Activity Series

Activity Series

Worked with Jake Tae, Jeremy Chang

  1. Introduction

Oxidation reduction reactions are reactions in which electrons are transferred from element to element. Furthermore, in order to express the extent that an element has been oxidized, oxidation number is used. Oxidation number, expressing the oxidation state of an atom, shows the total number of electrons which have been removed or added to the element. When an electron is removed from an element, the oxidation number increases and the increase of oxidation number means that the atom has lost electron which is called oxidation. Likewise, when an electron is added to an element and the oxidation state decreases, the process of gaining an electron is called reduction.

Reduction potential is another concept that illustrates the losing and gaining of electron. However, reduction potential not only show the element’s electron loose and gain but also shows the tendency of atoms to lose and gain electrons when it is exposed to change with an introduction of a new element. Thus, if an atom has higher reduction potential than another, the former is more likely to be reduced thus gaining electron from the other when they two react. In other words, the latter is more likely to be oxidized and donate electron to the other when reaction occurs between the two. However, the tendency of losing or gaining electron is not absolute but is dependent on one another. This is because the definition of reduction potential states that reduction potential do involve introduction of other element meaning that all measurements will also depend on the later introduced element too. Thus, experimented reduction potential of hydrogen is set to 0 so that other atoms’ reduction potential, when reacted with hydrogen, can be easily measured and denoted.

Furthermore, reduction potential as the tendency to gain electron does have a close connection to the electronegativity which is the tendency to gain electron. As it seems, since electronegativity is the tendency to gain electron and reduction potential is the exact same, they are proportional to each other which means that an electron with high electronegativity will have a high reduction potential.

However, unlike the reduction potential, activity series shows the tendency of an element to lose electrons when it is exposed to change with an introduction of a new element. Furthermore, activity series, unlike the reduction potential, does not have values but are series of metals in order of their tendency to react where the most reactive metals are on top and the least reactive at the bottom. Thus, if a metal is higher on the activity series than another, this means that the former metal wants to give away electron more readily thus be oxidized quicker.

In this experiment, six different metals will be tested with acid and their ions present on varieties of solutions to experimentally determine the tendency of oxidation and eventually create an experimented activity series of those six metals.

2. Materials

    1. Apparatus
      1. Small Test Tubes
      2. Test Tube Rack
    2. Chemicals
  1. 0.2M Ca(NO3)2
  2. 0.2M Zn(NO3)2
  3. 0.2M Mg(NO3)2
  4. 0.2M Fe(NO3)3
  5. 0.2M FeSO4
  6. 0.2M CuSO4
  7. 0.2M SnCl4
  8. 6M HCl
  9. 7 Small Pieces Each of Calcium, Magnesium, Zinc, Iron Wool, Tin, Copper

3. Procedure

    1. Reactions of Metal with Acid
      1. Add 0.5mL of dilute 6M HCl to 6 test tubes
      2. Add each piece of metal into each test tube
      3. Record any change
    2. Reactions of Metal with Solutions of Metal Ion
      1. Add 0.5mL of Ca(NO3)2 in 6 test tubes
      2. Add a small piece of calcium, copper, magnesium, iron wool, tin, and zinc
      3. Observe any color changes, evolution of gas and whether the metal dissolves or not
      4. Repeat the process with 6 other solutions and observe
    3. Reactive Activity Series
      1. Rank the metals in order of their reactivity judging from the most reactive to the least reactive

4. Data

    1. Reaction of Metals with Acid
      1. Ca
  • Observation: Extremely hot, visible gas evaporation, bubble eruption, still transparent.
      1. Cu
  • Observation: N.R.
      1. Mg
  • Observation: Hot temperature, rapid gas visibly erupts, bubble evolves, violent reaction
      1. Fe
  • Observation: Very slow reaction, bubbles however at very slow rate Gas evolves, Slowly rises up to the surface
      1. Sn
  • Observation: N.R.
      1. Zn
  • Observation: Bubble eruption. Is rapid reaction but not as much as calcium
    1. Reaction of Metals with Solution of Metal Ion
      1. Ca
  • Ca2+ N.R.
  • Cu2+ White Gas evolve, bubble eruption, becomes blueish green in color, blue precipitate
  • Fe3+ Bubble eruption, White precipitate, turquoise
  • Fe2+ Bubble erupts, visible white gas, color of rust, temperature increase
  • Mg2+ Bubble erupts, Gas evolution, violent reaction
  • Sn4+ Bubble, White gas
  • Zn2+ Hot, white bubble form, gas form, White precipitate
      1. Cu
  • Ca2+ N.R.
  • Cu2+ N.R.
  • Fe3+ N.R.
  • Fe2+ N.R.
  • Mg2+ N.R.
  • Sn4+ N.R.
  • Zn2+ N.R.
      1. Fe
  • Ca2+ N.R.
  • Cu2+ Solution turned murky yellow, the iron turned copper red, very slow reaction
  • Fe3+N.R.
  • Fe2+ N.R.
  • Mg2+ N.R.
  • Sn4+ N.R.
  • Zn2+ N.R.
      1. Mg
  • Ca2+ N.R.
  • Cu2+ Bubble erupts, liquid turns greenish yellow
  • Fe3+ slow bubble eruption, turned murky orange with a taste of white
  • Fe2+ Bubble erupts extremely slowly
  • Mg2+ N.R.
  • Sn4+ White Gas evolves, Solution turned yellowish translucent
  • Zn2+ solution turned murky
      1. Sn
  • Ca2+ N.R.
  • Cu2+ N.R.
  • Fe3+ N.R.
  • Fe2+ N.R.
  • Mg2+ N.R.
  • Sn4+ N.R.
  • Zn2+ N.R.
      1. Zn
  • Ca2+ N.R.
  • Cu2+ Liquid turned transparent, zinc turned black, little bubble eruption
  • Fe3+ Solution turned red
  • Fe2+ N.R.
  • Mg2+ N.R.
  • Sn4+ Bubble erupts
  • Zn2+ N.R.

5. Analysis

    1. Reactions of Metal with Acid
      1. Ca(s) + 2H+(aq) Ca2+(aq) + H2(g)
      2. Cu(s) + 2HCl(aq) Cu(s) + 2HCl(aq) : N.R.
      3. Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)
      4. Fe(s) + 2H+(aq) Fe2+(aq) + H2(g)
      5. Sn(s) + 4HCl(aq) Sn(s) + 4HCl(aq) : N.R.
      6. Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)
    2. Reactions of Metal with Solutions of Metal Ion
  • The reactions with metal and the same ion are not written since they do not and should not react with each other when put together.
      1. Ca(NO3)2(aq)
        1. Cu(s) + Ca+2(aq) Cu(s) + Ca+2(aq) : N.R.
        2. Mg(s) + Ca+2(aq) Mg(s) + Ca+2(aq) : N.R.
        3. Fe(s) + Ca+2(aq) Fe(s) + Ca+2(aq) : N.R.
        4. Sn(s) + Ca+2(aq) Sn(s) + Ca+2(aq) : N.R.
        5. Zn(s) + Ca+2(aq) Zn(s) + Ca+2(aq) : N.R.
      2. CuSO4(aq)
        1. Ca(s) + Cu2+(aq) Cu(s) + Ca2+(aq)
        2. Mg(s) + Cu2+(aq) Cu(s) + Mg2+(aq)
        3. Fe(s) + Cu2+(aq) Cu(s) + Fe2+(aq)
        4. Sn(s) + Cu2+(aq) Sn(s) + Cu2+(aq) : N.R.
        5. Zn(s) + Cu2+(aq) Cu(s) + Ca2+(aq)
      3. FeSO4(aq)
        1. Ca(s) + Fe2+(aq) Fe(s) + Ca2+(aq)
        2. Cu(s) + Fe2+(aq) Cu(s) + Fe2+(aq) : N.R.
        3. Mg(s) + Fe2+ (aq) Fe(s)+ Mg2+(aq)
        4. Sn(s) + Fe2+(aq) Sn(s) + Fe2+(aq) : N.R.
        5. Zn(s) + Fe2+(aq) Fe(s)+ Zn2+(aq)
      4. Fe(NO3)3(aq)
        1. 3Ca(s) + 2Fe3+(aq) 2Fe(s) + 3Ca2+ (aq)
        2. Cu(s) + Fe3+ (aq) Cu(s) + Fe3+ (aq) : N.R.
        3. 3Mg(s) + 2Fe3+ (aq) 2Fe(s) + 3Mg2+ (aq)
        4. Sn(s) + Fe3+ (aq) Sn(s) + Fe3+ (aq) : N.R.
        5. Zn(s) + Fe3+ (aq) Zn(s) + Fe3+ (aq) : N.R.
      5. Mg(NO3)2(aq)
        1. Ca(s) + Mg2+ (aq) Mg(s) + Ca2+ (aq)
        2. Cu(s) + Mg2+ (aq) Cu(s) + Mg2+ (aq) : N.R.
        3. Fe(s) + Mg2+ (aq) Fe(s) + Mg2+ (aq) : N.R.
        4. Sn(s) + Mg2+ (aq) Sn(s) + Mg2+ (aq) : N.R.
        5. Zn(s) + Mg2+ (aq) Zn(s) + Mg2+ (aq) : N.R.
      6. SnCl4(aq)
        1. 2Ca(s) + Sn4+(aq) Sn(s) + 2Ca2+(aq)
        2. Cu(s) + Sn4+(aq) Cu(s) + Sn4+(aq) : N.R.
        3. 2Mg(s) + Sn4+(aq) Sn(s) + 2Mg2+(aq)
        4. 2Fe(s) + Sn4+(aq) Sn(s) + 2Fe2+(aq)
        5. 2Zn(s) + Sn4+(aq) Sn(s) + 2Zn2+(aq)
      7. Zn(NO3)2(aq)
        1. Ca(s) + Zn2+(aq) Zn(s) + Ca2+(aq)
        2. Cu(s) + Zn2+(aq) Cu(s) + Zn2+ (aq) : N.R.
        3. Mg(s) + Zn2+(aq) Zn(s) + Mg2+ (aq)
        4. Fe(s) + Zn2+(aq) Fe(s) + Zn2+ (aq) : N.R.
        5. Sn(s) + Zn2+(aq) Sn(s) + Zn2+ (aq) : N.R.
    1. Experimental Activity Series
  1. Ca Ca2+ + 2e
  2. Mg Mg2+ + 2e
  3. Fe Fe3+ + 3e
  4. Zn Zn2+ + 2e
  5. Fe Fe2+ + 2e
  6. Sn Sn4+ + 4e
  7. Cu Cu2+ + 2e

Calcium seemed to be the most reactive amongst all the metal judging from the fact that it showed the most violent reaction with HCl and its metal reacted with every single solution present which means that calcium likes to oxidize more readily compared to other metals

Magnesium seemed to be the next reactive metal since its reaction with HCl was as violent as with calcium however did not readily oxidize compared to the calcium ion.

Zinc was predicted to be the next reactive amongst the metals since it did react with both HCl and many other solutions, however it neither reacted as violently as the previous two did nor reacted with solutions containing calcium and magnesium which means that it does not want to oxidize compared to them.

Iron generally seemed like the next reactive metal judging from the fact that it slowly reacted with HCl and with solutions containing copper and tin. In this experiment, two different ions of iron took part which were Fe(II) and Fe(III), ferrous and ferric. However, they both showed different traits of observation. Fe2+ in solution did react with zinc which means that zinc oxidizes more readily compared to Fe2+. Thus Fe2+ is lower on the activity series compared to zinc. On the other hand, Fe3+ did not react when zinc metal was added which means that Fe3+ oxidizes more readily than zinc. This locates Fe3+ higher on the activity series compared to zinc. However since Fe2+ did not react with tin and copper, Fe2+ stays between zinc and tin, when since Fe3+ did react with both magnesium and calcium, it stays between magnesium and zinc in the activity series.

Both tin and copper did not show much trait of reaction in both HCl reaction and reaction with varieties of ions. They neither reacted with HCl nor did with each other’s ions. Thus distinguishing their levels on the activity series is very difficult. However, the fact that iron only reacted with copper and not tin shows that tin oxidizes more readily compared to copper. Thus, tin has been put after Fe2+ and copper following tin on the activity series.

6. Discussion

The experimental activity series show similar traits as the actual activity series. However the major difference between them might be that known activity series does not distinguish the ions of irons but just categorizes them into iron so that further details cannot be seen. Furthermore, the known activity series unlike the experimental series has leveled out many more metals along the series that with experiment, would have been too vague to be differentiated.

Connecting the activity series to the periodic table, as discussed before in the introduction section, electronegativity is closely connected to the reduction potential thus the activity series. Electronegativity which shows the tendency of an atom to gain an electron is indirectly proportional to the activity series which shows the tendency of an atom to lose an electron since the elements as they go up the series prefers to oxidize more readily thus lose an electron. This allows the going up of the activity series the exactly opposite periodic trend as the electronegativity. Thus, the periodic trend of the electronegativity, is inversed in order to illustrate the periodic trend of metals going up the activity series. Furthermore this is proven through the experimental data and the expanded activity series. They show that calcium, a metal right below magnesium in the periodic table is more reactive compared to it and iron on the left of copper on the same period is more reactive than copper. This shows that the general trend of going up the activity series is from right to the left and from top to bottom of the periodic table.

However, there are exceptions on this trend as electronegativity do too, such as the filled and half-filled orbital exceptions do apply to this rule as well. To elaborate on the exceptions, the elements whose valence electrons either completely fill an orbital such as zinc whose d orbitals are filled tend to show a slight dent in the trend. This is due to their decreased electronegativity at the state of filled and half-filled orbitals. Thus, zinc, according to the general rule of periodic trend should have been under copper, however, because of the exceptions of the trend, saying that zinc has lower electronegativity than copper, it can be easily concluded that zinc is more reactive than copper.

Even though the experiment did generally show successful results, some of the observations were not as clear as it should have been. This mistake might have been committed due to lack of time. The experiment was performed under great time constraints since about 50 different reactions should have been carried out and all their observations should have been recorded at about an hour and a half. Furthermore, the acknowledgement given by the lab instructor at the beginning of the lab also created anxiety amongst the observers. Additionally, most of the reactions took very long time to even show some traits of reaction. For example, the reaction between iron and HCl might have been ignored if the test tubes were disposed before the first bubble erupt. This shows that there were ample number of chances that many of the observations made above couldn’t have been observed and vice versa, that there are chances that the observers might not have observed long enough for the reaction to be visible thus missing on some reactions.

Furthermore, during the experiment there were confusions regarding whether some particles visible were precipitates or residues of leftover metals. Due to lack of conciseness, the experiment was not carried out in controlled environment but was carried out by adding large pieces of metals to a small amount of solutions. This led to the ions present in the solution being the limiting reagent thus allowing the metals to leftover. This created small particles of unused metals. On the other hand, many of the above reaction created precipitates. However, since many precipitates are of similar color to the original metal, it was extremely hard for the observers to distinguish whether a particle is a residue or a precipitate thus creating mistakes on observation.

In order to improve on the mistakes made above, more time allotted would have helped the experiment. Increase of time would have helped the observers carry out the experiment in a more relaxed environment allowing better observations. Furthermore, time increase would have allowed the reactions to go to their completion allowing more accurate observations and results. Additionally, accurate measurement and calculation of the reactants would have helped to see a complete reaction. And finally, pre calculations of the reaction would have helped to determine whether a precipitate might form in a certain reaction or not allowing more accurate observations.

Work Cited

Brown and LeMay the Central Science Lab Manual

Brown, Lemay, Bursten, Murphy, Woodward, Stoltzfus. 2015. Chemistry the Central Science. 13th edition. United States of America: Pearson

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