Chemistry Case
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Explain what is meant by “electrolysis” and describe how it takes place in an electrolytic cell.
Answer: If we look at the latin roots of the word “electrolysis” we learn that it means, essentially, to “break apart” (lysis) using electricity.
In chemistry and manufacturing, electrolysis is a method of using an electric current to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially highly important as a stage in the separation of elements from naturally-occurring sources such as ores using an electrolytic cell.
In the case of an electrolytic cell, electrical energy is needed to drive non-spontaneous redox reactions. It is, as if, requiring energy to pump water uphill (since, water naturally flows downward).
How does the electrolysis of molten NaCl differ from the electrolysis of aqueous NaCl?
Answer: The electrolysis of an aqueous sodium chloride produces oxygen at the anode and hydrogen at the cathode. Meanwhile, the electrolysis of a molten sodium chloride produces sodium metal at cathode and chlorine gas at anode.
Describe the electrolysis of water. Why is sulfuric acid added to the water?
Answer: During the early history of the earth, hydrogen and oxygen gasses spontaneously reacted to form the water in the oceans, lakes, and rivers we have today. That spontaneous direction of reaction can be used to create water and electricity in a galvanic cell (as it does on the space shuttle). However, by using an electrolytic cell composed of water, two electrodes and an external source emf one can reverse the direction of the process and create hydrogen and oxygen from water and electricity. shows a setup for the electrolysis of water.
To illustrate, two cylinders containing sulfuric acid dissolved in water are connected with tubing. When electricity passes through the system, oxidation occurs at one electrode and reduction at the other. Electrolysis of the water generates hydrogen gas at one electrode and oxygen at the other. The role of the sulfuric acid is to allow transfer of charge from one platinum electrode to the other. As the time lapse shows, the volume of gas collected at each electrode differs. The volume of gas generated at each electrode can be determined from the volume readings on the cylinders. When gas from the left cylinder is collected in a test tube and a lighted match is brought near it, the gas ignites with a barking noise, indicating hydrogen. When a glowing match is brought near gas from the right cylinder, the match burns brightly, indicating the presence of oxygen.
Definition of terms:
An electrochemical cell is a device capable of either deriving electrical energy from chemical reactions, or facilitating chemical reactions through the introduction of electrical energy. A common example of an electrochemical cell is a standard 1.5-volt “battery”. (Actually a single “Galvanic cell”; a battery properly consists of multiple cells.)
An electrode is an electrical conductor used to make contact with a nonmetallic part of a circuit (e.g. a semiconductor, an electrolyte or a vacuum). The word was coined by the scientist Michael Faraday from the Greek words elektron (meaning amber, from which the word electricity is derived) and hodos, a way.[1]
An electrolytic cell decomposes chemical compounds by means of electrical energy, in a process called electrolysis; the Greek word lysis means to break up. The result is that the chemical energy is increased. Important examples of electrolysis are the decomposition of water into hydrogen and hydroxide, and bauxite into aluminium and other chemicals.
Voltaic cell A voltaic cell is created whenever dissimilar metals, connected in some way, are immersed in a conductive fluid.
electric cell
A device, such as a battery, that is capable of changing some form of energy, such as chemical energy or radiant energy, into electricity. Also called voltaic cell. An electric cell that converts light energy into electrical energy using the photoelectric effect is called a photoelectric or photovoltaic cell; such cells are used in the generation of solar power and are called solar cells. See also galvanic.
Electroplating is a plating process that uses electrical current to reduce cations of a desired material from a solution and coat a conductive object with a thin layer of the material, such as a metal. Electroplating is primarily used for depositing a layer of material to bestow a desired property (e.g., abrasion and wear resistance, corrosion protection, lubricity, aesthetic qualities, etc.) to a surface that otherwise lacks that property. Another application uses electroplating to build up thickness on undersized parts.
The process used in electroplating is called electrodeposition. It is analogous to a galvanic cell acting in reverse. The part to be plated is the cathode of the circuit. In one technique, the anode is made of the metal to be plated on the part. Both components are immersed in a solution called an electrolyte containing one or more dissolved metal salts as well as other ions that permit the flow of electricity. A rectifier supplies a direct current to the anode, oxidizing the metal atoms that comprise it and allowing them to dissolve in the solution. At the cathode, the dissolved metal ions in the electrolyte solution are reduced at the interface between the solution and the cathode, such that they “plate out” onto the cathode. The rate at which the anode is dissolved is equal to the rate at which the cathode is plated, vis-a-vis the current flowing through the circuit. In this manner, the ions in the electrolyte bath are continuously replenished by the anode.[1]
Other electroplating processes may use a nonconsumable anode such as lead. In these techniques, ions of the metal to be plated must be periodically replenished in the bath as they are drawn out of the solution.[2]
A half cell is a structure that contains a conductive electrode and a surrounding conductive electrolyte separated by a naturally-occurring Helmholtz double layer. Chemical reactions within this layer momentarily pump electric charges between the electrode and the electrolyte, resulting in a potential difference between the electrode and the electrolyte. The typical reaction involves a metal atom in the electrode being dissolved and transported as a positive ion across the double layer, causing the electrolyte to acquire a net positive