GOLDEN STARS MONTESSORI ACADEMYArial;
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CHEMISTRY: SS 2
TOPIC: Electrochemical Cells
Electrochemical Cells: Galvanic and Electrolytic Cells, Construction, Working Principles, Applications
CORE CONCEPTS
Electrochemical cells are devices that either convert chemical energy into electrical energy (Galvanic cells) or use electrical energy to drive non-spontaneous chemical reactions (Electrolytic cells). They are fundamental to many technologies, from batteries to industrial chemical production.
KEY POINTS
These are systems where chemical reactions produce electricity or where electricity causes chemical reactions. They consist of two electrodes (anode and cathode) and an electrolyte. Oxidation always occurs at the anode, and reduction always occurs at the cathode.
Galvanic cells convert chemical energy from spontaneous redox reactions into electrical energy. They are the basis of batteries.
Construction:
Two separate half-cells, each containing an electrode immersed in an electrolyte solution.
Anode: The electrode where oxidation occurs (negative terminal).
Cathode: The electrode where reduction occurs (positive terminal).
Salt Bridge: Connects the two half-cells, allowing ion flow to maintain electrical neutrality.
External Circuit: A wire connecting the electrodes, allowing electron flow.
Working Principles (e.g., Daniell Cell - Zinc-Copper Cell):
At the Anode (Zinc electrode): Zinc metal oxidizes, releasing electrons and forming Zn2+ ions.
Zn(s) Zn2+(aq) + 2e-
Electrons flow through the external circuit from the anode to the cathode.
At the Cathode (Copper electrode): Copper(II) ions from the solution gain electrons and deposit as copper metal.
Cu2+(aq) + 2e- Cu(s)
The salt bridge allows ions (e.g., K+ to cathode, Cl- to anode) to move between half-cells, preventing charge build-up.
Overall reaction: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Applications:
Dry cells (torch batteries), car batteries, fuel cells, portable electronic devices.
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Electrolytic cells use electrical energy from an external source to drive non-spontaneous chemical reactions. This process is called electrolysis.
Construction:
A single container with an electrolyte (molten ionic compound or aqueous solution).
Two electrodes (inert or active) immersed in the electrolyte.
Anode: Connected to the positive terminal of the external power source (oxidation occurs).
Cathode: Connected to the negative terminal of the external power source (reduction occurs).
Working Principles (e.g., Electrolysis of molten NaCl):
An external power source forces electrons onto the cathode and pulls them from the anode.
At the Anode (positive electrode): Chloride ions (Cl-) lose electrons to form chlorine gas.
2Cl-(l) Cl2(g) + 2e-
At the Cathode (negative electrode): Sodium ions (Na+) gain electrons to form liquid sodium metal.
2Na+(l) + 2e- 2Na(l)
Overall reaction: 2NaCl(l) 2Na(l) + Cl2(g)
Applications:
Electroplating (coating metals with a thin layer of another metal, e.g., chrome plating).
Extraction of reactive metals (e.g., aluminium from bauxite, sodium from molten NaCl).
Purification of metals (e.g., refining of copper).
Production of chemicals (e.g., chlorine gas, sodium hydroxide, hydrogen).
SUMMARY
Electrochemical cells are vital for converting energy between chemical and electrical forms. Galvanic cells produce electricity from spontaneous reactions, forming the basis of batteries. Electrolytic cells use electricity to force non-spontaneous reactions, enabling processes like electroplating, metal extraction, and chemical production. Understanding their construction and working principles is key to appreciating their widespread applications in modern technology and industry.