Metals & Non-metals Chapter Notes - Science Class 10

In this Class 10 chapter on "Metals and Non-Metals" we will explore the distinct properties and behaviors that define metals and non-metals. Our study will encompass the reactions of Metals and Non Metals with air, water, and acids, along with their reactivity series and formation of ionic compounds.
Additionally, we will delve into the occurrence of metals in minerals, extraction processes, corrosion prevention, alloy formation, and practical applications of both elements.

Physical Properties

Physical Properties of Metals:

Metals are known for their unique characteristics:

• Luster: Metals have a shiny and reflective surface, often described as metallic luster. This property makes them attractive and valuable for various applications, such as jewelry and coins.
  Metals are Lustrous
• Conductivity: Metals are excellent conductors of heat and electricity. They allow heat and electrical current to flow through them with minimal resistance. This property makes them essential in electrical wiring and cooking utensils.
  Metals conduct electricity due to presence of free electrons
• Malleability: Metals can be hammered or rolled into thin sheets without breaking. This property is known as malleability and allows metals to be shaped into various forms, like aluminum foils and copper sheets.
  Malleability
• Ductility: Metals can be drawn into thin wires without breaking. This property is called ductility and is crucial for making wires for electrical purposes, such as copper wires in electrical cables.

Ductility

• Density: Metals are generally dense materials, meaning they have a relatively high mass for their volume. This property makes them suitable for making heavy machinery and structural components.
  
• Melting and Boiling Points: Metals usually have high melting and boiling points. For example, iron melts at around 1,535 degrees Celsius. The metals have a closely packed structure and have a strong intermolecular force of attraction. To break these intermolecular force of attraction, more heat is required and therefore they have a high melting point except for some metals like gallium and caesium.

Physical Properties of Non-Metals:

Non-metals, on the other hand, have properties that are quite different from metals:

• Lack of Luster: Non-metals do not have the shiny appearance of metals; instead, they can be dull or have various colors.
  Non Metals Lack Lustre
• Poor Conductors: Non-metals are generally poor conductors of heat and electricity. They do not allow the easy flow of electrical current and heat.
  Sulfur is poor conductor of heat and electricity
• Brittleness: Many non-metals are brittle and tend to break or crumble when subjected to force. For example, sulfur and phosphorus are brittle non-metals. Non Metals are not Malleable or Ductile.
  Non Metals are Brittle i.e. Crumble into Pieces
• Low Density: Non-metals usually have lower density compared to metals. This means they have a lower mass for a given volume.
• Low Melting and Boiling Points: Non-metals often have low melting and boiling points. Non-metals are held by weak intermolecular force of attraction, so less heat required to break them. As a result, they have low melting points.
• Variety of States: Non-metals can exist in different states at room temperature. For example, oxygen is a gas, while bromine is a liquid, and sulfur is a solid.

Understanding these physical properties is essential as they form the basis for classifying elements as metals or non-metals. These properties also determine the various applications of these elements in our daily lives and various industries.
Below is a summary table containing the key properties of both Metals and Non-Metals for your reference.

Chemical Properties of Metals

Reaction of metals with air

Metals combine with oxygen to form metal oxide.
Metal + O₂ → Metal oxide
Examples:

The reactivity of different metals is different with O₂.

• Na and K react so vigorously that they catch fire if kept open so they are kept immersed in kerosene.
• Surfaces of Mg, Al, Zn, and Pb are covered with a thin layer of oxide which prevents them from further oxidation.
• Fe does not burn on heating but iron fillings burn vigorously.
• Cu does not burn but is coated with black copper oxide.
• Au and Ag do not react with oxygen.
• Amphoteric Oxides: Metal oxides that react with both acids, as well as bases to produce salts and water, are called amphoteric oxides.
  Examples:
  

Reaction of metals with water

• Metal + Water → Metal oxide + Hydrogen
• Metal oxide + Water → Metal hydroxide
  

Examples:
(i) 2Na + 2H₂O → 2NaOH + H₂ + Heat
(ii) Ca + 2H₂O → Ca(OH)₂ + H₂
(iii) Mg + 2H₂O → Mg(OH)₂ + H₂
(iv) 2Al + 3H₂O → Al₂O₃ + 3H₂
(v) 3Fe + 4H₂O → Fe₃O₄ + 4H₂

Reaction of metals with acids (Dilute)

• Metal + Dilute acid → Salt + H²
• Cu, Ag, and Hg do not react with dil. acids.

Examples
(i) Fe + 2HCl → FeCl₂ + H₂
(ii) Mg + 2HCl → MgCl₂+ H₂
(iii) Zn + 2HCl → ZnCl₂ + H₂
(iv) 2Al + 6HCl → 2AlCl₃ + 3H₂

Reaction of Metals with Solutions of other Metal Salts 

• Metal A + Salt solution B → Salt solution A + Metal B
•  Reactive metals can displace less reactive metals from their compounds in solution form.
  Fe + CuSO₄ →  FeSO₄ + Cu

The Reactivity Series

The reactivity series is a list of metals arranged in the order of their decreasing activities.

Reactivity Series

Reaction of Metals with Non-metals

• Reactivity of elements is the tendency to attain a filled valence shell.
• Atoms of the metals lose electrons from their valence shell to form cations. Atoms of non-metals gain electrons in the valence shell to form an anion.
  E.g.: Formation of NaCl
  
  Sodium cation
  
  Chloride anion
  

Ionic compounds

The compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds.

Formation of Ionic Compounds

Properties of Ionic Compounds

• Physical nature: Ionic compounds are solid substances and possess a certain level of hardness due to the powerful attraction between the positively and negatively charged ions. Typically, these compounds exhibit brittleness and tend to fracture into fragments when subjected to pressure.
• Melting and Boiling Point: Ionic compounds exhibit high melting and boiling points due to the presence of strong inter-ionic attractions, which require a significant amount of energy to be overcome.
• Solubility: Electrovalent compounds usually dissolve in water but do not dissolve in solvents like kerosene or petrol.
• Conduction of electricity: When an ionic compound is dissolved in water, the resulting solution contains ions that can migrate towards the electrodes when an electric current is applied. In contrast, solid ionic compounds do not conduct electricity because the ions are unable to move due to their fixed arrangement. However, when ionic compounds are in the molten state, they can conduct electricity because the heat overcomes the electrostatic forces holding the oppositely charged ions together. As a result, the ions become mobile and can freely move, facilitating the conduction of electricity.

Occurrence of Metals

• Minerals: The elements or compounds which occur naturally in the earth’s crust are called minerals.
• Ores: Minerals that contain a very high percentage of particular metal and the metal can be profitably extracted from it, such minerals are called ores.

Extraction of Metals from Ores

Step 1. Enrichment of ores.
Step 2. Extraction of metals.
Step 3. Refining of metals. 

Enrichment of Ores

Enrichment of ores refers to the process of eliminating impurities, such as soil, sand, stone, and silicates, from mined ores. Ores commonly contain a variety of unwanted substances called gangue. The primary objective of concentration is to separate these impurities from the ores. Concentration methods rely on the physical and chemical properties of the ores. Examples of such processes include gravity separation, electromagnetic separation, and froth flotation.

Extracting Metals Low in Activity Series

During self-reduction, when sulphide ores of metals such as Hg, Pb, Cu, etc., are heated in the presence of air, a portion of the ore is converted into an oxide. This oxide then reacts with the remaining sulphide ore to produce crude metal and sulphur dioxide. Notably, no external reducing agent is employed in this process.

• 2HgS(Cinnabar)+3O₂(g)+heat→2HgO(crude metal)+2SO₂(g)
  2HgO(s)+heat→2Hg(l)+O₂(g)
• Cu₂S(Copper pyrite)+3O₂(g)+heat→2Cu₂O(s)+2SO₂(g)
  2Cu₂O(s)+Cu₂S(s)+heat→6Cu(crude metal)+SO₂(g)
• 2PbS(Galena)+3O₂(g)+heat→2PbO(s)+2SO₂(g)
  PbS(s)+2PbO(s)→2Pb(crude metal)+SO₂(g)

Extracting Metals in the Middle of Activity Series

• Metals in the middle of the activity series, such as iron, zinc, lead, and copper, have a moderate level of reactivity. These metals are commonly found in nature as sulphides or carbonates. 
• Obtaining these metals from their oxides is easier compared to extracting them from sulphides and carbonates. 
• Therefore, before the reduction process, sulphide ores and carbonate ores need to be converted into metal oxides. 
• Sulphide ores are transformed into oxides through intense heating in the presence of an excess of air, a process called roasting. 
• Carbonate ores are converted into oxides by strong heating in a limited air environment, known as calcination.

The Aluminothermic reaction, also called the Goldschmidt reaction, is a strongly exothermic process where aluminium is heated together with metal oxides, typically iron (Fe) and chromium (Cr), at elevated temperatures.
Fe₂O₃+2Al→Al₂O₃+2Fe+heat
Cr₂O₃+2Al→Al₂O₃+2Cr+heat

Extracting Metals towards the Top of the Activity Series

• Metals that are high up in the reactivity series are extremely reactive. They cannot be extracted from their compounds through the process of heating with carbon. 
• For instance, carbon is unable to reduce the oxides of metals like sodium, magnesium, calcium, and aluminum to obtain the pure metals. 
• The reason behind this is that these metals have a stronger attraction to oxygen than carbon does. 
• Therefore, these metals are obtained through electrolytic reduction, where the process involves the electrolysis of their molten chlorides. 
• During this process, the metals are deposited at the cathode (the negatively charged electrode), while chlorine gas is liberated at the anode (the positively charged electrode).
• At the cathode (reduction):
  Na⁺(molten)+e⁻→Na(s)
  Metal is deposited
• At the anode (oxidation):
  2Cl⁻(molten)→Cl₂(g)+2e^–
  Chlorine gas is liberated.

Important terms
(a) Gangue: Ores are usually contaminated with large amounts of impurities such as soil, sand etc. called gangue.
(b) Roasting: The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is called roasting.
2ZnS + 3O₂ →(Heat) 2ZnO + 2SO₂
(c) Calcination: The carbonate ores are changed into oxides by heating strongly in limited air. This process is called calcination.
ZnCO₃ →(Heat) ZnO + CO₂
(d) Reduction: Metal oxides are reduced to corresponding metals by using a reducing agent like carbon.
ZnO + C → Zn + CO

Refining of metals

The most widely used method for refining impure metal is electrolytic refining.

Electrolytic Refining

(i) Anode: Impure copper
(ii) Cathode: Strip of pure copper
(iii) Electrolyte: Solution of acidified copper sulphate

• On passing the current through the electrolyte, the impure metal from the anode dissolves into the electrolyte.
• An equivalent amount of pure metal from the electrolyte is deposited at the cathode.

  The insoluble impurities settle down at the bottom of the anode and are called anode mud.

Corrosion

The surface of some metals gets corroded when they are exposed to moist air for a long period. This is called corrosion.

CorrosionExamples:
(i) Silver becomes black when exposed to air as it reacts with air to form a coating of silver sulphide.
(ii) Copper reacts with moist carbon dioxide in the air and gains a green coat of copper carbonate.
(iii) Iron when exposed to moist air acquires a coating of a brown flaky substance called rust.

• Prevention of Corrosion
  • The rusting of iron can be prevented by painting, oiling, greasing, galvanizing, chrome plating, anodizing or making alloys.

• Galvanization: It is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc.
• Alloy: An alloy is a homogeneous mixture of two or more metals or a metal and a non-metal.
• Examples of alloy:
  (i) Iron: Mixed with a small amount of carbon becomes hard and strong.
  (ii) Steel: Iron + Nickel and chromium
  (iii) Brass: Copper + Zinc
  (iv) Bronze : Copper + Tin (Sn)
  (v) Solder: Lead + tin
  (vi) Amalgam: If one of the metals is mercury (Hg).

Frequently Asked Questions (FAQs)

Q1. What are the physical properties of metals?
Ans. Metals are typically hard, dense, shiny, malleable, ductile, and good conductors of heat and electricity. They have a high boiling and melting point and are usually solid at room temperature. $$$

Q2. What are the physical properties of non-metals?
Ans. Non-metals are usually dull, brittle, and poor conductors of heat and electricity. They have low boiling and melting points and can exist in all three states of matter at room temperature. Non-metals are not malleable or ductile. 

Q3. What is the difference between metals and non-metals?
Ans. The primary difference between metals and non-metals is their physical properties. Metals are usually hard, dense, and good conductors of heat and electricity, while non-metals are brittle, dull, and poor conductors of heat and electricity. Additionally, metals tend to form cations, while non-metals form anions. 

Q4. What are some common uses of metals?
Ans. Metals have a variety of uses, including construction, transportation, electrical wiring, and the production of tools and appliances. Some metals, such as gold and silver, are used in jewelry and currency. Iron and steel are commonly used in construction and manufacturing. 

Q5. What are some common uses of non-metals?
Ans. Non-metals are used in a variety of applications, such as the production of plastics, ceramics, and semiconductors. Non-metals such as oxygen and nitrogen are essential for life and are present in the atmosphere. Non-metals such as sulfur and chlorine are used in the production of a variety of chemicals.