11.0     DEFINITION OF ELECTROLYSIS AND OTHER TERMS INVOLVED

ELECTROLYSIS – It can be defined as the chemical decomposition of a compound (electrolyte) into its component ions by the passage of electric current through a solution of that compound or the molten compound.

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ELECTROLYTE – It is a compound which when molten or in aqueous solution will conduct electricity and decomposes by it. An electrolyte could be classified into three:

• Strong electrolyte e.g. NaCl, NaOH, HCl, HNO₃ etc.
• Weak electrolyte e.g. ethanoic acid and its salts, NH₃ solution etc.
• Non-electrolyte e.g. benzene, sugar, ethanol, trichloromethane etc.

ELECTRODE – It is a conductor in the form of wire, plates or rods, through which an electric current enters or leaves the electrolyte. It is divided into two; anode and cathode.

          11.1     TYPES OF ELECTRIC CELL

• Electrolytic cell – A device used for electrolysis, consisting of an electrolyte, electrodes, and an external power source.
• Electrochemical cell – A device/cell that converts chemical energy to electrical energy. Examples include galvanic or voltaic cell.

         11.2     FARADAY’S LAW OF ELECTROLYSIS

1^ST LAW – It states that the mass of a substance deposited at the electrode during electrolysis, is directly proportional to the quantity of electricity passed through the electrolyte.

• Mass(M) = Z × Current(I) × Time(t)
• Z = Electrochemical equivalent (Z = R.A.M / CF; where F = 96500C)
• Faraday’s constant (F) represents the charge of one mole of electrons, approximately 96,500 C/mol. One faraday discharge one mole of univalent element, A⁺, while two moles of faraday discharges one mole of divalent element, B²⁺.

2^ND LAW – It states that when the quantity of electricity is passed through different electrolytes, the number of moles of the elements deposited are inversely proportional to the charges on the ions of the elements.

         11.3     FACTORS AFFECTING PREFERENTIAL DISCHARGE OF IONS IN ELECTROLYSIS

• Position of the ion in the electrochemical series.
• Concentration of the ion in the electrolyte.
• Nature of electrode used.

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         11.4     ELECTROLYSIS OF SOME SUBSTANCES

SUITABLE ELECTRODES FOR DIFFERENT ELECTROLYTES

• INERT ELECTRODES – Do not react with the electrolyte or products, such as platinum (Pt) or graphite (carbon). Suitable for electrolysis of solutions like NaCl, H₂SO₄, CuSO₄.
• ACTIVE ELECTRODES – Participate in the reaction, such as copper (Cu) electrodes in CuSO₄ electrolysis, where copper dissolves from the anode.

         11.5     APPLICATIONS OF ELECTROLYSIS

• Purification of elements/metals e.g copper.
• Production of elements and compounds (Al, Na, O2, Cl2 and NaOH).
• Electroplating/coating of metals to prevent corrosion.
• Extraction of metals from ores (e.g., aluminum, copper).
• Electrolytic reduction of metals from their compounds.

         11.5     REDOX SERIES

The redox series, also known as the electrochemical series or activity series, is a list of elements arranged according to their standard electrode potentials. It ranks elements (and their ions) based on how easily they ionize. The redox series are as follows: K, Ca, Na, Mg, Al, Zn, Fe, Sn, Pb, H, Cu, Hg, Ag, Au, and Pt.

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• Elements at the top of the redox series (like potassium, calcium, and sodium) ionize easily by losing electrons, making them strong reducing agents. These elements are more reactive, have less standard electrode potential and are easily oxidized (acts as anode).

• Elements at the bottom of the series (like copper, silver, gold and platinum) ionize less easily. They tend to gain electrons, acting as oxidizing agents, have high standard electrode potential and are less reactive. These elements are easily reduced (acts as cathode).

         11.6     HALF-CELL REACTIONS AND ELECTRODE POTENTIALS

A HALF-CELL REACTION represents either the oxidation or the reduction process that occurs in an electrochemical cell. There are two half-cells involved redox (reduction-oxidation) reaction. They are:

Oxidation half-cell reaction – In this reaction, an element loses electrons (oxidized). This occurs at the anode of the cell. Example: Zinc loses two electrons and becomes a zinc ion.

Reduction half-cell reaction – In this reaction, an element gains electron (reduced). This occurs at the cathode of the cell. Example: Copper ion gains two electrons and becomes copper metal.

When combined, these two half-cell reactions form the complete redox reaction, where electrons are transferred from the anode to the cathode through an external circuit, generating electrical current.

THE STANDARD ELECTRODE POTENTIAL of metal ions/metal system is the potential difference set up between the metal and a one-molar solution of its ions at 25^oC, (arbitrarily taking the standard electrode potential of the hydrogen ions/hydrogen gas system as zero volt).

The potential difference set up between a metal and its solution is called the electrode potential for that metal ions/metal system.

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          11.7     CORROSION AS AN ELECTROLYTIC PROCESS

Corrosion occurs when metals, especially iron, react with environmental factors (e.g., oxygen and moisture) in an electrolytic process, forming rust (Fe₂O₃·nH₂O). it could be prevented by using reactive metals like zinc for protection, applying paint to serve as shield, electroplating to create barriers, and using grease or oil to block moisture and prevent rust formation.