The objectives of this laboratory are:
b) To become familiar with some of the observable signs of these reactions,
c) To predict and identify the products formed in each of these reactions,
d) To write balanced chemical equations for each single replacement reaction.
e) To make three voltaic cells by using a salt bridge, a citrus fruit and potato.
f) To understand the function of a salt bridge.
During a chemical reaction both the form and composition of matter are changed. Old substances are converted to new substances, which have unique physical and chemical properties of their own. Some of the observable signs that a chemical reaction has occurred include the following:
· A metallic deposit appears
· Bubbles appear
· A temperature change occurs
· A color change occurs
· A precipitate (cloudy, tiny particles) appears
Note that there are many other observable signs for chemical reactions, but these are some of the ones most likely to be encountered in this lab.
In Part A of this lab we will examine Single Replacement Reactions. This is one type of oxidation reduction reaction, or redox (pronounced ree-dox) reaction, because it occurs via a transfer of electrons. All single replacement reactions have the general form: A + BC ® B + AC
Here, A is an element and BC is usually an aqueous ionic compound or an acid (consisting of B+ and C- aqueous ions). Element A replaces B in BC, resulting in the formation of a new element B and a new ionic compound or acid, AC. If the new element B is a metal, it will appear as a metallic deposit. If it is a gas, it will appear as bubbles.
An Activity Series of elements is often used to determine if A will displace B in a single replacement reaction. An Activity Series is provided below. As a rule, if A has a higher activity that B, a single replacement reaction will occur. However, if A has lower activity than B, a single replacement reaction will not occur.
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Example 1: magnesium metal + aqueous aluminum chloride
Since Mg is more active than Al, a single displacement reaction will occur. The predicted products are aluminum metal and aqueous magnesium chloride
Reaction Equation: 3 Mg (s) + 2 AlCl3 (aq) ® 2 Al (s) + 3 MgCl2 (aq) A B C B A C |
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Example 2: aluminum metal + aqueous magnesium chloride Since Al is not more active than Mg, a single displace reaction will NOT occur.
Reaction Equation: Al (s) + MgCl2 (aq) ® NO REACTION A B C |
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ACTIVITY SERIES FOR METALS (and HYDROGEN)
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Oxidation numbers tell how many electrons each atom in a compound possesses relative to the free atom. It serves as a “book-keeping” system so that the flow of electrons can be observed. If the oxidation number of any atom changes during a chemical reaction, a transfer of electrons has occurred. When an atom loses electrons, it is being oxidized (or has undergone oxidation). Oxidation is a loss of electrons. Conversely, when an atom gains electrons, it is been reduced (or has undergone reduction). Reduction is a gain of electrons. When describing redox reactions relative to the entire molecule, we use the terms oxidizing agent and reducing agent. A reducing agent is a substance that is being oxidized and thus causes another substance to be reduced. An oxidizing agent is a substance that is being reduced and thus causes another substance to be oxidized. We can remember these concepts with the helpful mnemonic “OIL RIG”:
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O I L |
xidation s The REDUCING AGENT oss of Electrons |
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R I G |
eduction s The OXIDIZING AGENT ain of Electrons |
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Example 1: magnesium metal + aqueous aluminum chloride Revisited )) Oxidation Number: 0 +3 -1 0 +2 -1 Reaction Equation: 3 Mg (s) + 2 AlCl3 (aq) ® 2 Al (s) + 3 MgCl2 (aq) A B C B A C |
In the above example, magnesium is oxidized (the reducing agent) because it has lost electrons. In addition, the aluminum atom in aluminum chloride is reduced because it has gained an electron. Thus, aluminum chloride is the oxidizing agent.
Electricity can be described as a flow of electrons through a wire. This form of energy is caused by the motion of electrons. A device that creates electrical current from redox reactions is called an electrochemical cell, voltaic cell, galvanic cell or battery. Batteries serve as a source of energy for flashlights, radios, as well as car motors. In Part B we will build several such batteries. It will consist of two metals (called electrodes) connected by a salt bridge between individual half-cells. A salt bridge, which contains a strong electrolyte, will join the two electrochemical reactions and complete the electrochemical circuit. It allows for the overall mixing of the two solutions. The metal strip where oxidation occurs is call the anode and is labeled with a negative (-) sign. The metal strip where reduction occurs is called the cathode and is labeled with a (+) sign. Such designations can be observed on the batteries that we us everyday. Electrons flow from anode to the cathode. This can be remembered by using the visual mnemonic:
A (anode) -------------> O (oxidation)
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| e flow
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ν
C (cathode) R (reduction)
(Developed by Paul Dement Spring 07)
Procedure
Safety
Be especially cautious when using the 6M HCl, and 3M H2SO4. These substances can burn your skin. Also be aware that skin discoloration will result from contact with AgNO3. If you feel any tingling sensations or see any color changes on your skin, flush with water immediately for a minimum of 15 minutes. Inform your instructor of any chemical contact as soon as possible.
Materials and Equipment
Solids: Copper metal, zinc metal, magnesium metal, solid sodium bicarbonate, copper metal plate, clean and shiny pennies, nickels and dimes
Solutions: 3M sulfuric acid, 6M hydrochloric acid, 1M sodium chloride, all other solutions are 0.1M and include silver nitrate, lead(II) nitrate, copper(II) sulfate, zinc nitrate, nickel(II) nitrate, aluminum sulfate
Equipment: 5 medium test tubes, plastic test tube rack, voltmeters, paper towels,. aluminum foil
1. Use the medium sized test tubes. Always use clean test tubes that have been rinsed with distilled water. The test tubes do not have to be dry.
2. Use approximately 3-mL quantities of all solutions. A good estimate is to use two full dropper squirts of each chemical.
3. Place one piece of metal in the test tube first, and then add the solution. The metal should be completely immersed in the solution used. If results are not obtained immediately, give the reaction some time. Some reactions take longer than others.
4. Perform the following reactions, and record your observations for each on the data sheet. All waste is to be disposed of in the plastic container in the hood!
a. Zinc metal + hydrochloric acid
b. Copper metal + aqueous silver nitrate
c. Copper metal + aqueous zinc nitrate
d. Zinc metal + aqueous lead(II) nitrate
e. Magnesium metal + sulfuric acid
Part B: Batteries
Salt Bridge Battery
1. Into a 250-mL beaker add approximately 25 mL of 1M copper(II) sulfate. (This is the cathode.) Label as Beaker #1.
2. Into another 250-mL beaker, add approximately 25 mL of 1M zinc sulfate. (This is the anode.) Label as Beaker #2.
3. Connect the solutions in the beakers by placing one end of a 10- inch piece of cotton twine into Beaker #1 and the other end into the Beaker #2. Obtain twine which has been soaked in a concentrated potassium sulfate solution from your instructor. (The twine is the salt bridge.)
4. Place a polished piece of copper metal plate into Beaker #1. Connect to voltmeter using wire clips.
5. Place a polished piece of zinc metal plate into Beaker #2. Connect to voltmeter using wire clips.
6. Turn voltmeter to 2V and record voltage.
7. Remove salt bridge. Record voltage.
Citrus Cell Battery
1. Cut a lemon or grapefruit in half across the segments.
2. Place the polished piece of zinc metal plate into one half of the fruit. Connect to voltmeter using wire clips.
3. Place a polished piece of copper metal plate into the other half of the fruit. Connect to voltmeter using wire clips.
4. Turn voltmeter to 2V and record voltage.
5. Repeat 1- 4 with a potato. Record observations.
Coin Battery