper sulfate solution and a small piece of copper metal into the zinc nitrate solution. Let the metals sit in the solutions for 2 minutes. Remove the two metals and record your observations.
4.3 Connect the zinc metal that was dipped in the CuSO 4 solution to the positive post of a 9V battery using an alligator clip. Connect the graphite pencil lead to the negative post and place both solids into the copper sulfate solution without touching the solids together. Take the solids out of the solution after 1 minute. Record your observations. 4.4 Repeat 4.3 connecting the copper metal to the negative post and the graphite pencil lead to the positive post and place both solids in the zinc nitrate solution without touching. Take the solids out of the solution after 1 minute and record your observations. 4.5 See me to help with the following procedure. I will perform the preparation for your group to complete this part in the hood for safety. 4.5.1 Pierce a disposable pipette bulb with two needles, making sure the needles do not touch each other. 4.5.2 Fill the pipette with 1M sulfuric acid solution. 4.5.3 Place the pipette in an small Erlenmeyer flask with the bulb out of the flask. 4.5.4 Connect one needle to the positive terminal and one to the negative terminal of a 9V battery and observe what happens. 4.5.5 When the bulb is full of gas, ignite the gaseous mixture as instructed. 4.5.6 Record your observations.