Copper Cycle Background

The Copper Cycle Introduction Many aspects of our lives involve chemical reactions—from the batteries that power our cars and cell phones to the thousands of processes occurring within our bodies. We cannot even begin to identify the millions of chemical reactions occurring around us all the time; yet, most of these reactions can be classified into one of six types of chemical reactions: double replacement (precipitation), single replacement, acid/base neutralization, combination, decomposition, and combustion. While you should already be familiar with these reactions, those reactions that require “rules” have been included in this background as a quick review. A review of acid/base chemistry has also been provided along with an explanation of how to determine whether a solution is basic or acidic using red and blue litmus paper. Aqueous Solutions(aq) Many reactions occur in an aqueous environment (i.e., in a solution where ions and compounds are dissolved in water). When we indicate that a reactant or product has the physical state (aq), we mean the substance is dissolved in water. When an ionic compound is in aqueous solution, the individual ions are present in solution; for example, NaCl(aq) exists as Na+ and Cl– ions moving around in water. Acids and Bases In general, acids are compounds that produce hydrogen ions (H+), also called protons. Acids can also be defined as substances that produce hydronium ions (H3O+), which is a hydrogen ion combined with a water molecule: H+(aq) + H2O(l)  H3O+(aq). The two equations below both represent the ionization of hydrochloric acid, HCl(aq): HCl(aq)  H+(aq) + Cl–(aq) HCl(aq) + H2O(l)  H3O+(aq) + Cl–(aq) Acids are usually easy to recognize since their formulas start with H—e.g. HCl(aq), HNO3(aq), and H2SO4(aq) are all acids. Note that the physical state aqueous, (aq), must be included to distinguish an acid from other forms of a substance. For example, the formula “HCl” can also be used for hydrogen chloride gas, HCl(g), so to indicate hydrochloric acid, one must specify HCl(aq). For now, bases are compounds that produce hydroxide ions (OH–) when dissolved in water. The dissociation of sodium hydroxide, NaOH, is shown below: NaOH(aq)  Na+(aq) + OH–(aq) Acid-Base Neutralization Reactions In an acid-base neutralization reaction, an acid reacts with a base to produce water and a salt (an ionic compound): HCl(aq) + NaOH(aq)  acid base H2O(l) + NaCl(aq) water salt Because water is always formed, acids will always react with bases, regardless of whether the salt is soluble or insoluble. page 1 of 6 When carrying out an acid-base neutralization reaction in the laboratory, we observe that most solutions of acids and bases are colorless, and the resulting water and soluble salt solutions are also colorless. Thus, it is impossible to monitor the progress of an acid-base reaction based solely on the appearance of the solutions. To help us monitor acid-base reactions, we use litmus paper to determine if a solution is acidic or basic. Red Litmus Paper Blue Litmus Paper Before adding a drop of solution After adding a drop of solution A drop of basic solution on red litmus turns it blue A drop of an acidic solution on blue litmus turns it red Figure 1 Litmus paper changes color depending on the presence of H3O+ or OH– ions in the substance being tested. Blue litmus paper turns red in acidic solutions containing H3O+ ions, and red litmus paper turns blue in basic solutions containing OH– ions. Solubility Rules Many ionic compounds are soluble—i.e., they dissolve in water. Others generally do not dissolve in water and are considered insoluble. To determine if an ionic compound is soluble—i.e., will dissolve—in water, we use the Solubility Rules: Solubility Rules for Ionic Compounds in Water The compound is SOLUBLE if it has: + + + + 1. Li , Na , K , or NH4 ion (ALWAYS!) – – – 2. C2H3O2 , NO3 , ClO4 – – – 3. Cl , Br , or I , except compounds + +2 +2 with Ag , Pb , and Hg2 are insoluble 2- 4. SO4 except compounds with Ag2SO4, CaSO4, SrSO4, BaSO4, PbSO4, and Hg2SO4 are insoluble The compound is INSOLUBLE if it has: 2– 2– 3– 5. CO3 , CrO4 , PO4 , except compounds + + + + with Li , Na , K , NH4 are soluble 2– + + + 6. S , except compounds with Li , Na , K , + +2 +2 +2 NH4 , Ca , Sr , Ba are soluble – 7. Hydroxide ion, OH , except compounds + + + + with Li , Na , K , NH4 are soluble The Solubility Rules indicate which compounds are soluble, and thus are represented as aqueous: e.g., KI(aq), BaCl2(aq), NaOH(aq), etc. The Solubility Rules also indicate which compounds are insoluble—i.e., do not dissolve in water and remain as solids: e.g. BaSO4(s), AgCl(s), CaCO3(s), etc. Precipitation Reaction For example, consider the reaction between aqueous lead(II) nitrate with aqueous potassium bromide, as shown below: page 2 of 6 Pb(NO3)2(aq) + KBr(aq)  PbBr2 + KNO3 Note that the chemical formulas for the products formed are based on their charges, not how they appear on the reactant side of the chemical equation. Based on Solubility Rules #4 and #1, we find that PbBr2 is insoluble and KNO3 is soluble. Thus, the complete, balanced equation is: Pb(NO3)2(aq) + 2 KBr(aq)  PbBr2(s) + 2 KNO3(aq) We can cancel the spectator ions from the ionic equation and write the net ionic equation: Pb2+(aq) + 2 Br -(aq)  PbBr2(s) This reaction produces a cloudy mixture with small particles of the solid suspended in the solution. When enough solid has formed, it will begin to settle at the bottom of the beaker. Thus, a clear solution becoming cloudy when another solution is added is often taken as experimental evidence of a solid or precipitate forming. Acids and Bases In general, acids are compounds that produce hydrogen ions (H+), also called protons. Acids can also be defined as substances that produce hydronium ions (H3O+), which is a hydrogen ion combined with a water molecule: H+(aq) + H2O(l)  H3O+(aq). The two equations below both represent the ionization of hydrochloric acid, HCl(aq): HCl(aq)  H+(aq) + Cl–(aq) HCl(aq) + H2O(l)  H3O+(aq) + Cl–(aq) Acids are usually easy to recognize since their formulas start with H—e.g. HCl(aq), HNO3(aq), and H2SO4(aq) are all acids. Note that the physical state aqueous, (aq), must be included to distinguish an acid from other forms of a substance. For example, the formula “HCl” can also be used for hydrogen chloride gas, HCl(g), so to indicate hydrochloric acid, one must specify HCl(aq). For now, bases are compounds that produce hydroxide ions (OH–) when dissolved in water. The dissociation of sodium hydroxide, NaOH, is shown below: NaOH(aq)  Na+(aq) + OH–(aq) Single Replacement: A + BC  AC + B In this type of reaction, a more “active metal” displaces another element in solution. These reactions can be further classified as a solid metal reacting with a) a metal ion solution, b) an acid solution, or c) water. An example of each is provided below: a) When a solid metal reacts with a metal solution, the solid metal’s ions go into solution while the metal ions originally in solution plate out onto the surface of solid metal—e.g., Zn (s) + CuSO4 (aq)  ZnSO4 (aq) + Cu (s). b) When a solid metal reacts with an acid, the metal replaces hydrogen in the acid to produce hydrogen gas while the metal ion goes into solution with the anion from the acid—e.g., Mg (s) + 2HCl page 3 of 6 (aq)  H2 (g) + MgCl2 (aq). c) When a solid metal reacts with water, the metal replaces hydrogen in the water to produce hydrogen gas while the metal ion goes into solution with hydroxide ion—e.g., Ca (s) + 2 H2O (aq)  H2 (g) + Ca(OH)2 (aq). To predict whether or not a single-replacement reaction will occur, we refer to the Activity Series for Metals (shown on the next page). Activity Series for Metals Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Hg > Au If a solid metal is “more active” – i.e., higher on the Activity Series – than the metal ion in solution or hydrogen for acids, the more active metal will displace the less active ion, so a reaction occurs. The more active metal goes into solution while the less active ion either plates out for a metal ion or bubbles out as a gas for hydrogen ion. If the solid metal is less active than the metal ion or hydrogen in the compound, then no reaction occurs. Only six metals (Li, K, Ba, Sr, Ca, and Na) – called “active metals” – react directly with water to produce hydrogen gas and a metal hydroxide solution. These active metals are the first six metals in the Activity Series. The Chemistry behind each step of the Copper Cycle The different copper species obtained in each part is shown in Equation 1 below: Cu(s) Part I Cu2+(aq) Part II Part IV Cu(OH)2(s) Cu2+(aq) Part V Part III CuO(s) Cu(s) Step I: Chemistry I. Oxidizing Copper Metal with Concentrated Nitric Acid, HNO3(aq) The first step involves transforming Cu metal to copper(II) ions, Cu2+, using concentrated nitric acid, HNO3(aq). At the same time, the nitrate ions (NO3–) are converted to nitrogen dioxide, NO2, a brown gas. The presence of Cu2+(aq) makes the solution blue. Figure 2 (below) When the reaction mixture is diluted with water, the Cu2+ ions are hydrated (surrounded by water) to form the octahedral complex ion, [Cu(H2O)6]2+, as shown below. Six water molecules (shown as red O and white H atoms) are bonded to a Cu2+ ion (shown in gray as the central atom). Cu2+(aq) + 6 H2O(l)  [Cu(H2O)6]2+(aq) page 4 of 6 In this lab, we will refer to this complex simply as Cu2+ since complex ions are beyond the scope of CHM151. Step II: Chemistry II. Precipitating Cu(OH)2(s) with NaOH(aq) In Part II, two reactions are carried out by adding NaOH(aq). In the first reaction, the hydroxide ions (OH–) from the NaOH(aq) neutralize the excess hydronium ions (H3O+) left over from the previous part: H3O+(aq) + OH–(aq)  2 H2O(l) Once all the H3O+ ions are neutralized, additional OH– ions react with the Cu2+ ion to form Cu(OH)2 precipitate. Once all the Cu2+ ions have reacted, no more precipitate forms. Adding more OH– ions makes the solution basic, so it can turn red litmus paper blue. Figure 2 on the next page shows the step-wise reaction of Cu2+ with NaOH. Figure 3 (below) Step-wise Illustration of the Precipitation of Cu(OH)2 in Part II Cu(OH)2 Cu(OH)2 1st Beaker: At the end of Part I, Cu2+ ions are present, making the solution blue. Excess hydronium ions (H3O+) remain from the nitric acid used. 2nd Beaker: Adding NaOH(aq) to the blue solution results in the OH– ions neutralizing the H3O+ ions to form water: H3O+(aq) + OH–(aq)  2 H2O(l). The Na+ ions and resulting water molecules are not shown since they are merely spectator ions. 3rd and 4th Beakers: Once all the H3O+ are neutralized, adding more NaOH(aq) results in the OH– ions reacting with the Cu2+ to form the blue Cu(OH)2(s) precipitate shown at the bottom of the beaker. 5th Beaker: When all of the Cu2+ ions have been converted to Cu(OH)2(s) precipitate, adding more NaOH(aq) results in unreacted OH– ions in solution, which makes the solution basic. Red page 5 of 6 litmus paper can be used to confirm the solution is basic. Note that the solution is no longer blue since no Cu2+ ions are present in the solution. In reality, your solution may still look blue because of the dispersion of the Cu(OH)2 by mixing. Step III: Chemistry III.Converting solid Cu(OH)2 to solid CuO In Part III of the sequence, the reaction mixture is heated. This transforms the Cu(OH)2 precipitate to CuO precipitate. The CuO precipitate is separated from the solution, called the supernatant liquid, using a method called gravity filtration. The mixture is filtered using a filter funnel, and the solid is collected on filter paper. The supernatant liquid runs through the filter paper and collects in a beaker. This resulting filtered solution is called the filtrate. Step IV: Chemistry IV. Dissolving CuO(s) with sulfuric acid, H2SO4(aq) In Part IV, the cleaned CuO precipitate is dissolved using sulfuric acid, H2SO4(aq). This reaction returns copper to its aqueous phase as Cu2+. Step V: Chemistry V. Reducing Cu2+ ions with Zinc Metal In Part V, zinc metal (Zn) is added to the copper solution to convert the copper ions back to copper metal, Cu(s). The resulting solution will contain colorless zinc ions, Zn2+(aq) and copper solid. Visible evidence of this reaction is observed as bubbles of gas being released from the solution. (Since the H3O+ ions do not dissolve the Cu metal, the amount of copper yielded is not affected by excess acid.) Identify the gas displaced from the acid in this reaction. When the solution becomes colorless, all of the Cu2+ ions have been converted to Cu metal. All of the excess Zn metal is also converted to Zn2+ ion by the excess H3O+ ions from the sulfuric acid, H2SO4(aq),used to dissolve the CuO precipitate in Part IV. Once all the Zn metal is dissolved, the Cu metal can be isolated by decanting, or pouring off, the supernatant liquid. The Cu will then be rinsed, dried, and weighed as described in the procedure. page 6 of 6