Exploring the Remarkable Properties of Metals: Conductivity, Malleability, and More

Metals exhibit various distinctive properties, including conductivity, malleability, ductility, luster, and high melting points. They tend to form positive ions and participate in metallic bonding, contributing to their widespread use in construction, electronics, and manufacturing.

The Physical Marvels of Metals: Strength, State, and Shining Characteristics

• Tensile Strength: Metals are hard and have a high tensile strength.
  • Carbon is the only non-metal with very high tensile strength.
• Solid State: Metals are in solid state at room temperature.
  • Mercury is the metal that is liquid below or near room temperature.
  • One non-metal, Bromine, is a liquid at room temperature. 
  • The other non-metals are solids at room temperature, including carbon and sulphur.
• Sonorous: Metals produce a typical ringing sound when hit by something.
• Conductivity: Metals have good conductivity of heat and electricity. 
  • Example: Graphite has good conductivity of heat and electricity.
• Malleable: Metals can be beaten into thin sheets.
• Ductile: Metals can be drawn into thin wires.
• Melting and Boiling Points: Because of conductivity of Metals have 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 a solid state at room temperature.
• Dense: Density of metals is high (except alkali metals). Osmium has the highest density, and lithium has 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 which are lustrous. 

• • Note that carbon is lustrous only in certain forms like diamond and graphite.

• Colour: Metals are silver-grey in colour (except gold and copper). 

Chemical Properties of Metals:

• Exploring the Formation of Metal Oxides through Reaction with Oxygen

• • A 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.
  • Metal + Oxygen→ Metal oxide (basic) . For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide.

                                                   2Cu + O₂ → 2CuO

                                                (Copper)               (Copper (II) oxide)

• • Similarly, aluminium forms aluminium oxide:

                                                        4Al + 3O₂ → 2Al₂O₃

                                                 (Aluminium)           (Aluminium oxide)

• 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. 
    • 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) → 2NaOH(aq)
    • K₂O(s) + H₂O(l) → 2KOH(aq)

• Amphoteric Oxides of Metals:

• • Amphoteric oxides are metal oxides which react with both acids as well as bases to form salt and water.
    • Al₂O₃(s)+6HCl→2AlCl₃+3H₂O(l)
    • Al₂O₃(s)+2NaOH→2NaAlO₂+H₂O(l)

• Reaction of Metals with Water or Steam (Refer Figure 9.3):

• • Aluminium, iron, and zinc are metals that do not react with either cold or hot water. 
    • However, when they come into contact with steam, they produce metal oxide and hydrogen. 
  • Lead, copper, silver, and gold are metals that do not react with water.
  • Calcium starts floating because it comes in contact with the hydrogen bubbles.
  • Metal + Water →Metal Hydroxide or Metal oxide + Hydrogen
    • 2Na+2H₂O (cold)→2NaOH+H₂+heat
    • 2K(s) + 2H₂O(l) → 2KOH + H₂(g) + heat energy
    • Ca+2H₂O (cold)→Ca (OH)₂+H₂
    • 2Al+3H₂O (steam)→Al₂O₃+3H₂
    • 3Fe+4H₂O (steam)→Fe₃O₄+4H₂

• Metal Reactivity with Acids: Hydrogen Evolution and Reactivity Trends  in the Context of Properties of Metals

• • 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 + dilute acid → Salt + Hydrogen Gas
  • Hydrogen gas is not evolved when a metal reacts with nitric acid.
    •  It is because HNO₃ is a strong oxidising agent. 
  • It oxidises the H₂ produced to water and itself gets reduced to any of the nitrogen oxides (N₂O, NO,NO₂).
    • But magnesium (Mg) and manganese (Mn) react with very dilute HNO₃ to evolve H₂ gas.
  • The rate of formation of bubbles is the fastest in the case of magnesium.
    • The reactivity decreases in the order Mg > Al > Zn > Fe.
  • In the case of copper, no bubbles were seen and the temperature also remained unchanged. 
    • This shows that copper does not react with dilute HCl.

• Reaction of Metals with solutions of other Metal Salts: 

• •  Metal Displacement Reaction: 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

• 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

• The Reactivity Series: Understanding the Descending Order of Metal Activities and Protective Oxide Layers

• • Metals: The reactivity series is a list of metals arranged in the order of their decreasing activities.
  • Arrangement of Metals: The reactivity series of metals, also known as the activity series, refers to the arrangement of metals in the descending order of their reactivities. 
  • Vigorous Activity: Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. 
    • Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil.
  • At ordinary temperature, the surfaces of metals such as magnesium, aluminium, zinc, lead, etc., are covered with a thin layer of oxide.
    • The protective oxide layer prevents the metal from further oxidation.
  • Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner.
  • Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide.
  • Silver and gold do not react with oxygen even at high temperatures. exemplifying the noble Properties of metals.

• Electrostatic Bonding: The Formation of Sodium Chloride through Metal and Non-Metal Interaction

• • Sodium atom has one electron in its outermost shell. 
  • If it loses the electron from its M shell then its L shell now becomes the outermost shell and that has a stable octet. 
    • The nucleus of this atom still has 11 protons but the number of electrons has become 10, so there is a net positive charge giving us a sodium cation Na+.
  • Chlorine has seven electrons in its outermost shell and it requires one more  electron to complete its octet.
  • If sodium and chlorine were to react, the electron lost by sodium could be taken up by chlorine. 
  • After gaining an electron, the chlorine atom gets a unit negative charge.
  • Because its nucleus has 17 protons and there are 18 electrons in its K, L and M shells. 
  • This gives us a chloride anion Cl–. So both these elements can have a give-and-take relation between them.
    • Na →Na+  + e–
    • Cl +e– →Cl–
    • Na + Cl →NaCl (Formation of Sodium Chloride)
  • Sodium and chloride ions, being oppositely charged, attract each other and are held by strong electrostatic forces of attraction to exist as sodium chloride (NaCl).

• It should be noted that sodium chloride does not exist as molecules but aggregates of oppositely charged ions. The unique properties of metals contribute to the stability and electrostatic bonding observed in the formation of compounds like sodium chloride.

Properties of Ionic Compounds

Physical Nature of Ionic Compounds: Exploring the Solid State, Hardness, Melting Points and Electrical Conductivity of Electrovalent Compounds

• State of Matter and Hardness: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. 
  • These compounds are generally brittle and break into pieces when pressure is applied.
• Melting and Boiling points: Ionic compounds have high melting and boiling points. 
  • This is because a considerable amount of energy is required to break the strong inter-ionic attraction.
• Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
• Conductivity of Electricity: The conductivity of electricity through a solution involves the movement of charged particles. 
  • A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution.

Understanding these properties Of metals and its Conductivity provides insights into the diverse behaviors of ionic compounds, influencing their applications across various scientific and industrial domains.

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