Redox Reactions Experiment Essay

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Pages: 11

Experiment 6: Redox Reactions- The Activities Series
March 9, 2015
John Rios
Professor Judy George
Chemistry 141-6658

Objective: This experiment was designed to aid the student in understanding the activity of certain elements. In addition, this experiment also explained the reason to which certain reactions did and did not occur between the salt and metal solutions, acidic and metallic solutions, halogen and halide solutions and halide ion and metal ions. The student will be able to determine the relative strengths of certain oxidizing and reducing agents.
Introduction:
A reaction in which the oxidation state of the atoms change is formerly known as a redox reaction. When an ion in a solution is replaced by another ion then this occurrence is known as a specific type of redox reaction, being the Single displacement reaction. During oxidation, an element loses an electron which increases the oxidation state. The elements that loses the electrons claim the title: Reducing Agent or Oxidizing Element. A reduction state is when an element gains an electron to decrease the oxidation state. Oxidizing agent and reduced elements are usually the elements that are gaining electrons. Oxidation has an increased oxidation state after the chemical reaction, whereas, reduction has a decreased oxidation state after the chemical reaction.
For example: in the Lab Experiment 7: Copper Reaction, the aluminum dissolved when placed into the solution of Copper (II) Chloride, which was a blue solution. A cloudy blue color was formed and copper residue appeared at the bottom of the beaker after the aluminum dissolved into the solution.
2 Al(s) + 3 CuCl2 (aq) 2AlCl3 (aq) + 3Cu(s)
Oxidizing process: Al0 to Al +3 (AlCl3) (the oxidation state of Aluminum switched from 0 to positive 3)
Al0(s) Al+3(aq) + 3e-
Reduction process: Cu+2 (CuCl2) to Cu0 (the oxidation state of Copper switched from a positive 2 to 0)
Cu+2(aq) + 2e- → Cu (s)
Copper +2 gained the electrons, so Copper +2 is the oxidizing agent, or OA for short (i.e. reduced element).
Aluminum lost the electrons, so Aluminum is the reducing agent, or RA for short (i.e. oxidized element).
Redox half reaction: 3Cu+2(aq) + 2Al (s) 3Cu(s) + 2Al+3(aq)
Copper2+ (Strong OA)
Aluminum (strong RA)
Copper (weak RA)
Al3+ (weak OA) The copper is replaced by aluminum in the copper (II) chloride solution, therefore, we can state that aluminum is easier to oxidize (RA) than the element, copper. From this, we can deduct that aluminum is more active.
Safety: The basic safety protocols: close toed shoes, goggles, hazmat, etc...
All inorganic/organic wastes put into the waste container.
Procedure:
1. Reaction between a metal and salt solution
2. Reaction between a metal and an acid solution
3. Reaction between a halogen (I2, Br2) and halide ion
4. Reaction between a halide ion and a metal ion
Reference:
Lehman, J. & Olmstead, T. et al (2002). Experiment 6: Redox Reactions- the Activities Series. In Grossmont College Chemistry 141 Laboratory Manual (6th Edition, pp. 53-58). El Cajon, California.

Results and calculations:
Observation:
Copper Sulfate 0.1M (CuSO4(aq)): a light blue solution.
Iron (II) Sulfate 0.1M (FeSO4(aq)): a clear and colorless solution.
Zinc Sulfate 0.1M (ZnSO4(aq)): a clear and colorless solution.

Reaction between a metal and salt solution
Metal and Salt Solution Reactions
Reactant
Observation
Net ionic reaction
Element oxidized
Element reduced
Stronger oxidizing agent
More active
Cu
Zn2+
Nothing occurred
No reaction none none
Cu
Zn
Cu
Fe2+
Nothing occurred
No reaction none none
Cu
Fe
Zn
Cu2+
Zinc strip turned black. Immediate reaction occurred. The blue solution faded away.

Zn(s) + Cu2+(aq) Zn2+(aq)+ Cu(s)
Zn
Cu
Cu
Zn
Zn
Fe2+
Bubbles appeared and resided around the zinc strip.

Zn(s) + Fe2+(aq) Zn2+(aq)+ Fe(s)
Zn
Fe
Fe
Zn
Fe
Cu2+
The iron rusted and the light blue…