Activity Series Single Replacement Reactions

Activity Series: Single Replacement Reactions -Molecular and Ionic equations -Displacement reactions -Balancing redox equations Learning Goals: • To observe reactions of metals with acids and salt solutions using activity series of metals. • To be able to write and balance single replacement reactions using activity series of metals. • To be able to write ionic and net ionic equations for the above reactions. Background information and theory: One of the most important characteristics of a metal is its activity (reactivity). The activity of a metal is its ability to react with nonmetals. In such reaction, a metal typically loses electrons to the nonmetal to form cations and the product is an ionic compound. Different metals have different activity. Some metals are very active such as lithium, sodium, potassium, and cesium; some are slowly active such as tin, lead, and iron; and some are not active at all (inactive) such as copper, silver, and mercury. Inactive metals are those which resist oxidation. These are often called the noble metals. An activity series is the ranking of metals according to their reactivity. In an activity series metals are arranged in order of decreasing ability to lose electrons. Generally, the more reactive metals are placed at the left of the list. Li >K >Ba >Ca >Na >Mg >Al >Mn >Zn >Cr >Fe >Co >Ni >Sn >Pb > H >Cu >Ag >Hg >Pt >Au More active less active Any metal before hydrogen will displace hydrogen gas, H2(g), from an acid, and any metal after hydrogen will not displace hydrogen from an acid. For example: Active metals (left to H) ; Inactive metals (right to H) ; Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g) Cu (s) + HCl (aq) → no reaction (NR) In performing this experiment, you will observe and compare the reaction of some common metals with acids and salt solutions and establish an order of reactivity, from most active to least active. PART I: Reactivity of metals with HCl Very active metals will react with strong acids such as hydrochloric acid. An example considers magnesium: Molecular equation: Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g) Ionic equation: Mg (s) + 2H +(aq) + 2 Cl- (aq) → Mg 2+ (aq) + 2 Cl- (aq) + H2 (g) Net-ionic equation Mg (s) + 2H +(aq) → Mg 2+ (aq) + H2 (g) Note: Very active metals react vigorously and even explosively in acid. Thus Li, K, and Na should not be tested with hydrochloric acid. Somewhat less reactive metals will slowly react with a strong acid such as hydrochloric acid. An example considers aluminum: Molecular equation: 2Al (s) + 6HCl (aq) → 2AlCl3 (aq) + 3H2 (g) Ionic equation: 2Al (s) + 6 H+ (aq) + 6 Cl -(aq) → 2 Al 3+(aq) + 3 Cl -(aq) + 3H2 (g) Net-ionic equation: 2Al (s) + 6 H+ (aq) → 2 Al 3+(aq) + 3H2 (g) PART II: Reactivity of metal with aqueous salt solution The reaction of metals with a solution of salts of another metal is also revealing. The relative activities of the less active metals can be determined by observing reactions in which a more reactive metal (higher activity), such as magnesium, Mg, displaces ions of a less reactive metal, such as iron sulfate, FeSO4, as a typical metallic salt or aqueous solutions. Molecular equation: Ionic equation: Net-ionic equation: → Mg (s) + FeSO4 (aq) MgSO4 (aq) + Fe (s) Mg (s) + Fe2+(aq) + SO42-(aq) → Mg2+(aq) + SO42- (aq) + Fe (s) Mg (s) + Fe2+(aq) → Mg2+(aq) + Fe (s) In this example, magnesium is said to be more active than iron because in the activity series magnesium is to the left of iron. The reverse of the reaction will not occur because Fe is less reactive and cannot displace ions of a more reactive metal, such as magnesium sulfate, MgSO4, as aqueous solution. Fe (s) + MgSO4 (aq) → N.R. or Fe (s) + Mg2+ (aq) → N.R. Because a solution of many metal ions is colorless, often we must examine the reaction of a mixture carefully to establish whether or not a reaction has occurred. Changes in the appearance of the metal, a solution-color change, or the formation of precipitate all indicate that a reaction of this type has occurred. Sometimes, because metallic salt solutions are acidic, hydrogen gas, H2, is produced when the metal is added. In such cases, the formation of hydrogen gas, does not indicate a reaction between the metal and the salt solution. The fact that a metal does or does not react with a solution of another metallic salt can be used to establish the comparative reactivity of several metals. In this experiment, you will examine relative activities of common metals and compare your initial and final observations of the metals and the solutions. Based on these observations, you will determine whether or not a reaction has occurred, and, if so, with what intensity, slowly or rapidly. Using these results, you will arrange the metals in decreasing or increasing order of reactivity. Activity Series Data Data Table 1: Reactivity of metals with Hydrochloric Acid, HCl Metals Would Observation it Reactions (molecular, ionic, net-ionic) with HCl work? Molecular equation: Ca(s) + 2HCl(aq) → H2(g) + CaCl2(aq) yes Very carbonated, + 2+ lots of Ionic equation: Ca(s) + 2H (aq) + 2Cl (aq) → H2(g) + Ca (aq) + 2Cl (aq) Ca continuous Net-ionic equation: Ca(s) + 2H+(aq) + → H2(g) + Ca2+(aq) bubbles no Molecular equation: Mg(s)+2HCl(aq) → MgCl (aq) + H2(g) Very carbonated, Ionic equation: Mg+2H++2Cl- → Mg2++2Cl-+H2 lots of Net ionic equation: Mg(s)+2H+(aq) → Mg2+(aq)+H2(g) steady bubbles no No bubbles Molecular equation: Fe(s)+2HCl(aq) → FeCl2(aq)+H2(g) at all Ionic equation: Fe+2H++2Cl- → Fe2++Cl-+H2 Net ionic equation: Fe(s)+2H+(aq) → Fe2+(aq)+H2(g) no No bubbles yes Lots bubbles yes No bubbles Mg Fe NO REACTION Cu Zn Sn of Molecular equation: Zn(s)+2HCl(aq) → ZnCl+H2 Ionic equation: Zn+2H++2Cl- → Zn2++Cl-+H2 Net ionic equation: Zn(s)+2H+(aq) →Zn2+(aq)+H2(g) Molecular equation: Sn(s)+2HCl(aq) → SnCl+H2 Ionic equation: Sn+2H++2Cl- → Sn2++2Cl-+H2 Net ionic equation: Sn(s)+2H+(aq) → Sn2++H2(g) Data Table 2: Observations of Reactivity of metals with metallic salt solutions Indicate your prediction by placing a checkmark on the corner of each box for those reactions that you think will work. Write your observations inside each box. Boxes in gray should remain empty as these reactions are not presented. Metals Ca2+ as Mg2+ as Fe3+ as Cu2+ as Zn2+ as Sn4+ Ca(C2H3O2)2 MgSO4 Fe(NO3)3 CuSO4 Zn(NO3)2 SnCl4 ✓ ✓ ✓ ✓ ✓ Ca ✓ Mg ✓ ✓ Fe ✓ ✓ Cu ✓ Zn ✓ Sn as Data Table 3: Reactions and activity series Complete equation Example: Ca(s) + 2HCl(aq) → H2(g) + CaCl2(aq) Net ionic equation Ca(s)+CuSO4(aq) → CaSO4+Cu(s) Ca(s) + Cu2+(aq) → Cu2(g) + Ca2+ Mg(s)+CuSO4(ag) → MgSO4+Cu Mg(s)+Cu2+(aq) → Mg2++Cu Zn(s)+CuSO4(aq) → ZnSO4+Cu Zn(s)+Cu2+(aq) → Zn2++Cu 3Ca(s)+2Fe(NO3)3(aq) → 3Ca(NO3)2(aq) +2Fe 3Ca(s)+2Fe3+(aq) → 3Ca2+(aq)+2Fe(s) 3Mg(s)+2Fe(NO3)3(aq) → 2Fe(s)+3Mg(NO3)2(aq) 3Mg(s)+2Fe3+(aq) →2Fe(s)+3Mg2+(aq) Ca(s)+MgSO4(aq) → CaSO4(aq)+Mg(s) Ca2+(s)+Mg-(aq) → Ca(s)+Mg+(aq) 2Ca(s) + SnCl4(aq) → 2CaCl2(aq) + Sn(s) 2Ca(s)+Sn(aq) → Sn2+(s)+2Ca(aq) 2Mg(s) + SnCl4(aq) → 2MgCl2(aq)+ Sn(s) 2Mg(s)+Sn(aq) → Sn2+(s)+2Mg(aq) 2Zn(s) + SnCl4(aq) → 2ZnCl2(aq) + Sn(s) 2Zn(s)+Sn(aq) → Sn2+(s) +2Mg(aq) 2Ca + 2Zn(NO3)2 →2Ca(NO3)2 + Zn2 2Ca(s)+Zn3+(aq) → 2Ca3+(aq)+Zn(s) Most reactive 1. CuSO4 2. SnCl4 3. Fe(NO3)3 Ca(s) + 2H+(aq) + → H2(g) + Ca2+(aq) 4. MgSO4 Least reactive 5. Zn(NO3)2 6. Ca(C2H3O2)2 Activity Series Postlab questions: Your instructor may ask you to answer these in your lab notebook, or to answer directly on this page and turn it in, or to include these answers in a formal lab report. Follow your instructor’s directions. 1. Balance the following reactions and indicate if they will occur (use activity series). a) Co+2HCl → CoCl2+H2: reaction will occur b) Hg + 2HCl → HgCl2 + H2: a reaction will occur c) 3Zn + 2Ni(NO3)3 → 3Zn(NO3)2 + 2Ni: a reaction will occur d) 2Li + 2HCl → 2LiCl + H2: a reaction will occur 2. What observations would you look for to determine if a metal undergoes reaction? You would look for the formations of salts/crystals or precipitates along with types of gaseous reactions like bubbles. 3. For each of the following reactions, identify; a. oxidized species b. reduced species c. oxidizing agent I) 2 Fe (s) + • • • → 2 FeCl3(s) Oxidized/reduced species: 2Fe+3Cl Oxidizing agent: 3Cl2 Reducing agent: 2Fe II) Mg (s) • • • 3 Cl2 (g) d. reducing agent + 2 H2SO4 (aq) → MgSO4 (aq) + Oxidized/reduced species: Mg+2H2SO4 Oxidizing agent: Mg Reducing agent: 2H2SO4 SO2 (g) + 2 H2O (l)