Metals and Non Metals Class 10: In the arrangement known as the periodic table, the elements present on our planet are systematically organized according to their ascending atomic numbers.
Our current understanding recognizes a grand total of 118 elements, among which 92 occur naturally, while the remaining 26 are synthesized artificially within laboratory settings.
These elements can be categorized into three distinct groups—Metalloids, Metals, and Non-Metals—based on their unique physical and chemical characteristics.
Metals and Non Metals Class 10 Notes: NCERT Science Chapter 3
Metals and Non Metals Class 10
Physical Properties: Metals and Non Metals Class 10
A quantifiable attribute that signifies a state of a physical system is known as a physical property. The physical properties of a system serve to define its shifts between momentary conditions. The term “observables” is employed to denote these discernible physical attributes.
Physical Properties of Metals
- Possessing significant hardness and high tensile strength – Carbon stands out as the sole non-metal with remarkably high tensile strength.
- Maintaining solidity at room temperature – While one non-metal, bromine, defies this by being a liquid at room temperature, others such as carbon and sulfur remain solid under these conditions.
- Exhibiting sonority – When struck, metals generate a distinctive ringing sound.
Proficient conductors of heat and electricity – Graphite notably excels as a conductor of both heat and electricity. - Demonstrating malleability, enabling the shaping into thin sheets.
Displaying ductility, allowing the drawing into fine wires. - Exhibiting notably high melting and boiling points (with exceptions like Caesium (Cs) and Gallium (Ga)) – An illustrative instance is graphite, a non-metal form of carbon, which boasts a high boiling point and remains solid at room temperature.
- Possessing density (excluding alkali metals), with Osmium exhibiting the highest density and lithium displaying the lowest.
- Exhibiting luster – The capacity to reflect light and be polished is characteristic of metals, exemplified by gold, silver, and copper. Iodine and carbon, although non-metals, also possess luster, albeit specific to certain forms like diamond and graphite for carbon.
- Generally showcasing a silver-grey hue (aside from gold and copper) – The typical coloration of metals tends to be silver or grey in nature.
Non-Metals
Nonmetals are elements that lack the characteristic properties of metals.
Physical Properties of Non-metals
- Exist in states of matter including solids, liquids, and gases at standard room conditions.
- Fragile
- Not capable of being shaped into thin sheets
- Not capable of being drawn into thin wires
- Lack a resonating sound when struck
- Poor 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 caesium have 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
- Hydrogen – Gas
- Nitrogen – Gas
- Oxygen – Gas
- Fluorine – Gas
- Chlorine – Gas
- Bromine – Liquid
- Iodine – Solid
- Carbon – Solid
- Sulphur – Solid
- Phosphorous – Solid
- Silicon – Solid
Chemical Properties: Metals and Non Metals Class 10
Chemical Properties of Metals
Alkali metals (such as Li, Na, K, etc.) display vigorous reactions with water and oxygen or air.
- Mg reacts with heated water.
- Al, Fe, and Zn undergo reactions with steam.
- Cu, Ag, Pt, and Au show no reactivity with water or weak acids.
Reaction of Metals with Oxygen (Burnt in Air)
When metals undergo combustion in the presence of atmospheric oxygen, they give rise to metal oxides. These oxides are a foundational type of substance discovered naturally, capable of altering the hue of red litmus paper to blue. To prevent interactions with oxygen, moisture, and carbon dioxide in the atmosphere, sodium and potassium metals are stored in kerosene oil.
Metal + Oxygen → Metal oxide (basic)
● Sodium (Na) and potassium (K) are preserved within kerosene oil due to their vigorous reactivity with air, leading to combustion.
4K(s) + O2(g) → 2K2O(s) (highly energetic reaction)
● Magnesium (Mg), aluminum (Al), zinc (Zn), and lead (Pb) exhibit a gradual reaction with air, developing a protective layer that thwarts corrosion.
2Mg(s) + O2(g) → 2MgO(s) (Mg combustion produces intense white light)
4Al(s) + 3O2(g) → 2Al2O3(s)
● Silver, platinum, and gold remain inert, devoid of combustion or reactivity with air.
Basic Oxides of Metals
Metallic oxides are structured crystalline solids comprising a metal cation and an oxide anion. Commonly, they engage in reactions with water to generate bases or combine with acids to produce salts. The equation MO + H2O → M(OH)2 (where M represents a group 2 metal) exemplifies this pattern. Consequently, these compounds are frequently referred to as basic oxides.
Certain metallic oxides are soluble in water, forming alkalis. Their aqueous solutions exhibit the ability to convert red litmus paper to blue.
Examples:
Na2O(s) + H2O(l) → 2NaOH(aq)
K2O(s) + H2O(l) → 2KOH(aq)
Amphoteric Oxides of Metals
Amphoteric oxides display the unique ability to react with both acids and bases, resulting in the formation of salts and water.
For instance, compounds like Al2O3, ZnO, PbO, and SnO fall under this category.
Examples of reactions:
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 activity series of metals, often referred to as the reactivity series, entails the organization of metals in a sequence based on their decreasing levels of reactivity.
The subsequent table presents the reactivity of metals, ranked from greatest to least reactivity.
| 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
Undergoing vigorous heating in the presence of an abundant supply of air, it transforms sulphide ores into oxides. This process serves to eliminate volatile impurities as well.
Equation:
2ZnS(s) + 3O2(g) + Heat → 2ZnO(s) + 2SO2(g)
Calcination
Through intense heating within a restricted air supply, it transforms carbonate and hydrated ores into oxides while simultaneously purging volatile impurities.
Equations:
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 Water or Steam
Aluminium, iron, and zinc are metals that exhibit no reaction with water, be it cold or hot. However, when they encounter steam, they generate metal oxide along with hydrogen gas. In contrast, lead, copper, silver, and gold are metals that remain unreactive when in contact with water.
Metal + Water → Metal hydroxide or Metal oxide + Hydrogen
Examples:
2Na + 2H2O (cold) → 2NaOH + H2 + heat
Ca + 2H2O (cold) → Ca(OH)2 + H2
Mg + 2H2O (hot) → Mg(OH)2 + H2
2Al + 3H2O (steam) → Al2O3 + 3H2
Zn + H2O (steam) → ZnO + H2
3Fe + 4H2O (steam) → Fe3O4 + 4H2
Reaction of Metals with Acid
Upon immersion of a metal into an acid, it undergoes a reduction in size due to the consumption within a chemical process. Concurrently, the appearance of gas bubbles is evident, signifying the generation of hydrogen gas as a byproduct. This reactive hydrogen gas can be further demonstrated through its combustible nature, as evidenced by igniting it with a burning splint.
Metal + Dilute Acid → Salt + Hydrogen gas
Examples:
2Na(s) + 2HCl(dilute) → 2NaCl(aq) + H2(g)
2K(s) + H2SO4(dilute) → K2SO4(aq) + H2(g)
Only magnesium (Mg) and manganese (Mn) exhibit reactivity with highly diluted nitric acid, resulting in the liberation of hydrogen gas.
Examples:
Mg(s) + 2HNO3(dilute) → Mg(NO3)2(aq) + H2(g)
Mn(s) + 2HNO3(dilute) → Mn(NO3)2(aq) + H2(g)
Displacement Reaction
A higher-reactivity element replaces a lower-reactivity element within its compound or solution.
How Do Metals React with the Solution of Other Metal Salts
A metal that possesses greater reactivity has the ability to replace a metal with lesser reactivity within a solution of its salt, leading to what is termed a displacement reaction. This kind of reaction is commonly referred to as a metal displacement reaction. The reactivity of frequently employed metals has been organized in a descending order, constituting the reactivity or activity series.
The reaction takes the form:
Metal A + Salt of metal B → Salt of metal A + Metal B
For instance:
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
This process finds application in thermite welding, where it contributes to the displacement of iron from its oxide using aluminum.
Moreover, it is a pivotal aspect of steel production, where carbon displaces iron from its oxide.
Furthermore, this principle finds predominant use in various metal extraction processes.
Reaction of Metals with Bases
Bases are characterized by a bitter flavor and a smooth, slick sensation. When a base is dissolved in water, it is termed an alkali. Upon interacting with acids, these substances yield salts through chemical reactions. Bases are recognized for their ability to shift the color of red litmus paper to blue.
The reaction follows the pattern:
Base + Metal → Salt + Hydrogen
Examples:
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
Metal Extraction
Production of Steel
Thermite Reaction: Al(s) + Fe2O3(s) → Al2O3 + Fe(molten)
The thermite reaction finds application in the welding of railway tracks, repairing cracked machine parts, and more.
Occurrence of Metals
The majority of elements, particularly metals, are naturally found in a bonded state alongside other elements. These combinations of metal compounds are categorized as minerals. However, only a select few among these minerals serve as practical reservoirs of the respective metal. These specific utilitarian reserves are referred to as ores.
Notably, gold (Au) and platinum (Pt) are examples of elements that exist in their native or free form.
Extraction of Metals
The activity of retrieving metal ores from deep subterranean deposits is referred to as Mining. Within the Earth’s crust, metal ores are present in varying quantities. The extraction of metals from these ores is the pivotal step that enables us to harness the minerals concealed beneath the surface. It’s important to note that ores contrast significantly from the refined metals visible in structures like buildings and bridges. Ores comprise the sought-after metal compounds along with impurities and earthly materials collectively termed Gangue.
Enrichment of Ores
Purification involves the elimination of impurities or gangue from ore by employing diverse physical and chemical methods. The selection of a specific technique for a given ore hinges on the distinctions between the ore and the gangue in their properties.
In the realm of chemistry, gangue signifies an unwanted element or impurity that encases minerals within an ore deposit, which can be anything from sand and rock to other materials. This mineral presence is a frequent occurrence in mining activities.
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)
2HgO(s)+heat→2Hg(l)+O2(g)
2. Cu2S(Copper pyrite)+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 signifies a process in which ore is subjected to heat in the absence of air or with limited air supply. Roasting, on the other hand, involves heating ore in the presence of air or oxygen, but below its melting point. This process is utilized to enhance or modify the properties of ores.
Smelting is a procedure that entails heating the previously roasted or calcined ore (metal oxide) to an elevated temperature alongside a suitable reducing agent. This results in the production of the raw metal in its molten form.
Example:
Fe2O3 + 3C (coke) → 2Fe + 3CO2
An aluminothermic reaction, known as the Goldschmidt reaction, is a profoundly exothermic reaction involving the heating of metal oxides, typically those of Fe and Cr, in the presence of aluminum at high temperatures.
Examples:
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 undergoes electrolysis within a specialized apparatus.
At the cathode (reduction):
Na+(molten) + e− → Na(s)
Metal is precipitated.
At the anode (oxidation):
2Cl−(molten) → Cl2(g) + 2e–
Chlorine gas is liberated.
2. Hall’s Process: An amalgamation of molten alumina and a fluoride solvent, typically cryolite (Na3AlF6), is subjected to electrolysis.
At the cathode (reduction):
2Al3+ + 6e– → 2Al(s)
Metal is deposited.
At the anode (oxidation):
6O2– → 3O2(g) + 12e–
Oxygen gas is released.
Metals positioned atop the reactivity series exhibit pronounced reactivity. These metals cannot be extracted from their compounds through heating with carbon, as their affinity for oxygen surpasses that of carbon. Thus, the electrolytic reduction method is employed for the extraction of such metals.
Refining of Metals: Metals and Non Metals Class 10
Metal refining involves the elimination of impurities or gangue from raw metal. This final phase in metallurgy relies on distinguishing the properties of the metal from those of the gangue.
Electrolytic Refining
Copper, zinc, nickel, silver, tin, gold, and similar metals undergo electrolytic refinement.
Anode: Contains impure or crude metal
Cathode: Comprises a slender strip of pure metal
Electrolyte: Consists of an aqueous solution of a metal salt
During anode oxidation: Metal ions are released into the solution
At cathode reduction: An equivalent quantity of metal from the solution gets deposited
Impurities gather at the base 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 Non-Metals React?
Metals shed valence electrons, leading to the creation of cations.
Non-metals acquire these electrons in their valence shell, resulting in the formation of anions.
The cation and anion are drawn together by a potent electrostatic force, culminating in the establishment of an ionic bond.
For instance, in calcium chloride, the ionic bond emerges from the attraction between oppositely charged calcium and chloride ions.
The calcium atom loses two electrons, achieving the electronic arrangement akin to the nearest noble gas (argon), resulting in a net charge of +2.
Ionic Compounds
Ionic compounds are balanced compounds comprising positively charged cations and negatively charged anions. Binary ionic compounds, those containing just two distinct elements, are named by listing the cation’s name followed by the anion’s name.
The compound’s cohesion arises from the electrostatic attractions between ions of opposing charges.
Illustrations include: MgCl2, CaO, MgO, NaCl, and more.
Properties of Ionic Compound
Ionic compounds:
Typically exist as crystalline solids composed of ions.
Exhibit elevated melting and boiling points.
Display electrical conductivity in aqueous solutions and when liquefied.
Generally dissolve readily in water and polar solvents.
Electric Conduction of Ionic Compounds
Ionic compounds exhibit electrical conductivity when they are in a molten or aqueous state, during which the ions are liberated and function as charge carriers. In their solid form, ions are firmly bound by electrostatic attraction and lack mobility, resulting in the inability to conduct electricity.
CBSE Class 10 Science notes Chapter 3 – 5
CBSE Class 10 Science notes Chapter 3 – 4
As an illustration, consider ionic compounds like NaCl, which do not conduct electricity while in a solid state. However, when these compounds are dissolved in water or in a molten condition, they become capable of conducting electricity.
Melting and Boiling Points of Ionic Compounds
Ionic compounds possess robust electrostatic forces that demand a substantial energy input to disrupt. Consequently, the melting and boiling points of an ionic compound are typically elevated.
Solubility of Ionic Compounds
The majority of ionic compounds tend to dissolve in water, a phenomenon attributed to the dispersion of ions within the water medium. This outcome is a result of water’s polar characteristic.
For instance, consider NaCl, a three-dimensional salt crystal comprised of Na+ and Cl− ions held together by electrostatic forces of attraction. When a NaCl crystal makes contact with water, the partially positively charged ends of water molecules interact with the Cl− ions, while the negatively charged ends of water molecules interact with the Na+ ions. This ion-dipole interaction between ions and water molecules contributes to the disruption of the robust electrostatic forces binding the crystal. As a result, the crystal becomes soluble in water.
Corrosion: Metals and Non Metals Class 10
Alloys
Alloys represent uniform blends of a metal combined with either other metals or nonmetals. The creation of alloys serves to amplify advantageous material properties like hardness, tensile strength, and resistance to corrosion.
Here are a few instances of alloys:
– Brass: amalgamation of copper and zinc
– Bronze: fusion of copper and tin
– Solder: combination of lead and tin
– Amalgam: mixture involving mercury and other metals
Corrosion
Progressive degradation of a substance, often a metal, due to the influence of moisture, air, or chemicals in its ambient surroundings.
Rusting:
4Fe(s) + 3O2(from air) + xH2O(moisture) → 2Fe2O3. xH2O(rust)
Copper Corrosion:
Cu(s) + H2O(moisture) + CO2(from air) → CuCO3.Cu(OH)2(green)
Silver Corrosion:
Ag(s) + H2S(from air) → Ag2S(black) + H2(gas)
Prevention of Corrosion
Preventions
1. Application of Coatings: Metal surfaces can be safeguarded from corrosion by applying protective coatings like paints, oil, or grease. These coatings create a barrier that excludes air and moisture.
2. Alloy Formation: Incorporating metals into alloys enhances their corrosion resistance. For instance, stainless steel is a notable example of an alloyed metal.
3. Galvanization: This involves the deposition of molten zinc onto iron articles. Zinc forms a protective layer that acts as a barrier against corrosion.
4. Electroplating: Through the application of an electric current, one metal can be coated onto another, providing not only protection but also an improved metallic appearance. Illustrations include silver plating and nickel plating.
5. Sacrificial Protection: Magnesium, being more reactive than iron, can be used as a sacrificial layer on iron or steel articles. Acting as a cathode, magnesium undergoes a reaction, sacrificing itself instead of the underlying iron or steel, thus shielding the articles from corrosion.
Read Also
- Chemical Reactions and Equations
- Acids, Bases and Salts
- Life Processes
- Control and Coordination
- How Do Organisms Reproduce?
- Heredity and Evolution
- Electricity
- Magnetic Effects of Electric Current
- Our Environment
- Human Eye and the Colourful World Notes Chapter 10 Science
Frequently Asked Questions on Metals and Non Metals Class 10
1. What are metals and non-metals?
Metals are elements that typically exhibit properties like malleability, ductility, and good conductivity of heat and electricity. Non-metals, on the other hand, often lack these properties and may be brittle and poor conductors.
2. How are metals and non-metals classified on the periodic table?
Metals are generally found on the left side and in the middle of the periodic table, while non-metals are primarily located on the right side.
3. What are metalloids?
Metalloids are elements that have properties intermediate between metals and non-metals. They show a combination of characteristics from both categories.
4. What is the reactivity series of metals?
The reactivity series is a ranking of metals based on their reactivity with water, acids, and other substances. It helps us understand their behavior in various chemical reactions.
5. How do metals react with water and acids?
Many metals react with water to form metal hydroxides and hydrogen gas. They can also react with acids to produce salts and hydrogen gas.
6. What are ionic compounds?
Ionic compounds are compounds generated through the transfer of electrons from a metal to a non-metal, culminating in the creation of positively charged cations and negatively charged anions.
7. What is the difference between metals and non-metals in terms of electron gain and loss?
Metals tend to lose electrons and form positively charged ions (cations), while non-metals tend to gain electrons and form negatively charged ions (anions).
8. How are alloys different from pure metals?
Alloys are mixtures of two or more metals, or a metal and a non-metal, that are combined to enhance specific properties. They often possess improved strength, durability, and resistance to corrosion compared to pure metals.
9. What is the process of corrosion?
Corrosion is the gradual degradation of metals due to the influence of factors like moisture, air, and chemicals in the environment. It often results in the formation of oxides, hydroxides, or other compounds on the metal’s surface.
10. How can we prevent corrosion?
Corrosion prevention can be achieved through techniques such as applying protective coatings like paints or oils, creating alloys, galvanizing, electroplating, and employing sacrificial protection, which involves utilizing a more reactive metal to safeguard a less reactive metal.