Metals and Non-metals For Class 10 Science Summary notes

Physical Properties

Any measurable property whose value describes a condition of a physical system is referred to as

 a physical property. A system’s physical properties can be utilized to characterize its transitions between brief states. 

Observables are a term used to describe physical qualities. 

Physical Properties of Metals

● Hard and have a high tensile strength – Carbon is the only nonmetal with very high tensile strength.

● Solid at room temperature – One non-metal, bromine, is a liquid at room temperature. The other non-metals are solids at room temperature, including carbon and sulfur.

● Sonorous – Metals produce a typical ringing sound when hit by something.

● Good conductors of heat and electricity – Graphite is a good conductor of heat and electricity.

● Malleable, i.e., can be beaten into thin sheets.

● Ductile, i.e., can be drawn into thin wires.

● High melting and boiling points (except Caesium (Cs) and Gallium (Ga)) – Graphite, a form of carbon (a non-metal), has a high boiling point and exists in the solid state at room temperature.

● Dense, (except for alkali metals). Osmium – has the highest density and lithium – has the least density.

● Lustrous – Metals have the quality of reflecting light from their surface and can be polished e.g. gold, silver, and copper. Iodine and carbon are non-metals that are lustrous. Note that carbon is lustrous only in certain forms like diamond, and graphite.
● Silver-grey in color, (except gold and copper) – Metals usually have a silver or grey color.

Non-Metals

Non-metals are those elements that do not exhibit the properties of metals.

Physical Properties of Non-metals

• Occur as solids, liquids, and gases at room temperature
• Brittle
• Non-malleable
• Non-ductile
• Non-sonorous
• Bad conductors of heat and electricity

Exceptions in Physical Properties

• Alkali metals (Na, K, Li) can be cut using a knife.
• Mercury is a liquid metal.
• Lead and mercury are poor conductors of heat.
• Mercury expands significantly for the slightest change in temperature.
• Gallium and cesium have a very low melting points.
• Iodine is non-metal but it has lustre.
• Graphite conducts electricity.
• Diamond conducts heat and has a very high melting point.

Examples of Non-metals

1. Hydrogen – Gas
2. Nitrogen – Gas
3. Oxygen – Gas
4. Fluorine – Gas
5. Chlorine – Gas
6. Bromine – Liquid
7. Iodine – Solid
8. Carbon – Solid
9. Sulphur – Solid
10. Phosphorous – Solid
11. Silicon – Solid

Chemical Properties

Chemical Properties of Metals

● Alkali metals (Li, Na, K, etc) react vigorously with water and oxygen or air.
● Mg reacts with hot water.
● Al, Fe, and Zn react with steam.
● Cu, Ag, Pt, and Au do not react with water or dilute acids.

The reaction of Metals with Oxygen (Burnt in Air)

Metal oxide is formed when metals are burned in the air and react with oxygen in the air. Metal oxides are a type of basic material found in nature. They change the color of red litmus to blue. To avoid reactions with oxygen, moisture, and carbon dioxide in the air, sodium, and potassium metals are kept in kerosene oil.

Metal + Oxygen→ Metal oxide (basic)
● Na and K are kept immersed in kerosene oil as they react vigorously with air and catch fire.
4K (s)+O₂ (g)→2K₂O (s) (vigorous reaction)
● Mg, Al, Zn, and Pb react slowly with air and form a protective layer that prevents corrosion.

2Mg (s)+O₂ (g)→2MgO (s) (Mg burns with white dazzling light)

4Al (s)+3O₂ (g)→2Al₂O₃ (s)

● Silver, platinum, and gold don’t burn or react with air.

Basic Oxides of Metals

Metal oxides are crystalline solids that contain a metal cation and an oxide anion. They typically react with water to form bases or with acids to form salts.

MO + H₂O → M(OH)₂ (where M = group 2 metal) .

Thus, these compounds are often called basic oxides.

Some metallic oxides get dissolved in water and form alkalis. Their aqueous solution turns red litmus blue.

Na₂O (s)+H₂O (l)→2 NaOH (aq)

K₂O (s)+H₂O (l)→2 KOH (aq)

Amphoteric Oxides of Metals

Amphoteric oxides are metal oxides that react with both acids as well as bases to form salt and water.
For example – Al₂O₃, ZnO, PbO, SnO

Al₂O₃ (s)+6HCl (aq)→2AlCl₃ (aq)+3H₂O (l)

Al₂O₃ (s)+2NaOH (aq)→2NaAlO₂ (aq)+H₂O (l)

ZnO (s)+2HCl (aq)→ZnCl₂ (aq)+H₂O (l)

ZnO (s)+2NaOH (aq)→Na₂ZnO₂ (aq)+H₂O (l)

Reactivity Series

The reactivity series of metals, also known as the activity series, refers to the arrangement of metals in the descending order of their reactivities.

The below table illustrates the reactivity of metals from high order to low order.

Symbol	Element
K	Potassium ( Highly Active Metal)
Ba	Barium
Ca	Calcium
Na	Sodium
Mg	Magnesium
Al	Aluminium
Zn	Zinc
Fe	Iron
Ni	Nickel
Sn	Tin
Pb	Lead
H	Hydrogen
Cu	Copper
Hg	Mercury
Ag	Silver
Au	Gold
Pt	Platinum

Roasting

Converts sulphide ores into oxides on heating strongly in the presence of excess air.
It also removes volatile impurities.

2 ZnS (s)+3O₂ (g) + Heat → 2ZnO (s)+2SO₂ (g)

Calcination

Converts carbonate and hydrated ores into oxides on heating strongly in the presence of limited air. It also removes volatile impurities.

ZnCO₃ (s) + heat → ZnO (s)+ CO₂ (g)

CaCO₃ (s) + heat → CaO (s) + CO₂ (g)

Al₂O₃.2H₂O(s) + heat → 2Al₂O₃ (s) + 2H₂O (l)

2Fe₂O₃.3H₂O(s) + heat → 2Fe₂O₃ (s) + 3H₂O (l)

Reaction of Metals with Acid

When a metal is immersed in acid, it becomes smaller and smaller as the chemical process consumes it. Gas bubbles can also be detected at the same moment. Hydrogen gas bubbles are formed as a result of the reaction. Because hydrogen is combustible, this can be demonstrated with a burning splint.

Metal + diluteacid → Salt + Hydrogen gas

2Na (s) + 2HCl (dilute) → 2NaCl (aq) + H₂ (g)

2K (s) + H₂SO₄ (dilute) → K₂SO₄ (aq) + H₂ (g)

Only Mg and Mn, react with very dilute nitric acid to liberate hydrogen gas.

Mg (s) + 2HNO₃ (dilute) → Mg(NO₃)₂ (aq) + H₂ (g)

Mn (s) + 2HNO₃ (dilute) → Mn(NO₃)₂ (aq) + H₂ (g)

Displacement Reaction

A more reactive element displaces a less reactive element from its compound or solution.

How Do Metals React with Solution of Other Metal Salts

A more reactive metal can displace a less reactive metal from its salt solution in a displacement reaction. Metal displacement reaction is a common name for this reaction. The reactivity of certain regularly used metals has been ordered in decreasing order. This is referred to as the reactivity or activity series.

Metal A+Salt of metal B → Salt of metal A + Metal B

Fe (s) + CuSO₄ (aq) → FeSO₄ (aq) + Cu (s)
Cu(s) + 2AgNO₃ (aq) → Cu(NO₃) (aq) + 2Ag (s)

• It’s a component of thermite welding. Aluminium displaces iron from its oxide in this process.
• It is used in the production of steel. In which iron is displaced from its oxide by carbon.
• It is mostly utilised in metal extraction.

Reaction of Metals with Bases

The base has a bitter taste and a slippery texture. A base dissolved in water is called an alkali. When chemically reacting with acids, such compounds produce salts. Bases are known to turn blue on red litmus paper.

Base + metal → salt + hydrogen
2NaOH(aq)+Zn(s) → Na₂ZnO₂(aq)+H₂(g)
2NaOH(aq)+2Al(s)+2H₂O(l) → 2NaAlO₂(aq)+2H₂(g)

Extraction of Metals and Non-Metals

Applications of Displacement Reaction

Uses of displacement reaction

1. Extraction of metals
2. Manufacturing of steel
3. Thermite reaction: Al(s) + Fe₂O₃(s) → Al₂O₃ + Fe (molten)

The thermite reaction is used in the welding of railway tracks, cracked machine parts, etc.

Occurrence of Metals

Most of the elements, especially metals occur in nature in the combined state with other elements. All these compounds of metals are known as minerals. But out of them, only a few are viable sources of that metal. Such sources are called ores.
Au, Pt – exists in the native or free state.

Extraction of Metals

The process of extracting metal ores buried deep underground is called Mining. The metal ores are found in the earth’s crust in varying abundance. The extraction of metals from ores is what allows us to use the minerals in the ground! The ores are very different from the finished metals that we see in buildings and bridges. Ores consist of the desired metal compound and the impurities and earthly substances called Gangue.

Metals of high reactivity – Na, K, Mg, Al.

Metals of medium reactivity – Fe, Zn, Pb, Sn.

Metals of low reactivity – Cu, Ag, Hg

Extracting Metals Low in Reactivity Series

By self-reduction- when the sulphide ores of less electropositive metals like Hg, Pb, Cu etc., are heated in air, a part of the ore gets converted to oxide which then reacts with the remaining sulphide ore to give the crude metal and sulphur dioxide. In this process, no external reducing agent is used.

1. 2HgS (Cinnabar) + 3O₂(g) + heat → 2HgO (crude metal) + 2SO₂ (g)
2 HgO (s) + heat → 2Hg(l) + O₂ (g)

2. Cu₂S(Copperpyrite)+3O₂(g)+heat→2Cu₂O(s)+2SO₂(g)
2Cu₂O(s)+Cu₂S(s)+heat→6Cu(crude metal)+SO₂(g)

3. 2PbS (Galena) + 3O₂ (g) + heat → 2PbO (s) + 2SO₂ (g)
PbS (s) + 2PbO (s) → 2Pb (crudemetal) + SO₂ (g)

Extracting Metals in the Middle of Reactivity Series

Calcination is a process in which ore is heated in the absence of air or air might be supplied in limited quantity. Roasting involves heating of ore lower than its melting point in the presence of air or oxygen. Calcination involves thermal decomposition of carbonate ores.

Smelting – involves heating the roasted or calcined ore (metal oxide) to a high temperature with a suitable reducing agent. The crude metal is obtained in its molten state.

Fe₂O₃ + 3 C (coke) → 2 Fe + 3 CO₂

Aluminothermic reaction – also known as the Goldschmidt reaction is a highly exothermic reaction in which metal oxides usually Fe and Cr are heated to a high temperature with aluminium.

Fe₂O₃ + 2Al → Al₂O₃ + 2Fe + heat
Cr₂O₃ + 2Al → Al₂O₃ + 2Cr + heat

Extraction of Metals Towards the Top of the Reactivity Series

Electrolytic reduction:

1. Down’s process: Molten NaCl is electrolyzed in a special apparatus.

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.

2. Hall’s process: Mixture of molten alumina and a fluoride solvent usually cryolite, (Na₃AlF₆) is electrolysed.

At the cathode (reduction):

2 Al ³⁺+ 6 e ^–→ 2Al (s)
Metal is deposited.

At the anode (oxidation):

6O₂ ^– → 3O₂ (g) + 12 e ^–

Oxygen gas is liberated.

The metals at the top of the reactivity series are highly reactive. They cannot be obtained from their compounds by heating with carbon, because these metals have more affinity for oxygen than carbon. Hence, for the extraction of such metals electrolytic reduction method is used.

Refining of Metals

Refining of metals – removing impurities or gangue from crude metal. It is the last step in metallurgy and is based on the difference between the properties of metal and the gangue.

Electrolytic Refining

Metals like copper, zinc, nickel, silver, tin, gold, etc., are refined electrolytically.
Anode: impure or crude metal
Cathode: a thin strip of pure metal
Electrolyte: aqueous solution of metal salt

From anode (oxidation): metal ions are released into the solution
At the cathode (reduction): the equivalent amount of metal from the solution is deposited
Impurities deposit at the bottom of the anode.

Electronic Configuration

Group 1 elements – Alkali metals

Element	Electronic Configuration
Lithium(Li)	2,1
Sodium(Na)	2,8,1
Potassium(K)	2,8,8,1
Rubidium(Rb)	2,8,18,8,1

Group 2 elements – Alkaline earth metals

Element	Electronic Configuration
Beryllium(Be)	2,2
Magnesium(Mg)	2,8,2
Calcium(Ca)	2,8,8,2
Stronium(Sr)	2,8,18,8,2

How Do Metals and Nonmetals React

Metals lose valence electron(s) and form cations.
Non-metals gain those electrons in their valence shell and form anions.
The cation and the anion are attracted to each other by strong electrostatic force, thus forming an ionic bond.
For example: In calcium chloride, the ionic bond is formed by oppositely charged calcium and chloride ions.
The calcium atom loses 2 electrons and attains the electronic configuration of the nearest noble gas (Ar). By doing so, it gains a net charge of +2.

The two Chlorine atoms take one electron each, thus gaining a charge of -1 (each) and attain the electronic configuration of the nearest noble gas (Ar).

Ionic Compounds

Ionic compounds are neutral compounds that are made up of positively charged cations and negatively charged anions. Binary ionic compounds (ionic compounds containing only two types of elements) are named by first writing the name of the cation, then the name of the anion.

The electrostatic attractions between the oppositely charged ions hold the compound together.
Examples: MgCl₂, CaO, MgO, NaCl, etc.

Properties of Ionic Compound

Ionic compounds

1. Are usually crystalline solids (made of ions).
2. Have high melting and boiling points.
3. Conduct electricity when in an aqueous solution and when melted.
4. Are mostly soluble in water and polar solvents.

Physical Nature

Ionic solids usually exist in regular, well-defined crystal structures.

Electric Conduction of Ionic Compounds

Ionic compounds conduct electricity in the molten or aqueous state when ions become free and act as charge carriers.
In solid form, ions are strongly held by electrostatic forces of attraction and are not free to move; hence do not conduct electricity.

For example, ionic compounds such as NaCl do not conduct electricity when solid but when dissolved in water or in a molten state, they will conduct electricity.

Melting and Boiling Points of Ionic Compounds

In ionic compounds, the strong electrostatic forces between ions require a high amount of energy to break. Thus, the melting point and boiling point of an ionic compound are usually very high.

Solubility of Ionic Compounds

Most ionic compounds are soluble in water due to the separation of ions by water. This occurs due to the polar nature of water.
For example, NaCl is a 3-D salt crystal composed of Na⁺ and Cl⁻ ions bound together through electrostatic forces of attraction. When a crystal of NaCl comes into contact with water, the partially positively charged ends of water molecules interact with the Cl− ions, while the negatively charged end of the water molecules interacts with the Na+ ions. This ion-dipole interaction between ions and water molecules assists in the breaking of the strong electrostatic forces of attraction within the crystal and ultimately in the solubility of the crystal.

Corrosion

Alloys

Alloys are homogeneous mixtures of a metal with other metals or nonmetals. Alloy formation enhances the desirable properties of the material, such as hardness, tensile strength, and resistance to corrosion.

Examples of a few alloys:
Brass: copper and zinc
Bronze: copper and tin
Solder: lead and tin
Amalgam: mercury and other metal

Corrosion

Gradual deterioration of material usually a metal by the action of moisture, air, or chemicals in the surrounding environment.
Rusting:

4Fe(s)+3O₂(from air)+xH₂O(moisture)→2Fe₂O₃. xH₂O(rust)

Corrosion of copper:

Cu(s)+H₂O(moisture)+CO₂(from air)→CuCO₃.Cu(OH)₂(green)

Corrosion of silver:

Ag(s)+H₂S(from air)→Ag₂S(black)+H₂(g)

Prevention of Corrosion

Prevention:
1. Coating with paints or oil or grease: The application of paint or oil or grease on metal surfaces keep out air and moisture.

2. Alloying: Alloyed metal is more resistant to corrosion. Example: stainless steel.

3. Galvanization: This is a process of coating molten zinc on iron articles. Zinc forms a protective layer and prevents corrosion.

4. Electroplating: It is a method of coating one metal with another by the use of an electric current. This method not only lends protection but also enhances the metallic appearance.
Examples: silver plating, nickel plating.

5. Sacrificial protection: Magnesium is more reactive than iron. When it is coated on articles made of iron or steel, it acts as the cathode, undergoes reaction (sacrifice) instead of iron, and protects the articles.