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Introduction to Chemical Reactions Results (You may incorporate your group observations or copy them into these tables): Chemical Reactivity – Precipitation Reactions Acid/Base Reactions Parts 1 & 2 (20 pts) AgNO3 Na2CO3 Ba(NO3)2 Pb(NO3)2 Zn(NO3)2 Unknown Blue Litmus Red Litmus Phth. Indicator Universal Indicator HCl H2SO4 NH4Cl NaOH Na2CO3 Na2HPO4 Na2SO4 NaCl NH3 CuSO4 CuSO4 + NH3 Identity of Unknown: (10 pts) Unknown number:Identity: Brief statement of reasoning (Required): Observations: (10 pts) With no NH3 With NH3 CuSO4 Oxidation-Reduction reactions: (10 pts) Observations Zn + CuSO4 (No current) Zn + CuSO4 (9V – positive terminal) graphite (9V – negative terminal) Cu + Zn(NO3)2 (No Current) Cu + Zn(NO3)2 (9V – negative terminal) graphite (9V – positive terminal) H2SO4 (9V) H2SO4 (ignition of gas) Questions: (50 pts) 1. This experiment was run on a small scale to minimize chemical waste. The alternative method is to use approximately 1 mL of each reagent in a test tube to observe the reactions, test for acidity, etc (8 pts) 1.1 How many drops of reagents did you use in Parts 1 & 2? 1.2 Assuming that 15 drops of reagent equals 1 mL, what is the total volume of reagent you used in Parts 1 & 2? 1.3 What volume of reagent would you have used if you used the alternative method described above? 1.4 Assuming seven CHM201 classes per year and eight lab groups for each class, compute the total reduction of waste accomplished by using this small-scale method. 2. In your observations, you described what happened when you placed zinc metal in the copper sulfate solution and what happened when you attached it to the battery. (8 pts) 2.1 Write the equation describing the chemical change that occurred when no battery was used. 2.2 Write the equation describing the chemical change that occurred when the 9V battery was used. 3. You also described what happened when you placed the copper metal in the zinc nitrate solution with and without the battery. (6 pts) 3.1 Write the equation describing the chemical change that occurred when no battery was used. 3.2 Write the equation describing the chemical change that occurred when the 9V battery was used. 4. Categorize the reagents used in Parts 1 & 2 in the table below: (10 pts) Reagent Acid Base Neither HCl H2SO4 NH4Cl NH3 NaOH Na2CO3 Na2HPO4 Na2SO4 NaCl 5. Write the full ionic and net ionic equations for the following: (8 pts) 5.1 HCl/AgNO3 full ionic equation: net ionic equation: 5.2 HCl/Na2CO3 full ionic equation: net ionic equation: 5.3 Ba(NO3)2 /Na2SO4 full ionic equation: net ionic equation: 5.4 Ba(NO3)2 /NaCl full ionic equation: net ionic equation: 6. Write the chemical equation for the decomposition of sulfuric acid, H2SO4 solution, using the 9V battery (note that sulfate ion is not involved in this reaction). Indicate the oxidation numbers for all atoms that change during the reaction and indicate which compounds are oxidized and which are reduced. (10 pts) N= NO ronne Chemical Reactivity — Precipitation Reactions Acid/Base Reactions Unknown | Blue Red Phth. Universal Parts 1&2 | AgNO, | NayCO, | Ba(NOy): | Pb(NOs: | Zn(NOs): | #3 Litmus | Litmus __| Indicator | Indicator oe ory Rab NE Bas ne [I p nve-p | me fe le a et I EO a on ne NC Ye No Vo Be ip we [om a a Yolaw | VC ux Fes an Mo 2 | B p-te doc. Oyen Na:COs Titi ne os Sas a Broke a B-pe | B pe Sor Na;HPO, Bn nw Shp ap fon, He. BRe| R p yee NassO, [Bw | AC Pa fo il % Bye | Me np | Ne eee NaCl 3 Ne fe ee oe Vo Boxe |ne -p | HO ore NH, Ne #e ft T0 = ne Ne, RB p So Observations: Part 3 With no NH; With NH; 3 Cus0, [EC Bie, Som wo blz Oxidation-Reduction reactions: Part 4 Observations Zn + CuSOy (No current) Black | Oz : Zn + CuSOq (9V — positive terminal) R pany Cranse AN aol ig blak LZ) Wa OI Graphite (9V — negative terminal) oo STUN Glog SweMy lshs (US [hig Nowd? ZN. \ Cu + Zn(NO;): (9V — negative terminal) F55oWy Who SUNN | Pagemting Sen Yd Gratin cous NAW, S220 | odacked deoocd S qin graphite (9V — positive terminal) Vagieny WRS EC Chasey In @0( po bad | cing H:S04 (9V) sezfied) | Ulogomting 8 Pre iq verkions H2S0; (ignition of gas) | PoRRAD we us ) Wee we | hed d WE feo Vogl to Jos. Introduction to Chemical Reactions CHM201 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) 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: C1A1(aq) + C2A2(aq) → C1A2(s) + C2A1(aq) Where C1 is the cation of solution 1, A1 is the anion of solution 1, C2 is the cation of solution 2, and A2 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: AgNO3(aq) + NaCl → AgCl(s) + (aq) (Note that C1 = Ag +, A1 = NO3 -, C2 = Na +, A2 = Cl - in this example) Since AgCl is an insoluble material we might suspect that the combination of AgNO3 and NaCl would give an insoluble material (You will do this in the experiment. Check it out!). Notice that the other possible product, NaNO3, 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, CuSO4, 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) + CuSO4(aq) → CuCl2(aq) + Na2SO4(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 (H2O). The chemical reaction we would write for this chemical change would be: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) The net ionic equation for this reaction is: H+(aq) + OH-(aq) → H2O(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 (CO3 2-) or hydrogen carbonate, also known as bicarbonate, (HCO3-) to give water and carbon dioxide. An example is given below for the reaction of sulfuric acid with potassium carbonate: H2SO4(aq) + K2CO3(aq) → K2SO4(aq) + H2O(l) + CO2(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) + CO3 2-(aq) → H2O(l) + CO2(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.