Extraction of Highly Reactive Metals
A category of metal elements known as reactive metals can react with acids, water, mineral acids, and severe oxidising acids. The activity or reactivity series, which ranks the most reactive metals from highest to lowest, can be used to identify this group.
The reactivity of metals increases and loses electrons to form positive ions as the series progresses from the bottom to the top. Those at the top of the series tarnish and corrode faster and require more energy to separate and become electron donors.
What is the Reactivity Series of Metals?
The metal reactivity series, also known as the activity series, is the arrangement of metals in descending order of reactivity. The reactivity series data can be used to predict whether one metal can displace another in a single displacement reaction. It can also be used to determine the reactivity of metals to water and acids.
The first five items on the chart are regarded as highly reactive. These substances can react with hot or cold water to produce steam containing hydroxides and hydrogen gas. Following the list from magnesium to chromium, the next four metals are considered active metals because they can react with hot water or steam to produce hydrogen gas and oxides. Hydrogen ion reduction is resistant to all oxides formed under the first two metal groups.
From iron to lead, the six elements can recycle hydrogen from nitric, sulfuric, and hydrochloric acids. These metals’ oxides can be reduced by heating them with carbon, hydrogen gas, or carbon monoxide.
Order of Decreasing Reactivity of metals:
Caesium (Cs) > Potassium (K) > Sodium (Na) > Lithium (Li) > Calcium (Ca) > Magnesium (Mg) > Aluminum (Al) > Zinc (Zn) > Iron (Fe)
Extraction of Highly Reactive Metals
Electrolytic reduction is used to extract metals with high reactivity from their ores, such as sodium, calcium, magnesium, and aluminium. As carbon is less reactive than these metals, it cannot be used to reduce them.
- Electrolytic Reduction
Electrolytic reduction is a type of electrolysis in which an electric current is passed through a molten or dissolved ionic substance, causing the electrodes to chemically react and the materials to decompose. The hydroxides, oxides, and chlorides of metals in their combined state are electrically reduced using this process.
Metals are collected at the cathode. Some metals, such as K, Na, and Al, are obtained through the electrolytic reduction process. When an electric current is passed through molten sodium chloride or a solution of sodium chloride, sodium metal is deposited over the cathode.
Na+ + e− ⇢ Na
2Cl− − e− ⇢ Cl2
2NaCl ⇢ 2Na + Cl2
Electrolysis of Aqueous Sodium Chloride: In an aqueous solution, sodium chloride is dissociated and exists as sodium and chloride ions. In an aqueous solution, the electrolysis of sodium chloride is simpler. Water, on the other hand, can undergo reduction and oxidation reactions at various potentials. As a result, the substance that is oxidized or reduced is not only sodium and chloride ions, but also the water molecule. At both the cathode and the anode, two competing reactions are possible.
At Cathode: A reduction reaction occurs when the pH is 7. Water can be converted to hydrogen gas, and sodium ions can be converted to sodium metal.
2H2O (l) + 2e– → H2 (g) + 2OH–
Na+(l) + e– → Na (l)
At anode: With a pH of 7, the oxidation reaction occurs. Water can be oxidised to produce oxygen, or a chloride ion can be oxidised to produce a chlorine molecule.
2H2O → O2(g) + 4H+
2Cl– → Cl2 + 2e–
As a result, the product of aqueous sodium chloride electrolysis can be anything between sodium metal or hydrogen gas at the cathode and chlorine or oxygen gas at the anode, with a sodium hydroxide byproduct resulting from the reaction of sodium and water. The electrolysis product is determined by the concentration of sodium chloride aqueous solution.
- Electrolytic Reduction of Alumina
The negative electrode is a steel container that has been coated with carbon. Aluminium oxide is a type of ionic compound. The Al+3 and O-2 ions are free to move and conduct electricity when it is melted. The cathode of the electrolysis of the alumina/cryolite solution yields aluminium and the anode yields oxygen.
Cathode Reaction: 4Al+3 + 12e– → 4Al
Anode Reaction: 6O-2 – 12e– → 3O2
Because aluminium is denser than alumina, it settles to the bottom of the cell and can be tapped off as pure liquid metal. At the positive carbon anode, oxygen is released. Carbon dioxide is also produced at the carbon anode as a result of hot oxygen reacting with the carbon anode to produce carbon dioxide gas.
Carbon + Oxygen → Carbon dioxide.
C (s) + O2 (g) → CO2 (g)
Carbon anodes gradually degrade because each molecule of carbon dioxide emitted carries a small amount of carbon with it. When the carbon anodes become too small, they must be replaced.
Question 1: What does a reactivity series show?
Metals are arranged in descending order of reactivity in the reactivity series. The reactivity of a metal can be determined by studying its reactions to competition and displacement.
Question 2: What metal is the least reactive?
Transition metals are elements in the periodic table that are much less reactive, and metals like gold and platinum are at the bottom of the list, exhibiting little chemical reaction with any common reagents.
Question 3: Which is the most reactive element?
Alkali metals are the most reactive elementary group (situated far apart from intermediate metals and noble gases). Cesium is second from the bottom of this group, has six electron shells, and has the characteristics of a reactive atom, making it the most reactive element.
Question 4: Why electrolytic reduction is done to obtain aluminium?
Since aluminium has a high affinity for oxygen, it is a stable compound. It is resistant to common reducing agents such as carbon. As a result, electrolytic reduction is used to obtain aluminium.
Question 5: Define electrolytic reduction.
Electrolytic reduction is the process by which highly reactive metals are removed from their compound state. It is usually done to obtain the purest form of the desired metal. Magnesium and aluminium are two examples.