Metals And Non Metals-IV

Aluminium, iron, and zinc are moderately reactive metals. They do not react with cold water (or react extremely slowly due to protective oxide layers). They also do not react vigorously with hot water. However, when steam (water vapour at high temperature) is passed over these heated metals, they react to form the metal oxide and hydrogen gas. The reactions are: 2Al + 3H₂O(g) → Al₂O₃ + 3H₂; 3Fe + 4H₂O(g) → Fe₃O₄ + 4H₂; Zn + H₂O(g) → ZnO + H₂. This distinguishes them from very reactive metals (which react with cold water) and unreactive metals (which don’t react with steam).
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The reactivity series (also called the activity series) is a list of metals arranged in order of their decreasing chemical reactivity. The most reactive metal (potassium) is placed at the top, and the least reactive metal (gold) is placed at the bottom. This arrangement allows chemists to predict the behavior of metals, such as their reaction with water, acids, oxygen, and their ability to displace other metals from compound solutions. Density, alphabetical order, or atomic mass do not correlate with chemical reactivity.
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The rate of reaction with dilute acids (like HCl) follows the reactivity series. Magnesium (Mg) is highly reactive and produces hydrogen bubbles very rapidly and vigorously. Aluminium (Al) is also highly reactive, but it initially reacts slowly due to its protective oxide layer (Al₂O₃); once this layer is removed, it reacts faster than zinc. Zinc (Zn) reacts moderately, producing a steady stream of bubbles. Iron (Fe) reacts slowly, producing bubbles at a much slower rate. Thus, the correct decreasing order is Mg > Al > Zn > Fe.
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Nitric acid (HNO₃) is a powerful oxidizing agent, especially when concentrated. When a metal reacts with dilute or concentrated nitric acid, the hydrogen that would normally be produced (Metal + Acid → Salt + H₂) is immediately oxidized by the nitric acid to water. Instead of hydrogen, nitric acid itself gets reduced to various products such as nitrogen dioxide (NO₂, brown gas), nitric oxide (NO), or nitrous oxide (N₂O), depending on the concentration of the acid and the reactivity of the metal. This oxidizing property prevents hydrogen gas evolution.
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An exothermic reaction releases heat. The reaction between magnesium and dilute hydrochloric acid is highly exothermic. The temperature of the solution rises significantly and quickly. The equation is: Mg + 2HCl → MgCl₂ + H₂ + Heat. The vigor of the reaction (rate of bubble formation) correlates directly with the amount of heat released per mole of metal. Magnesium, being the most reactive among the options given (above aluminium, zinc, and iron in the reactivity series), produces the most heat in the shortest time, making it the most exothermic reaction.
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Calcium reacts with cold water to form calcium hydroxide and hydrogen gas: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). The reaction is moderately vigorous but not explosive. As hydrogen gas is produced at the surface of the calcium metal, it forms tiny bubbles that adhere to the metal’s surface. These bubbles reduce the density of the calcium piece (by increasing its volume without significantly increasing its mass), causing the metal to become buoyant and float on the water surface. It is not because calcium is light (its density is 1.55 g/cm³, greater than water’s 1.0 g/cm³).
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In solid ionic compounds (like NaCl), the positive and negative ions are held together in a fixed, rigid three-dimensional lattice by strong electrostatic forces. For electrical conductivity, charged particles must be able to move freely. In the solid state, the ions are not mobile; they can only vibrate in their fixed positions. They cannot move from one place to another. Therefore, solid ionic compounds do not conduct electricity. When melted (molten) or dissolved in water, the ions become free to move and conduct electricity.
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Sodium and potassium are extremely reactive metals (alkali metals). At room temperature, they react so quickly with the oxygen and moisture present in the air that the reaction generates enough heat to ignite the metal itself. The heat of reaction first melts the metal (due to its low melting point) and then ignites it. Sodium burns with a golden-yellow flame, and potassium burns with a lilac flame. To prevent this, they must be stored completely immersed in kerosene oil to cut off contact with air and moisture.
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Oxygen is the most abundant element in the Earth’s crust (about 46% by mass). Therefore, many metals occur naturally in the form of their oxides. Metal oxides are stable compounds formed when metals combine with oxygen. Examples include bauxite (Al₂O₃), hematite (Fe₂O₃), and rutile (TiO₂). While other ores like sulphides (e.g., ZnS, PbS) and carbonates (e.g., CaCO₃, ZnCO₃) also exist, oxide ores are extremely common due to the high abundance and reactivity of oxygen.
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Sodium and potassium are extremely reactive metals. When placed in an acid (like dilute HCl or H₂SO₄), the reaction is dangerously explosive. They first react with the water present in the dilute acid (as they would with water), producing a large amount of heat. The acid then reacts even more vigorously. The heat generated is sufficient to ignite the hydrogen gas produced, causing a fire or explosion. For safety reasons, these metals are never used in school laboratory acid reactions. Even with concentrated acids, the reaction is too violent.
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Ionic compounds are generally soluble in polar solvents like water. Water molecules are polar (having a partial positive charge on hydrogen and a partial negative charge on oxygen). These polar water molecules attract the positive and negative ions of the ionic compound. The water molecules surround and hydrate the individual ions, overcoming the strong electrostatic forces holding the crystal lattice together and pulling the ions into solution. Ionic compounds are insoluble or only slightly soluble in non-polar solvents like petrol, oil, and kerosene.
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Sodium reacts spontaneously with oxygen and water vapor present in the air, generating enough heat to catch fire. Storing it under kerosene oil serves as a physical barrier, cutting off contact with both air (oxygen) and moisture. Kerosene is a hydrocarbon that does not react with sodium. This prevents accidental fires and ensures safe storage. Sodium does not rust (rusting is specific to iron); the term for sodium’s degradation is simply “reaction” or “oxidation.”
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The reaction between sodium and cold water is highly exothermic, meaning it releases a large amount of heat energy. The equation is: 2Na + 2H₂O → 2NaOH + H₂ + Heat. This heat is so intense that it melts the sodium (which has a low melting point) and ignites the hydrogen gas produced. The combination of the rapid production of flammable hydrogen gas and the significant exothermic heat is what makes the reaction violent. An endothermic reaction would absorb heat and would not cause such violence.
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Copper is below hydrogen in the reactivity series. This means that copper atoms have a very low tendency to lose electrons (oxidize) compared to hydrogen atoms. For a metal to displace hydrogen from an acid, it must be a stronger reducing agent than hydrogen. Since copper is a weaker reducing agent, it cannot donate electrons to H⁺ ions in the acid. Therefore, no reaction occurs. The principle is: only metals above hydrogen in the series can displace hydrogen from dilute acids, regardless of their physical properties like heaviness, brittleness, or ductility.
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The rate of a chemical reaction depends on the surface area of the reactants. Iron, in its bulk form (like a nail or sheet), does not burn in air; it simply rusts slowly. However, when iron is finely divided into filings, the total surface area exposed to oxygen increases dramatically. When heated, the increased surface area allows oxygen to react with a large number of iron atoms simultaneously, causing the filings to burn vigorously, producing sparks of iron(III) oxide (Fe₂O₃ or Fe₃O₄). Iron is not considered “very reactive” compared to alkali metals.
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In typical Class 10 Activity 7.9 (comparing reactions of metals with water), sodium is observed to react most violently. Sodium reacts with cold water immediately, producing a lot of heat, fizzing, and hydrogen gas that catches fire. Magnesium reacts very slowly with cold water and only vigorously with steam. Aluminium shows almost no reaction with cold water due to its oxide layer. Copper does not react with water at all. Therefore, sodium is consistently the most reactive metal observed in such an activity.
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Sodium (Na) has an atomic number of 11 with an electronic configuration of 2,8,1. It has one valence electron in its outermost shell. To achieve the stable octet configuration of the nearest noble gas (neon, which has 2,8), it is easier for sodium to lose its single valence electron than to gain seven electrons. By losing one electron, it becomes a positively charged sodium ion (Na⁺). The process is: Na → Na⁺ + e⁻. This is a characteristic property of metals.
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The reactivity series ranks metals from most reactive to least reactive. Metals that are very low in the series (below hydrogen) have extremely low chemical reactivity. Lead (Pb) reacts with steam but very slowly with water. Copper (Cu), silver (Ag), and gold (Au) are noble metals that do not react with cold water, hot water, or even steam under normal conditions. They have high reduction potentials and cannot displace hydrogen from water molecules. In contrast, Na, K, Mg, Ca, Al, and Zn show some reaction with water or steam.
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When copper metal is heated strongly in air (or oxygen), it undergoes a combination reaction to form copper(II) oxide. The chemical reaction is: 2Cu(s) + O₂(g) → 2CuO(s). Copper(II) oxide (CuO) is a black, powdery solid. While copper(I) oxide (Cu₂O) is red and may form initially at lower temperatures, continued heating in ample air converts it to the more stable black CuO. Cu(OH)₂ (copper hydroxide) is blue and formed by reacting copper salts with bases, not by heating copper in air.
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Chlorine (Cl) has an atomic number of 17 with an electronic configuration of 2,8,7. It has seven valence electrons and needs one more electron to achieve the stable octet configuration of the nearest noble gas (argon, which has 2,8,8). It is easier for chlorine to gain one electron than to lose seven electrons. By gaining one electron, it becomes a negatively charged chloride ion (Cl⁻). The process is: Cl + e⁻ → Cl⁻. This is a characteristic property of non-metals.
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Sodium chloride (NaCl) is an ionic compound. It does not exist as discrete molecules or individual atoms. Instead, it forms a giant three-dimensional crystal lattice structure consisting of alternating positive sodium ions (Na⁺) and negative chloride ions (Cl⁻). These ions are held together by strong electrostatic forces of attraction (ionic bonds). In the solid state, each Na⁺ is surrounded by six Cl⁻ ions, and each Cl⁻ is surrounded by six Na⁺ ions. This arrangement extends throughout the crystal.
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A displacement reaction occurs when a more reactive metal (which loses electrons more easily) donates electrons to the ions of a less reactive metal. This causes the less reactive metal to come out of its compound (as a solid) while the more reactive metal forms a new compound. For example: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s). Iron (more reactive) displaces copper (less reactive) from copper sulphate solution. This is a fundamental principle of the reactivity series.
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For electrical conductivity, charged particles must be able to move freely. In the solid state, the ions in an ionic compound are held in fixed positions within a rigid crystal lattice and cannot move. However, when the compound is melted (heated to its melting point), the lattice breaks down. The ions (cations and anions) are no longer fixed and become free to move throughout the molten liquid. When an electric potential is applied, these mobile ions migrate towards the oppositely charged electrodes, conducting electricity.
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Moderately reactive metals (such as iron, zinc, copper, and lead) are not found in their free (native) state in nature because they react with atmospheric oxygen, moisture, and other elements over geological time scales. Instead, they occur as stable compounds called minerals. If these minerals contain a high enough percentage of the metal to be economically viable, they are called ores. Common forms of ores for moderately reactive metals include oxides (e.g., Fe₂O₃), sulphides (e.g., ZnS, PbS), and carbonates (e.g., ZnCO₃, CuCO₃).
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The rate of hydrogen bubble formation in dilute hydrochloric acid is a direct indicator of the metal’s reactivity. Among the common metals, magnesium (Mg) is the most reactive towards dilute acids (among the options given). When a piece of magnesium ribbon is dropped into dilute HCl, it reacts immediately, producing a vigorous stream of tiny hydrogen bubbles very rapidly. The reaction is highly exothermic. Zinc reacts steadily, iron reacts slowly, and aluminium (though more reactive than zinc) initially reacts slowly due to its protective oxide layer.
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Ionic compounds are hard because the strong electrostatic forces of attraction (ionic bonds) between oppositely charged ions hold the lattice together firmly. They are brittle because when a force is applied, like a hammer strike, it causes a shift in the layers of ions. This shift brings ions of the same charge (positive-positive or negative-negative) next to each other. The strong repulsion between these like-charged ions causes the crystal lattice to shatter or crack along cleavage planes, rather than deforming.
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Displacement reactions provide direct, unambiguous evidence of the relative reactivity of two metals. If metal A displaces metal B from a solution of B’s compound, it proves that metal A is more reactive than metal B. This is because a more reactive metal can reduce the ion of a less reactive metal (it loses electrons more readily, and the less reactive metal ion gains them). This provides a quantitative comparison that is often clearer than qualitative observations of reactions with water or acids.
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Iron is more reactive than copper (iron is above copper in the reactivity series). When an iron nail is placed in a blue solution of copper sulphate (CuSO₄), a displacement reaction occurs. Iron atoms lose electrons to become Fe²⁺ ions, and Cu²⁺ ions gain those electrons to become copper metal. The equation is: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s). The blue color fades (as Cu²⁺ is removed), and a reddish-brown coating of copper metal appears on the iron nail. Iron cannot displace zinc (more reactive) or sodium (much more reactive).
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In chemistry, a cation is a positively charged ion. In the ionic compound magnesium chloride (MgCl₂), the compound is formed by the transfer of two electrons from a magnesium atom to two chlorine atoms. The magnesium atom loses two electrons to become a magnesium ion with a 2+ charge, written as Mg²⁺. This Mg²⁺ is the cation. The Cl⁻ ions are anions (negatively charged ions). Ca²⁺ is the calcium ion, and Na⁺ is the sodium ion, which are not present in MgCl₂.
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Nitric acid is a strong oxidizing agent and usually does not produce hydrogen gas. However, when the nitric acid is extremely dilute (less than 2% concentration) and the metal is not too reactive, hydrogen can sometimes be evolved. Specifically, metals like magnesium (Mg) and manganese (Mn) can react with very dilute HNO₃ to produce hydrogen gas because the oxidizing power of the acid is significantly reduced at such low concentrations. For most other metals (like Zn, Fe), even dilute HNO₃ usually produces nitrogen oxides. Na and K react too violently, and Cu/Ag do not react.
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Magnesium, aluminium, and zinc are moderately reactive metals. When exposed to air at room temperature, their surfaces react with atmospheric oxygen to form a very thin, adherent, and stable layer of their respective metal oxides (MgO, Al₂O₃, ZnO). This layer is dense, impervious, and strongly bonded to the underlying metal. It acts as a physical barrier, preventing oxygen and moisture from reaching the metal surface. This phenomenon, called passivation, stops further oxidation and protects the metal from corrosion.
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Highly reactive metals (like potassium, sodium, calcium, magnesium, and aluminium) have a very strong tendency to lose electrons and form positive ions. They react readily with oxygen, moisture, and other elements in the environment. Therefore, they are never found in nature in their pure, uncombined (free) state. They always occur as compounds (minerals, ores) such as oxides, chlorides, carbonates, or silicates. Less reactive metals (like gold, silver, platinum) can be found in the free state because they do not react easily.
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Hydrogen is not a metal, but it is included in the reactivity series as a reference point. Its position is based on the ability of metals to displace hydrogen from dilute acids. Metals above hydrogen in the series can displace hydrogen from acids (e.g., Zn + H₂SO₄ → ZnSO₄ + H₂). Metals below hydrogen cannot displace hydrogen from acids. This inclusion allows chemists to predict whether a given metal will react with a dilute acid to produce hydrogen gas, making the series a practical tool for predicting chemical behavior.
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The reactivity series arranges metals from most reactive (potassium, sodium) to least reactive. Gold (Au) sits at the very bottom of the series, below silver, mercury, copper, and platinum. It is the most noble metal, meaning it has the least tendency to lose electrons and form positive ions. Gold does not react with oxygen, water, most acids, or even air at high temperatures. It can only be dissolved using a special mixture called aqua regia (a 3:1 mixture of concentrated HCl and concentrated HNO₃).
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The Earth’s crust (the outermost solid layer of the Earth) is the primary source of most metals and their ores. While non-metals like oxygen, silicon, and carbon are also abundant in the crust, the crust contains the vast majority of economically extractable metal resources. These metals are found combined with other elements in minerals such as oxides, sulphides, carbonates, and silicates. Gases are primarily obtained from the atmosphere, and acids are human-made or produced by natural processes but are not considered a “major source” from the crust.
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The general pattern for metal-water reactions depends on temperature. For most metals (except very reactive ones like Na, K, Ca), the reaction with water (especially steam) produces a metal oxide and hydrogen gas. For example: 2Al + 3H₂O(g) → Al₂O₃ + 3H₂; 3Fe + 4H₂O(g) → Fe₃O₄ + 4H₂. The production of metal hydroxide + hydrogen is specific to very reactive metals reacting with cold water (e.g., 2Na + 2H₂O → 2NaOH + H₂). The question asks “generally,” and for a broad range of metals (e.g., Fe, Zn, Al with steam), the oxide is the common product.
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Silver (Ag) and gold (Au) are noble metals located at the very bottom of the reactivity series. They have extremely high reduction potentials and are chemically very unreactive. Gold does not react with oxygen at any temperature, remaining pure and lustrous. Silver does not form silver oxide (Ag₂O) even at high temperatures because the oxide is unstable and decomposes back to the metal upon heating. In contrast, sodium, potassium, magnesium, zinc, iron, and copper all form oxides when heated in air.
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The general reaction of a metal (above hydrogen in the reactivity series) with a dilute acid (like HCl or H₂SO₄) is a single displacement reaction. The metal displaces hydrogen from the acid. The products are always a salt (formed from the metal and the acid’s anion) and hydrogen gas. The general equation is: Metal + Acid → Salt + Hydrogen. For example: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g). This is a fundamental method for preparing hydrogen gas in laboratories.
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The method used to extract a metal from its ore depends primarily on the metal’s position in the reactivity series. Highly reactive metals (K, Na, Ca, Mg, Al) have very stable compounds and require electrolytic reduction (electrolysis) for extraction. Moderately reactive metals (Zn, Fe, Sn, Pb) can be extracted by reduction with carbon or carbon monoxide (e.g., in a blast furnace). Less reactive metals (Cu, Ag, Au) occur in free state or can be extracted by simple heating or chemical processes. Different reactivities dictate different extraction methods.
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Calcium reacts with cold water to form calcium hydroxide and hydrogen gas. The balanced chemical equation is: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). Calcium hydroxide, also known as slaked lime, is only slightly soluble in water, forming a milky white suspension called limewater. Calcium oxide (CaO, quicklime) is formed by heating calcium carbonate, not by reacting calcium with water. Calcium chloride (CaCl₂) is formed by reacting calcium with hydrochloric acid.
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A mineral is a naturally occurring, inorganic solid with a definite chemical composition and crystal structure. When a mineral contains a sufficiently high percentage of a metal (or another valuable element) that it can be profitably extracted and processed, it is called an ore. Not all minerals are ores; the term “ore” implies economic viability. For example, bauxite is a mineral that is an ore of aluminium, while ordinary clay (also containing aluminium) is not considered an ore because extracting aluminium from it would be too expensive.
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Ionic compounds are formed by the complete transfer of one or more valence electrons from a metal atom (which becomes a positively charged cation) to a non-metal atom (which becomes a negatively charged anion). The resulting oppositely charged ions are then held together by strong electrostatic forces of attraction (ionic bonds). In contrast, covalent compounds (also called molecular compounds) are formed by the sharing of electrons between non-metal atoms, not by transfer.
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Gold (Au) and platinum (Pt) are noble metals located at the very bottom of the reactivity series. They are extremely unreactive and have a very low tendency to lose electrons and form positive ions. They do not readily react with oxygen, moisture, or most other elements present in the environment. Therefore, they are found in nature in their native or free (uncombined) state, often as nuggets or grains in alluvial deposits (placer deposits). Most other metals are found in combined forms (oxides, sulphides, carbonates).
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Noble gases (helium, neon, argon, krypton, xenon, radon) have a complete octet (or duplet for helium) in their outermost electron shell. This electronic configuration is extremely stable. They have no tendency to gain, lose, or share electrons because doing so would disrupt this stable configuration. As a result, noble gases are chemically inert or very unreactive under normal conditions. Their lack of reactivity is due to their stable electron configuration, not density or absence of electrons.
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Ionic compounds are generally solids at room temperature. This is because the strong electrostatic forces of attraction (ionic bonds) between the oppositely charged ions hold them together in a rigid, three-dimensional crystal lattice. A large amount of energy is required to overcome these forces and melt the compound (change it to a liquid) or vaporize it (change it to a gas). Therefore, most ionic compounds have high melting and boiling points and exist as solids under standard conditions. Examples include NaCl, MgO, CaCO₃, and KBr.
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Magnesium reacts very slowly with cold water, producing a few bubbles of hydrogen before stopping due to the formation of insoluble magnesium hydroxide (Mg(OH)₂) on its surface. With hot water, the reaction is still very slow. However, when steam (water vapour at high temperature) is passed over heated magnesium, the reaction proceeds readily. The reaction is: Mg(s) + H₂O(g) → MgO(s) + H₂(g). Therefore, to get a vigorous and complete reaction, magnesium requires steam, not liquid water (hot or cold).
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The reactivity series places potassium (K) at the very top, followed by sodium (Na), then calcium (Ca), magnesium (Mg), aluminium (Al), etc. Potassium is the most reactive metal among all metals. It reacts more violently than sodium with water, producing enough heat to ignite not only the hydrogen gas but also the potassium metal itself, burning with a lilac flame. This higher reactivity of potassium compared to sodium is due to its larger atomic size and lower ionization energy, making it easier to lose its single valence electron.
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Ionic compounds are composed of positive and negative ions held together by strong electrostatic forces of attraction (ionic bonds) in a giant crystal lattice. A large amount of thermal energy (heat) is required to overcome these powerful inter-ionic forces and to break the lattice structure so that the solid can melt into a liquid. The stronger the attraction between the ions (which depends on the charges and sizes of the ions), the higher the melting point. For example, MgO (Mg²⁺ and O²⁻) has a very high melting point (~2850°C).
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The reactivity series is arranged in order of decreasing reactivity from top to bottom. Therefore, metals at the top of the reactivity series (like potassium, sodium, calcium, magnesium) are the most highly reactive metals. They have a very strong tendency to lose electrons and form positive ions. They react violently with water and oxygen, are never found in the free state in nature, and require electrolysis for extraction. As we move down the series, reactivity decreases.
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The reaction between sodium and cold water is extremely violent and highly exothermic: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + Heat. The large amount of heat generated ignites the hydrogen gas produced. The hydrogen burns with a golden-yellow flame (colored by the sodium vapour). This is why sodium is never used in acid reactions and is stored under kerosene to prevent accidental contact with moisture. The hydrogen does not dissolve in water and is certainly produced.Close