Category: Chemistry – AP

Chemical Equilibrium: Le Châtelier’s Principle

Chemical Equilibrium: Le Châtelier’s Principle

  1. Introduction

Equilibrium is a state where the forward rate of the reaction is equal to the reverse rate. At this state, chemical reaction is not observed making it look as if the reaction has stopped.

However, the equilibrium can be easily changed when there is any stress in the system. The equilibria’s response to these stresses is explained by Le Châtelier’s principle: if equilibria are disturbed by an outside factor, the system will react in a direction to adjust to the new environment, reestablishing the equilibrium. However, the principle does not only state that the equilibrium will shift but it also predicts how it will change. Knowing the factors that disturb the equilibrium helps figuring out whether the equilibrium will move. The factors that are specified in Le Châtelier’s principle are pressure, concentration of the reactant or product, and temperature.

The equilibrium constant (Kc) and the reaction quotient (Qc) are also important in finding which direction the equilibrium will shift to when the system is upset. The equilibrium constant is the concentration of product divided by the concentration of the reactant raised to their orders accordingly, and the constant dependant on temperature when the system is in equilibrium (Kc=K([A]a[B]b)/([C]c[D]d) when cC+dD←→aA+bB). The reaction quotient has the equal way of calculating as the equilibrium constant. However, the reaction quotient can be calculated at any point on the progression of the reaction(Qc= K([A]a[B]b)/([C]c[D]d) when cC+dDaA+bB).

The significance of the equilibrium constant and the reaction quotient is that they decide which direction the equilibrium will shift to according to their comparative values. The reaction quotient being bigger than the equilibrium constant indicates that the product’s concentration is bigger than the ones of the reactant’s shifting the equilibrium towards the left. In exact opposite, the reaction quotient being smaller than the equilibrium constant indicates that the reactant’s concentration is bigger than the ones of the product’s shifting the equilibrium towards the right.

Aim of the experiment is to define Le Châtelier’s principle according to the factors that upset the equilibrium: concentration of reactants and product, and Temperature; to observe the equilibrium using sparingly soluble salts; and to explain the direction of the shifts and their reasons.

2. Materials

    1. Apparatus
  • Chemicals in a Medicine Dropper
  • Test Tubes
  • 50mL Graduated Cylinder
  • 10mL Beaker
  • Clamp Stand, Clamp
  • Heating Plate
  • 100mL Beaker
    1. Chemical
  • 0.1M CuSO4
  • 0.1M NiCl2
  • 1.0M CoCl2
  • 0.1M KI
  • 0.1M Na2CO3
  • 0.1M AgNO3
  • 6M HNO3
  • 0.1M HCl
  • 1M HCl
  • 12M HCl
  • 15M NH3

3. Procedures

    1. Changes in Reactant or Product Concentration
      1. Copper and Nickel Ion
        1. Place 1mL of 0.1M CuSO4 in a 10mL beaker and record the color
        2. Add 15M NH3 dropwise until color change occur and the solution is clear
      • Solids may occur (Cu(OH)2)
        1. Record observations
        2. Mix the solution by tickling it while adding NH3
        3. Add 1M HCl dropwise and mix the solution until the solution is clear
        4. Record observation
        5. Repeat the procedure using 0.1M NiCl2 in place of CuSO4 and record observation
      1. Cobalt Ions
        1. Place about 0.5mL (10 drops) of 1M CoCl2 in a clean 10mL beaker
        2. Record color
        3. Drop 12M HCl to the test tube until a distinct color change occurs
        4. Record observation
        5. Slowly add water to the test tube while mixing
        6. Record observation
    1. Equilibria Involving Sparingly Soluble Salt
        1. To 0.5mL (10 drops) of 1.0M Na2CO3 in a 10mL beaker
        2. Add 1 drop of 0.1M AgNO3
        3. Record observation
        4. Add 6M HNO3 dropwise until you observe a change in appearance
        5. Record observation
        6. To the above solution, add 0.1M HCl dropwise
        7. Record observation
        8. While mixing, add 15M NH3 dropwise until chemical reaction occurs
        9. Add 6M HNO3 dropwise until chemical change is no more change.
        10. Record observation
        11. Continue to add 0.1M KI dropwise until you see evident chemical change
        12. Observe reaction
        13. Dispose the waste accordingly
    2. Effect of Temperature on Equilibria
        1. Heat 75mL of water to boil in a 100mL beaker on a heating plate
        2. Place about 1mL of 1.0M CoCl2 in a small test tube in the boiling content using a clamp stand
        3. Compare the color of the cool cobalt solution to that of the hot solution
        4. Record observation

4. Data

    1. Changes in Reactant or Product Concentration
      1. Copper and Nickel Ions
        1. CuSo4 (aq) – Light Blue
        2. [Cu(NH3)4]2+ (aq) – Dark Blue
        3. After HCl reaction – Extremely Clear, almost transparent but slightly light bluish
        4. NiCl2 (aq) – Light Green, turquoise
        5. [Ni(NH3)6]2+ (aq) – Pale violet
        6. After HCl reaction – Light Blue
      2. Cobalt Ions
        1. CoCl2 (aq) – Rose pink, slightly red
        2. After HCl reaction – Dark Blue, Gas evolve, Temperature increase
        3. After H2O reaction – Rose pink, same as the color before the HCl reaction
    2. Equilibria Involving Sparingly Solution Salts
        1. Na2CO3 (aq) – Transparent, slightly pinkish
        2. After adding AgNO3 (aq) – White precipitate formed
        3. After HNO3 reaction – White precipitate seems to dissolve, pale white
        4. After HCl reaction – Solution turned white, precipitate form
        5. After NH3 (aq) reaction – The solution turned transparent, gas evolved, extremely small amount of precipitate left
        6. After HNO3 reaction – Turned again into milky white, precipitate reappear
        7. After NH3 reaction – Turned transparent, however no gas evolved this time, precipitate disappear
        8. After adding KI – The solution turned milky green, white precipitates remained
    3. The Effect of Temperature on Equilibria
        1. Cool CoCl2 – Rose Pink
        2. Hot [CoCl4]2- – Very dark blackish rose pink

5. Analysis

    1. Changes in Reactant
      1. Copper and Nickel Ions
        1. [Cu(H2O)4]2+(aq) + 4NH3 (aq) ←→ [Cu(NH3)4]2+(aq) + 4H2O (l)
        2. [Ni(H2O)6]2+(aq) + 6NH3 (aq) ←→ [Ni(NH3)6]2+(aq) + 6H2O (l)
      • Addition of NH3 made the equilibrium shift right
          1. H+(aq) + NH3 (aq) ←→ NH4+
      • Reaction with HCl made the equilibrium shift left
      1. Cobalt Ions
        1. [Co(H2O)6]2+(aq) + 4Cl (aq) ←→ [CoCl4]2-(aq) + 6H2O (l)
      • Reaction with HCl made the equilibrium shift right
      • Reaction with H2O made the equilibrium shift left
    1. Equibria Involving Sparingly Soluble Salt
        1. 2Ag+(aq) + CO32-(aq) ←→ Ag2CO3 (s)
          1. 2H+(aq) + CO32-(aq) ←→ H2CO3 (aq)
          2. H2CO3 (aq) CO2 (aq) + H2O (l)
      • Reaction with HNO3 made the equilibrium shift left
        1. Ag+(aq) + Cl(aq) ←→ AgCl (s)
      • Reaction with HCl creates a new precipitate of AgCl and moves the equilibrium to the left
          1. Ag+(aq) + 2NH3 (aq) ←→ [Ag(NH3)2]+(aq)
      • Reaction with NH3 shifts the equilibrium to the left
          1. H+(aq) + NH3 (aq) ←→ NH4+(aq)
      • Reaction with HNO3 moves the equilibrium to the right
          1. I(aq) + Ag+(aq) ←→ AgI (s)
      • Reaction with the iodide ion creates a new precipitate of AgI and shifts the equilibrium to the left
    1. The Effect of Temperature on Equilibria
        1. CoCl2+ΔH Δ [CoCl4]2-
      • When the solution is heated, the equilibrium shifts to the right
      • When the solution is cooled, the equilibrium shifts to the left

6. Discussion

[Cu(H2O)4]2+(aq) + 4NH3 (aq) ←→ [Cu(NH3)4]2+(aq) + 4H2O (l)

[Ni(H2O)6]2+(aq) + 6NH3 (aq) ←→ [Ni(NH3)6]2+(aq) + 6H2O (l)

For above reaction NH3 reacts with both Cu2+ and Ni2+ ions to create [Cu(NH3)4]2+(aq) and [Ni(NH3)6]2+(aq) respectively. Thus, by adding NH3 to the solution, the equilibrium shifted right since while adding NH3 the concentration NH3 increases allowing more NH3 to react with the ions to create [Cu(NH3)4]2+(aq) and [Ni(NH3)6]2+(aq).

H+(aq) + NH3 (aq) ←→ NH4+

However, when HCl is added to neutralize NH3, the concentration of NH3 decreases, shifting the equilibrium left, thus creating less [Cu(NH3)4]2+(aq) and [Ni(NH3)6]2+(aq).

[Co(H2O)6]2+(aq) + 4Cl (aq) ←→ [CoCl4]2-(aq) + 6H2O (l)

Similar to above reaction, adding more HCl, the concentration of Cl(aq) increases shifting the equilibrium right. Through this reaction, more [CoCl4]2-(aq) is being created. However, when more water(H2O) is added, the concentration of H2O increases shifting the equilibrium left, creating less [CoCl4]2-(aq)

2Ag+(aq) + CO32-(aq) ←→ Ag2CO3 (s) …………….(1)

In reaction (1), CO32- reacts with Ag+ to create Ag2CO3 (s). The addition of Na2CO3 shifts the equilibrium to the right.

2H+(aq) + CO32-(aq) ←→ H2CO3 (aq)

H2CO3 (aq) CO2 (aq) + H2O (l)

In the above reaction, H+ reacts with CO32- to form H2CO3 that quickly decomposes to CO2 and water. Thus, addition of HNO3 reduced the concentration of CO32- ions shifting the equilibrium (1) left. Decrease in production of Ag2CO3 (s) is the result of this reaction. This process is observed through decrease in precipitates, Ag2CO3.

Ag+(aq) + Cl(aq) ←→ AgCl (s) …………………….(2)

From equilibrium (1), Cl from HCl reacts with Ag+ to create AgCl(aq). This decreases the concentration of Ag+(aq) shifting the equilibrium left. However, if excess HCl is added until chemical reaction has settled, the reaction has changed to equilibrium (2). This creates a new precipitate of AgCl(s).

Ag+(aq) + 2NH3 (aq) ←→ [Ag(NH3)2]+(aq)…………(3)

NH3 reacts with the silver ion to form [Ag(NH3)2]+(aq) which is a soluble substance. Thus, addition of NH3 reduces the concentration of Ag+(aq) ion shifting the equilibrium (2) left. This leads to less production of AgCl (s) so that the precipitates seem to disappear.

H+(aq) + NH3 (aq) ←→ NH4+(aq)

However, H+ react with NH3 to form NH4+(aq) which does not take part in reaction. This reduces the concentration of NH3 in equilibrium (3) which eventually increases the concentration of Ag+ in equilibrium (2). This shifts the equilibrium (2) right creating more AgCl (s) precipitate.

Process of adding NH3 and HNO3 – in this context H+ – is repeated several times. For all those times the equilibrium shifts for the same reason as explained above.

I(aq) + Ag+(aq) ←→ AgI (s)……………………….(4)

The iodide ion, I, reacts with Ag+ to form AgI(s) precipitate. This decreases the concentration of Ag+ in equilibrium (2) shifting the equilibrium left. However, if excess KI is added to the solution, the reaction completely changes to equilibrium (4).

CoCl2+ΔH Δ [CoCl4]2-

The process of CoCl2←→ [CoCl4]2- is an endothermic reaction. Thus, when heated, heat is given into the system. Thus, letting more heat get absorbed into the system shifts the equilibrium to the right.

On the other hand, if the reaction is cooled down, heat is taken out of the system. Thus, not letting heat take part in reaction shifts the equilibrium to the left.

Thus overall, all the reaction’s shift in equilibrium is explained by Le Châtelier’s principle. The change in concentration was the main reason of shifts in equilibrium. Furthermore, the change in heat demonstrated in the last reaction was also another factor that shifts equilibrium. Thus, Le Châtelier’s principle which explains the shift in equilibrium due to change in concentration and heat is well demonstrated throughout the experiment.

However, some errors that might be present on the process of the experiment are insufficient amount of substances (or even heat) that did not lead the experiment to completion and lack of observation.

One error observed was in the process of observing the color change of CoCl2+ΔH Δ [CoCl4]2- reaction. According to the text book, the color change in CoCl2+ΔH Δ [CoCl4]2- should have been from rose pink to blue. However, the color change observed was from rose pink to very dark blackish pink. One reason to explain this error is that, the amount of heat exerted in the reaction was insufficient to completely shift the equilibrium to its end but made it stop in the middle of shifting the equilibrium.

Another error that was observed was in the process of adding HNO3 to 2Ag+(aq) + CO32-(aq) ←→ Ag2CO3 (s). According to the process and its analysis, CO2 gas is supposed to bubble when HNO3 is added. However, there were no such bubbles observed. Two reasons that explain this error might be that, first; there wasn’t enough amount of solution present for any bubble to be observed and second; the amount of CO2 produced was too less that most of CO2 being produced dissolved in water before being able to form gas. The amount added to the 10mL beaker was 0.5mL of Na2CO3 which equals to about 10 drops, and a drop of AgNO3 which approximates to about 0.05mL. In total, only 0.55mL of solution was present in a 10mL beaker. This might have led to having too much surface area for bubble to be seen thus no significant bubble was observed.

Another error that was observed was that in process of adding highly concentrated acids and bases, gases evolved when it was not supposed to: when adding 12M HCl and when adding 15M NH3. One way to explain this is that HCl and NH3, being a very strong acid and a very strong base, reacts readily with water to form H2 gas for HCl and and NH4+ for NH3.

In order to improve the experiment, more time would have helped observe more definite data. Using definite data, observations of formation of more justifiable substances would have been given leading to more clear shifts of equilibrium. More time in the last experiment would have provided chances for more CoCl2 to form [CoCl4]2- thus changing the color into definite blue. To add on, addition of time would have given opportunities to carry experiments out for more than one time solidifying the argument of equilibrium shifts.

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

John J. Carroll. John D. Slupsky. Alan E. Mather. June 13, 1991. The Solubility of Carbon Dioxide in Water at Low Pressure. Ref. Data, Vol 20, No. 6: University of Alberta


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