CBSE SCIENCE CHAPTER 3 NOTES

Metals & Non MetalsĀ 

Metals and Non-metals are two broad categories of elements in the periodic table. The physical properties of these elements are quite distinct from each other. Let’s take a look at the physical properties of metals and non-metals separately:

Physical properties of Metals:

  1. Hardness: The majority of metals are hard, with the exception of alkali metals such as sodium, potassium, lithium, which are very soft and can be easily cut with a knife.
  2. Strength: Most metals are strong and have high tensile strength, which makes them suitable for use in large structures such as copper and iron. However, sodium and potassium are soft metals and have low strength.
  3. State: With the exception of mercury, which is a liquid at room temperature, metals are solid.
  4. Sound: Metals are known for producing a ringing sound, which is why they are referred to as “sonorous.” This property makes metals suitable for use in musical instruments.
  5. Conduction: Metals are good conductors of both heat and electricity. Due to this property, electric wires are typically made of metals like copper and aluminium.
  6. Malleability: Metals are malleable, meaning they can be hammered into thin sheets. This property is why iron is used to make large ships.
  7. Ductility: Metals are ductile, meaning they can be drawn into thin wires. This property is why wires are made of metals.
  8. Melting and Boiling Point: Metals typically have high melting and boiling points, with the exception of sodium and potassium, which have relatively low melting and boiling points.
  9. Density: Most metals have a high density.
  10. Colour: Most metals are grey in colour, although there are exceptions such as gold and copper.

Exceptions:

  1. Mercury (Hg) is the only metal that exists in a liquid state at room temperature.
  2. Alkali metals (group 1 elements) have low densities and are very soft. For example, sodium can be easily cut with a knife.
  3. Group 11 elements, Copper, Silver, and Gold, have low reactivity and do not easily corrode in air or water.

Chemical Properties of Metals
1. Reaction with oxygen: Most of the metals form respective metal oxides when reacting with oxygen.
Metal + Oxygen ā†’ Metal Oxide
Examples:
Reaction of Potassium with Oxygen: Potassium metal forms potassium oxide when reacts with oxygen.

Reaction of Sodium with Oxygen: Sodium metal forms sodium oxide when reacts with oxygen.

Lithium, potassium, sodium, etc. are known as Alkali-metals. Alkali metals react vigorously with oxygen.

Reaction of Copper metal with Oxygen: Copper does not react with oxygen at room temperature but when burnt in air, it gives oxide.

Silver, gold and platinum do not combine with the oxygen of air even at high temperature. They are the least reactive.

2. Reaction of metals with water: When metals react with water, they form metal hydroxide and hydrogen gas. Metal + Water ā†’ Metal hydroxide + Hydrogen Most metals do not react with water, but alkali metals react vigorously with it.

For example, when sodium metal reacts with water, it forms sodium hydroxide and releases hydrogen gas along with a lot of heat. Similarly, calcium metal forms calcium hydroxide and hydrogen gas when it reacts with water. Magnesium metal reacts slowly with water to form magnesium hydroxide and hydrogen gas.

When steam is passed over magnesium metal, it reacts to form magnesium oxide and hydrogen gas. Aluminum metal reacts with water very slowly, but when steam is passed over it, it forms aluminum oxide and hydrogen gas.

Zinc metal reacts with steam to produce zinc oxide and hydrogen gas, but it does not react with cold water. Iron metal reacts with water very slowly, and rust (iron oxide) is formed when it reacts with moisture present in the atmosphere. When steam is passed over iron, it forms iron oxide and hydrogen gas.

Calcium and magnesium metals are heavier than water, but they still float on its surface. This is because when they react with water, hydrogen gas is evolved in the form of bubbles that stick to the metal surface, making it float. Other metals, such as lead, copper, silver, and gold, either do not react with water or react very slowly. Hence, the reactivity of different metals towards water can be arranged in the following order: K > Na > Ca > Mg > Ae > Zn > Fe > Pb > Cu > Ag > Au

Reaction of metal with dilute acid :

Chemical equation for the reaction between sodium metal and dilute hydrochloric acid is:

2Na + 2HCl ā†’ 2NaCl + H2

Where Na is the symbol for sodium, HCl is the chemical formula for hydrochloric acid, NaCl is the chemical formula for sodium chloride, and H2 is the chemical formula for hydrogen gas.

When magnesium reacts with dilute hydrochloric acid, it forms magnesium chloride (MgCl2) and hydrogen gas (H2), which is represented by the chemical equation:

Mg + 2HCl ā†’ MgCl2 + H2

hen zinc reacts with dilute sulfuric acid, zinc sulfate and hydrogen gas are produced. This reaction is commonly used in the laboratory to produce hydrogen gas. The chemical equation for the reaction is:

Zn + H2SO4 ā†’ ZnSO4 + H2

Reaction of Iron with dilute sulphuric acid: Iron reacts with dilute sulphuric acid to form iron(II) sulphate and hydrogen gas.

Fe + H2SO4 ā†’ FeSO4 + H2

This reaction is similar to the reaction of zinc with dilute sulphuric acid, where the metal forms its respective salt and hydrogen gas is produced. Iron(II) sulphate is a pale green solid and is commonly used in fertilizers and water treatment.

When metals are treated with nitric acid (HNO3), hydrogen (H2) gas is not evolved. This is because nitric acid is a strong oxidizing agent that oxidizes the hydrogen gas (H2) liberated into water (H2O) and gets reduced to some oxide of nitrogen like nitrous oxide (N2O), nitric oxide (NO), and nitrogen dioxide (NO2).

Copper, gold, and silver are referred to as noble metals because they do not react with water or dilute acids.

The order of reactivity of metals towards dilute hydrochloric acid or sulphuric acid is as follows: K > Na > Ca > Mg > Al > Zn > Fe > Cu > Hg > Ag

Aquaregia is a highly corrosive and fuming liquid, which is prepared by mixing concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1. It is one of the few reagents that can dissolve gold and platinum.

The reactivity series of metals is a list that arranges metals in the order of their decreasing reactivity with dilute acids, with potassium being the most reactive and gold being the least reactive among the common metals.

Metallergy :

Minerals are naturally occurring compounds that contain metals and other impurities.
Ores are minerals that can be extracted profitably and conveniently to obtain metals.
Gangue is the term used for earthly impurities such as silica and mud that are found along with the ore.
Metallurgy refers to the process used to extract metals in their pure form from their ores.

Enrichment of Ores:

  • The minerals from which metals are extracted profitably and conveniently are called ores.
  • Ores are usually contaminated with impurities like sand, soil, etc., which are known as gangue.
  • The removal of gangue from the ore before extracting the metal is necessary.
  • The method of removing gangue depends on the differences between the physical and chemical properties of the gangue and the ore.

Conversion of Concentrated Ore into Metal:

  • The extraction of a metal from its concentrated ore is essentially a process of reduction of the metal compound present in the ore.
  • The method of reduction used depends on the reactivity of the metal to be extracted

Extraction of Less Reactive Metals:
Metals at the bottom of the reactivity series are not highly reactive, and their oxides can be reduced by heating the ore itself.

Extraction of Mercury:
Mercury is extracted from cinnabar, an ore of mercury, by a two-step process. First, cinnabar is heated in the air to convert it into mercuric oxide and sulfur dioxide gas:

2HgS(s) + 3O2(g) ā†’ 2HgO(s) + 2SO2(g) (heat)

Next, the mercuric oxide is further heated to reduce it to mercury and oxygen gas:

2HgO(s) ā†’ 2Hg(s) + O2(g) (heat)

The extraction of moderately reactive metals

The extraction of moderately reactive metals from their ores involves the reduction of their oxides with carbon, aluminum, sodium, or calcium. The reduction process is simpler with metal oxides than with carbonates or sulfides, so the ores are first converted to metal oxides. This can be achieved through calcination or roasting of the concentrated ores.

Calcination and roasting are the two processes used to convert ores to their corresponding metal oxides. Calcination involves heating a carbonate ore strongly in the absence of air, while roasting involves strongly heating a sulphide ore in the presence of air.

The metal oxides obtained from calcination and roasting are then reduced to their corresponding metals using reducing agents such as carbon, aluminium, sodium, or calcium. For example, zinc oxide is reduced with carbon to obtain the metal zinc. Similarly, iron oxide is reduced with aluminium to obtain the metal iron, which is produced in the molten state due to the heat released.

The reaction between iron (III) oxide and aluminium is also used in the thermite reaction, which is employed to join railway tracks or repair broken machine parts.

Extraction of Highly Reactive Metals
Metals with high reactivity, found at the top of the reactivity series, have a strong affinity for oxygen. Therefore, oxides of these metals such as sodium, magnesium, calcium, and aluminium cannot be reduced by carbon. They are usually obtained by electrolytic reduction.
Sodium, magnesium, and calcium are obtained by the electrolysis of their molten chlorides. The process involves passing an electric current through the molten chloride, which causes the metal ions to move to the negative electrode (cathode) and gain electrons to form the pure metal. Aluminium is obtained by the electrolysis of a molten mixture of aluminium oxide and cryolite. During the process, oxygen ions move towards the positive electrode (anode) and react with the carbon anode to form carbon dioxide, while the aluminium ions move to the negative electrode (cathode) and gain electrons to form pure aluminium.

Sodium metal is extracted by the electrolytic reduction of molten sodium chloride in the Downs process.
2NaCl(l) ā†’ Electrolysis ā†’ 2Na(l) + Cl2(g)
At the cathode: Na+ + e- ā†’ Na
At the anode: 2Cl- ā†’ Cl2 + 2e-

Refining of Metals
ļ‚· Electrolytic refining is the most commonly used method for refining impure metals.
ļ‚· Electrolytic refining involves refining metals by electrolysis. Metals such as copper, zinc, tin, lead, chromium, nickel, silver, and gold are commonly refined electrolytically.
ļ‚· The process of refining impure metals by electrolysis involves:
o Using a thick block of impure metal as the anode.
o Taking a thin strip of pure metal as the cathode.
o Using a water-soluble salt as the electrolyte.
o Passing an electric current through the electrolyte to dissolve the impure metal from the anode into the electrolyte.
o Depositing an equivalent amount of pure metal from the electrolyte onto the cathode.
o Allowing the soluble impurities to go into the solution while the insoluble impurities settle at the bottom of the anode and are known as the ā€˜anode mud.

Corrosion

Corrosion occurs when the surface of a metal is attacked by substances such as air, moisture or other chemicals in the environment. The process of corrosion is commonly observed in metals and alloys. Copper, for instance, develops a greenish deposit on its surface when exposed to moist air, which is actually copper carbonate. Iron, on the other hand, reacts with air moisture to acquire a coating of brown flaky substance known as rust. Rust is a hydrated iron (III) oxide (Fe2O3.xH2O). Rusting of iron requires the presence of air (or oxygen) and water (or moisture) to occur.

The prevention of corrosion involves various steps, including:

  1. Coating: Coating the surface of the metal with a layer of paint, oil, or some other protective material can help prevent the metal from coming into contact with air and moisture. This prevents the process of oxidation and slows down the rate of corrosion.
  2. Galvanization: Galvanization is a process of coating a metal with a layer of zinc. This protects the metal from corrosion as zinc corrodes before the underlying metal. It is commonly used to protect iron and steel.
  3. Alloying: Alloying involves adding other metals or non-metals to the base metal to create an alloy. This improves the metal’s resistance to corrosion.
  4. Cathodic Protection: Cathodic protection involves connecting the metal to be protected to a more reactive metal, such as magnesium or zinc. The more reactive metal corrodes instead of the protected metal, thus preventing its corrosion.
  5. Using inhibitors: Corrosion inhibitors are chemicals that, when added to the environment around the metal, reduce the rate of corrosion. They work by forming a protective layer on the metal surface that prevents further corrosion.

Anodising

Anodising is an electrolytic process in which a metal (usually aluminium) is coated with a protective oxide layer. In this process, the metal is made the anode in an electrolytic cell and an acid solution containing an oxidizing agent is used as the electrolyte. When a current is passed through the cell, oxygen ions from the electrolyte combine with the metal atoms at the surface of the anode, forming a layer of metal oxide. This layer is highly resistant to corrosion and can also be dyed to produce a wide range of colors. Anodising is commonly used to protect aluminium components in various industries, such as aerospace, automotive, and construction.

nodizing is an electrolytic process that involves the formation of an oxide layer on the surface of a metal object. The process is commonly used to protect aluminum and its alloys from corrosion and to improve its wear resistance.

During anodizing, the metal object is made the anode in an electrolytic cell, while a cathode is placed in the electrolyte solution. A direct current is passed through the cell, causing oxygen ions to be attracted to the anode surface. The oxygen ions then react with the metal atoms, forming a layer of metal oxide on the surface of the object.

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