Comprehensive theory, key formulas, diagrams, and memory aids for Metals and Non-Metals.
Elements can be broadly classified into two categories based on their physical and chemical properties: metals and non-metals. Look around you—the wiring in your home, the vessels in your kitchen, the oxygen you breathe, and the carbon in your body—all these are made of elements belonging to these two categories. Understanding the distinction between them is fundamental to chemistry and material science.
Metals generally share a set of physical characteristics that make them useful for building structures, conducting electricity, and making tools. * Metallic Luster: In their pure state, metals have a shining surface (e.g., gold, silver, copper). * Hardness: Most metals are hard, but the hardness varies. (Exception: Sodium and potassium are so soft they can be cut with a knife). * Malleability: Metals can be beaten into thin sheets. Gold and silver are the most malleable metals. This is why we have aluminum foil for wrapping food. * Ductility: The ability of metals to be drawn into thin wires. Gold is the most ductile metal (a wire of about 2 km length can be drawn from just 1 gram of gold). * Conductivity: Metals are excellent conductors of heat and electricity. Silver and copper are the best conductors; lead and mercury are comparatively poor conductors. * Sonorous: Metals produce a ringing sound when struck on a hard surface.
Non-metals generally exhibit properties opposite to those of metals. They can be solid, liquid, or gas at room temperature. * State: Mostly solids (like carbon, sulfur) or gases (like oxygen, hydrogen). Bromine is the only non-metal that is liquid at room temperature. * Luster: Non-metals do not have luster; they are dull. (Exception: Iodine crystals are shiny). * Hardness and Brittleness: Solid non-metals are generally soft and brittle (they break when beaten). (Exception: Diamond, an allotrope of carbon, is the hardest naturally occurring substance). * Conductivity: They are poor conductors of heat and electricity. (Exception: Graphite, another allotrope of carbon, is an excellent conductor of electricity).
Metals are electropositive elements; they tend to lose electrons to form positive ions (cations). Their chemical behavior depends heavily on how easily they lose these electrons.
Almost all metals combine with oxygen to form metal oxides. * Metal + Oxygen → Metal Oxide * Example: 2Mg + O₂ → 2MgO (Magnesium burns with a white flame to form magnesium oxide). * Basic Nature: Most metal oxides are basic in nature and turn red litmus blue. Some, like aluminum oxide (Al₂O₃) and zinc oxide (ZnO), show both acidic and basic behavior and are called amphoteric oxides.
Different metals react with water at different rates depending on their reactivity. * Highly Reactive Metals: Potassium and sodium react violently with cold water, evolving hydrogen gas which immediately catches fire due to the heat generated. * 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + Heat * Moderately Reactive Metals: Magnesium reacts only with hot water. Aluminum, iron, and zinc do not react with cold or hot water, but they react with steam. * Least Reactive Metals: Metals like lead, copper, silver, and gold do not react with water at all.
Metals react with dilute acids to displace hydrogen gas and form metal salts. * Metal + Dilute Acid → Salt + Hydrogen Gas * Example: Mg + 2HCl → MgCl₂ + H₂ * Note: Hydrogen gas is not evolved when a metal reacts with dilute nitric acid (HNO₃) because it is a strong oxidizing agent and oxidizes the H₂ produced to water. (Exception: Mg and Mn react with very dilute HNO₃ to evolve H₂).
Not all metals are equally reactive. The reactivity series is a list of metals arranged in the order of their decreasing activities. * Top (Most Reactive): Potassium (K), Sodium (Na), Calcium (Ca), Magnesium (Mg), Aluminum (Al) * Middle (Moderately Reactive): Zinc (Zn), Iron (Fe), Lead (Pb) * [Hydrogen is often included as a reference point] * Bottom (Least Reactive): Copper (Cu), Mercury (Hg), Silver (Ag), Gold (Au)
A more reactive metal can displace a less reactive metal from its salt solution (Displacement Reaction). For instance, iron can displace copper from copper sulfate, but copper cannot displace iron from iron sulfate.
How do metals and non-metals react with each other? They do so by transferring electrons to achieve a stable, full outer electron shell (noble gas configuration). * Metals lose electrons from their valence shell to form positive ions (cations). * Non-metals gain electrons in their valence shell to form negative ions (anions).
When a metal reacts with a non-metal, the transfer of electrons forms ions. The strong electrostatic force of attraction between the oppositely charged cations and anions holds them together. These compounds are called ionic compounds or electrovalent compounds. * Example: Formation of Sodium Chloride (NaCl) * Sodium (Na) has 1 valence electron. It loses it to form Na⁺. * Chlorine (Cl) has 7 valence electrons. It gains 1 electron to form Cl⁻. * The oppositely charged Na⁺ and Cl⁻ attract each other to form the solid crystal NaCl.
Properties of Ionic Compounds: 1. Solid and relatively hard. 2. High melting and boiling points (due to strong inter-ionic attraction). 3. Generally soluble in water but insoluble in organic solvents (like petrol, kerosene). 4. Conduct electricity in the molten state or when dissolved in water (because ions are free to move), but do not conduct electricity in the solid state.
Most metals are found in the Earth's crust as compounds (oxides, sulfides, carbonates) because they are reactive. These naturally occurring compounds are called minerals. Minerals from which a metal can be extracted profitably are called ores.
The extraction of metals involves several steps: 1. Enrichment of Ores: Removing impurities (gangue) like soil, sand, etc., from the ore. 2. Extraction of Metals: * Metals low in activity series (e.g., Hg, Cu): Extracted by simply heating their sulfide ores in the presence of air (Roasting). * Metals in the middle of activity series (e.g., Zn, Fe): Usually present as sulfides or carbonates. They are first converted to metal oxides by roasting (heating in excess air) or calcination (heating in limited air). Then, the oxide is reduced to metal using a reducing agent like carbon (coke). * Metals high in activity series (e.g., Na, Mg, Al): These cannot be reduced by carbon because they have a higher affinity for oxygen. They are obtained by electrolytic reduction (passing electricity through their molten salts). 3. Refining: Purifying the extracted impure metal, often using electrolytic refining.
As discussed in previous chapters, metals deteriorate when exposed to moisture and air. * Silver items turn black due to the formation of silver sulfide. * Copper forms a green coating of basic copper carbonate. * Iron forms flaky brown rust. * Prevention: Galvanization (coating iron with a thin layer of zinc), painting, or oiling.
Alloying is an excellent method of improving the properties of a metal. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. * Steel: Iron mixed with a small amount of carbon (makes it hard and strong). * Stainless Steel: Iron mixed with nickel and chromium (prevents rusting). * Brass: Alloy of Copper and Zinc. * Bronze: Alloy of Copper and Tin. * Amalgam: Any alloy containing mercury as one of the components. * Note: The electrical conductivity and melting point of an alloy are generally lower than that of the pure metals.
The study of metals and non-metals is a study of contrasts. Metals, defined by their luster, malleability, and electron-losing nature, form the physical framework of modern civilization. Non-metals, diverse in state and electron-gaining in nature, are essential for life (oxygen, carbon) and form the basis of organic chemistry. Their interaction results in stable ionic compounds. Understanding the reactivity series allows us to predict chemical behavior and dictates how we extract these essential elements from the Earth through metallurgy.