Metals and Non-metals

Substance M is metal while substance N is non-metal

Metals are malleable (flatten into sheets without breaking when hammered) and sonorous (produce a deep ringing sound when struck). Non-metals are typically brittle and shatter into powder under mechanical stress.

Lustrous: Copper | Good Conductor: Graphite and silver | Malleable: Iron | Bad Conductor: Rubber

Metals like copper have metallic luster and iron is malleable. Silver is the best electrical conductor, and graphite is an allotrope of carbon that uniquely conducts electricity due to delocalised pi electrons. Rubber is an electrical insulator.

$$2\text{Fe}_2\text{O}_3(s)$$ (Ferric oxide / Iron(III) oxide)

Four atoms of iron react with three diatomic molecules of oxygen gas to yield two formula units of iron(III) oxide ($2\text{Fe}_2\text{O}_3$), balancing all 4 iron atoms and 6 oxygen atoms.

$$X - \text{P}_4\text{O}_{10}(s);\quad Y - 2\text{MgO}(s)$$

Tetraphosphorus burns in excess oxygen to synthesize acidic phosphorus pentoxide ($\text{P}_4\text{O}_{10}$). Magnesium metal combusts with a brilliant white flame to yield basic magnesium oxide ($2\text{MgO}$).

$$\text{Mg(NO}_3)_2 + \text{H}_2\text{O}$$

Magnesium oxide ($\text{MgO}$) is a basic metal oxide. It reacts with nitric acid ($\text{HNO}_3$) to undergo neutralisation, forming magnesium nitrate salt and water.

$$2\text{HCl} + \text{Ba(OH)}_2 \longrightarrow \text{BaCl}_2 + 2\text{H}_2\text{O}$$

Barium hydroxide neutralises two moles of hydrochloric acid to precipitate/form barium chloride ($\text{BaCl}_2$) and two moles of water.

$$\text{Ca(OH)}_2$$ (Calcium hydroxide / Slaked lime)

Calcium oxide combines exothermically with water to form calcium hydroxide: $$\text{CaO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$ No gas bubbles are evolved.

Test the basicity using a red litmus paper.

Dissolving metallic ash ($\text{MgO}$) in water forms magnesium hydroxide ($\text{Mg(OH)}_2$), a basic alkali solution that turns moist red litmus paper blue.

It catches fire and forms hydroxide (and hydrogen gas).

The reaction of sodium with water is violently exothermic: $$\text{2Na}(s) + 2\text{H}_2\text{O}(l) \longrightarrow 2\text{NaOH}(aq) + \text{H}_2(g) + \text{Heat}$$ The intense heat released instantly ignites the evolved hydrogen gas with a yellow flame.

Beaker 1: $\text{KOH}$ and $\text{H}_2$ gas; Beaker 2: No reaction

Potassium is an alkali metal high in the reactivity series and reacts violently with cold water ($2\text{K} + 2\text{H}_2\text{O} \rightarrow 2\text{KOH} + \text{H}_2$). Silver is near the bottom of the reactivity series (unreactive noble metal) and does not react with water even at elevated temperatures.

$$\text{Mg(NO}_3)_2 \text{ (Magnesium nitrate) and } \text{H}_2 \text{ (Hydrogen gas)}$$

Although concentrated nitric acid is a strong oxidising agent that oxidises $\text{H}_2$ to water, magnesium ($\text{Mg}$) and manganese ($\text{Mn}$) are unique metals that react with very dilute nitric acid ($1\%$) to evolve hydrogen gas: $$\text{Mg} + 2\text{HNO}_3(\text{very dil.}) \longrightarrow \text{Mg(NO}_3)_2 + \text{H}_2 \uparrow$$

Nitric acid is a strong oxidising agent that oxidises hydrogen to water and gets reduced to $\text{NO}_2$.

Nitric acid oxidises any nascent hydrogen produced during the acid-metal interaction to water ($\text{H}_2\text{O}$) while itself being reduced to oxides of nitrogen ($\text{NO}_2, \text{NO}$).

Lead is more reactive than copper

In the reactivity series, lead ($\text{Pb}$) lies above copper ($\text{Cu}$). Therefore, metallic lead displaces copper ions from aqueous copper chloride solution to form soluble lead chloride and precipitate reddish-brown metallic copper.

$$\text{Silver} < \text{Copper} < \text{Zinc}$$

Zinc displaces copper from $\text{CuSO}_4$, proving $\text{Zinc} > \text{Copper}$. Silver cannot displace copper, proving $\text{Copper} > \text{Silver}$. Hence, increasing order is: $\text{Silver} < \text{Copper} < \text{Zinc}$.

Strong electrostatic forces of attraction between oppositely charged ions require large thermal energy to overcome.

Ionic lattices are held together by powerful inter-ionic coulombic forces throughout a three-dimensional crystal lattice, imparting exceptionally high melting and boiling points.

Dissolved salt dissociates into free mobile ions ($\text{Na}^+$ and $\text{Cl}^-$) that carry electric charge through the solution.

In solid state, ions are locked in rigid lattice positions and cannot move. When dissolved in water, the crystal lattice breaks apart, releasing mobile hydrated ions that migrate towards oppositely charged electrodes to conduct current.

Roasting cinnabar in air to mercuric oxide, then heating to reduce to mercury metal.

Mercury is low in the reactivity series. Heating cinnabar ($2\text{HgS} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{HgO} + 2\text{SO}_2$) yields mercuric oxide, which upon continued heating readily decomposes thermally without any chemical reducing agent: $2\text{HgO} \xrightarrow{\Delta} 2\text{Hg} + \text{O}_2$.

Roasting metal sulphide to oxide, then reducing with carbon coke.

It is easier to obtain a metal from its oxide than from its sulphide or carbonate. Therefore, zinc blende is first roasted in excess air to form zinc oxide ($2\text{ZnO}$), followed by chemical reduction with carbon (coke): $$\text{ZnO}(s) + \text{C}(s) \xrightarrow{\Delta} \text{Zn}(s) + \text{CO}(g)$$

Copper from the impure anode dissolves into the electrolyte, and an equivalent amount of pure copper deposits onto the cathode.

• Anode (impure copper): $\text{Cu} \rightarrow \text{Cu}^{2+} + 2e^-$ (dissolution)
• Cathode (pure copper strip): $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$ (electrodeposition)
• Anode mud: Insoluble precious metals ($\text{Ag}, \text{Au}, \text{Pt}$) settle below the anode.

Galvanisation (coating the iron object with a thin protective layer of zinc).

Zinc is more electropositive than iron and acts as a sacrificial anode. Even if the zinc coating is scratched, zinc corrodes preferentially, shielding the underlying iron from oxidation.

Nail 4 (farthest from the heat source)

Metals conduct heat via lattice vibration and delocalised free electrons from the hot end to the cold end. The thermal energy reaches Nail 1 first (melting its wax first) and Nail 4 last because it is located farthest from the flame.

Thickness / quantity of wax coatings and distance between nails.

To reliably compare the rate of thermal conduction along the rod, all other parameters affecting melting time—specifically the quantity/thickness of wax holding each nail and the spacing between nails—must be controlled variables.

$$\text{ZnSO}_4\text{ (Zinc sulphate)}$$

Zinc is more reactive than copper and displaces copper from copper sulphate:
$$\text{Zn}(s) + \text{CuSO}_4(aq, \text{blue}) \longrightarrow \text{ZnSO}_4(aq, \text{colourless}) + \text{Cu}(s, \text{reddish-brown})$$

Use pieces of small zinc flakes / granulated zinc.

Finely divided zinc flakes offer a vastly greater effective surface area per unit mass than a solid plate. This increases the frequency of collisions between $\text{Cu}^{2+}$ ions and zinc atoms, dramatically speeding up the displacement rate.

Copper is less reactive than zinc.

A displacement reaction only occurs when a more reactive metal is placed in a salt solution of a less reactive metal. Because copper lies below zinc in the electrochemical reactivity series ($\text{Zn} > \text{Cu}$), copper cannot supply electrons to reduce $\text{Zn}^{2+}$ ions, resulting in no reaction.

Extremely low chemical reactivity (noble metals).

Gold is placed at the very bottom of the reactivity series. Because of its very high standard electrode potential and inert chemical character, it does not combine with oxygen, carbon dioxide, moisture, or atmospheric sulphur compounds under normal conditions, and is found in the uncombined native state.

$$3\text{H}_2\text{ (Hydrogen gas)}$$

Metals like aluminium, zinc, and iron do not react with cold or hot liquid water, but react with steam to form the corresponding metal oxide and liberate hydrogen gas:
$$2\text{Al}(s) + 3\text{H}_2\text{O}(g) \longrightarrow \text{Al}_2\text{O}_3(s) + 3\text{H}_2(g)$$

1: No | 2: Yes | 3: Yes

• Statement 1 (No): Ionic compounds are typically highly soluble in polar solvents like water (though insoluble in non-polar solvents like petrol/kerosene).
• Statement 2 (Yes): They have high melting points due to strong inter-ionic electrostatic forces.
• Statement 3 (Yes): They are formed by complete transfer of valence electrons from electropositive metals to electronegative non-metals.

Aluminium ($\text{Al}$)

Aluminium has a much higher chemical affinity for oxygen than carbon does. Carbon cannot overcome the intense binding energy of aluminium oxide (alumina). Hence, aluminium must be extracted by electrolytic reduction (Hall-Héroult process) of molten alumina dissolved in cryolite.

One atom (Formula: $\text{MgS}$)

Magnesium loses its 2 valence electrons to achieve a stable octet: $\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^-$. Sulphur requires exactly 2 electrons to complete its valence shell: $\text{S} + 2e^- \rightarrow \text{S}^{2-}$. Thus, one magnesium atom transfers 2 electrons directly to one sulphur atom to form magnesium sulphide ($\text{Mg}^{2+}\text{S}^{2-} = \text{MgS}$).

Neon ($\text{Ne}$, atomic number 10) has a complete octet configuration ($2, 8$). It exists as an inert, unreactive monatomic gas in Period 2.

Nitrogen forms oxides like $\text{NO}_2$ which dissolve in rainwater to form nitric acid ($\text{HNO}_3$), a primary constituent of acid rain. Carbon forms $\text{CO}_2$ which yields mildly acidic carbonic acid ($\text{H}_2\text{CO}_3$).

Nitrogen is a primary macronutrient for plant growth. It forms nitrates ($\text{NO}_3^-$) and ammonium salts ($\text{NH}_4^+$) in chemical fertilizers (e.g., urea, ammonium nitrate) to enhance soil fertility.

Lithium ($\text{Li}$, atomic number 3) is an alkali metal with electronic configuration $2, 1$. Because it oxidizes rapidly in air and reacts exothermically with moisture, it is stored under mineral oil or paraffin wax.

The atomic number equals the number of protons in the nucleus. Atomic number 4 corresponds to Beryllium ($\text{Be}$).

Fluorine has atomic number $Z=9$ (9 protons, 10 neutrons) giving a standard relative atomic mass of $18.998 \approx 19\text{ u}$.

• 5 valence electrons: Nitrogen (atomic number 7, configuration $2, 5$).
• Least metallic / most electronegative: Fluorine ($\text{F}$) has the highest electronegativity and non-metallic character across Period 2 (or Neon as an inert non-metal).

1. High melting and boiling points: Oppositely charged ions are packed in a 3D giant ionic lattice held by powerful coulombic attractions that require large thermal energy to break.
2. Electrical conductivity in molten/aqueous state: Ions are free to move and carry electric charge in aqueous solution or molten state (non-conductive in solid form).
3. Hard and brittle: Crystal lattices resist shearing; when external stress shifts ionic planes, like charges align and strongly repel, shattering the crystal.

Substances A, B, and D are covalent compounds.

• Reason: Covalent molecular compounds have weak intermolecular van der Waals forces, resulting in low melting points (all below $300\text{ K}$). Furthermore, they consist of neutral molecules with no free ions or mobile electrons, making them non-conductors of electricity in both solid and liquid states.
• Conversely, substances C and E have high melting points and conduct electricity in liquid state, identifying them as ionic compounds.

Step-by-step Stoichiometric Calculation:
1. Molar mass of $\text{Mg}$: $24\text{ g/mol}$
2. Moles of $\text{Mg}$ reacting: $n = \frac{12\text{ g}}{24\text{ g/mol}} = 0.5\text{ mol}$
3. Stoichiometric ratio: According to $2\text{Mg} \rightarrow 2\text{MgO}$, $1\text{ mol}$ of $\text{Mg}$ yields $1\text{ mol}$ of $\text{MgO}$. Thus, $0.5\text{ mol}$ of $\text{Mg}$ yields $0.5\text{ mol}$ of $\text{MgO}$.
4. Molar mass of $\text{MgO}$: $24 + 16 = 40\text{ g/mol}$
5. Mass of $\text{MgO}$ formed: $m = 0.5\text{ mol} \times 40\text{ g/mol} = \mathbf{20\text{ g}}$.

Most reactive: R  |  Least reactive: Q

• From reaction 1: $\text{P}$ displaces $\text{Q}$ from $\text{QX} \implies \mathbf{\text{P} > \text{Q}}$.
• From reaction 2: $\text{R}$ displaces $\text{P}$ from $\text{PY} \implies \mathbf{\text{R} > \text{P}}$.
• Combining both relations: $\mathbf{\text{R} > \text{P} > \text{Q}}$. Therefore, $\text{R}$ is the most reactive and $\text{Q}$ is the least reactive.

Containers Q and R will show a distinct colour change:

• Container P (No change): Zinc cannot displace itself from $\text{ZnSO}_4$.
• Container Q (Pale green to colourless): Zinc is more reactive than iron and displaces iron, forming colourless $\text{ZnSO}_4$ and depositing grey/black iron:
$$\text{Zn}(s) + \text{FeSO}_4(aq, \text{pale green}) \longrightarrow \text{ZnSO}_4(aq, \text{colourless}) + \text{Fe}(s)$$
• Container R (Deep blue to colourless): Zinc displaces copper, discharging the blue color as $\text{Cu}^{2+}$ is reduced to reddish-brown copper deposit:
$$\text{Zn}(s) + \text{CuSO}_4(aq, \text{blue}) \longrightarrow \text{ZnSO}_4(aq, \text{colourless}) + \text{Cu}(s)$$

Observation & Reasoning:
1. Because Metal B displaces iron from $\text{FeSO}_4$, Metal B is more reactive than iron ($\text{B} > \text{Fe}$).
2. In the reactivity series, iron is more reactive than copper ($\text{Fe} > \text{Cu}$).
3. By transitive property, Metal B is significantly more reactive than copper ($\text{B} > \text{Fe} > \text{Cu}$).
4. Therefore, Metal B will readily displace copper from $\text{CuSO}_4$ solution: the blue colour will fade/decolourise and a reddish-brown deposit of copper metal will form on metal B.

$$\text{D} < \text{C} < \text{A} < \text{B}$$

• $\text{D}$ cannot displace any metal $\implies$ least reactive.
• $\text{C}$ displaces only silver (a very unreactive metal).
• $\text{A}$ displaces copper (more reactive than silver).
• $\text{B}$ displaces iron (more reactive than copper).
Hence, order of increasing reactivity: $\mathbf{\text{D} < \text{C} < \text{A} < \text{B}}$.

1. Conservation of finite natural ores: Mineral reserves of ores like bauxite and hematite are non-renewable; recycling extends their lifespan.
2. Substantial energy conservation: Recycling scrap metal requires a small fraction (e.g., recycling aluminium saves $95\%$ energy) of the massive electrical/thermal energy needed for primary extraction.
3. Reduction in greenhouse gas and toxic emissions: Smelting and roasting release huge quantities of $\text{CO}_2$ and $\text{SO}_2$ (causing acid rain). Recycling produces minimal emissions.
4. Mitigation of landscape degradation and waste: Eliminates gigantic volumes of mining overburden, mine tailings, toxic slurry, and slag dumped in landfills.

• Cathode (reduction of aluminium cation):
$$\text{Al}^{3+} + 3e^- \longrightarrow \text{Al}(l)$$
Molten aluminium sinks to the bottom of the cell.
• Anode (oxidation of oxide anion):
$$2\text{O}^{2-} \longrightarrow \text{O}_2(g) + 4e^-$$
(Followed by reaction of oxygen with carbon anodes: $\text{C} + \text{O}_2 \rightarrow \text{CO}_2$).

1. Lowers melting temperature: Adding cryolite ($\text{Na}_3\text{AlF}_6$) drops the operating melting point from over $2050^\circ\text{C}$ down to roughly $950^\circ\text{C}$, drastically reducing fuel/energy costs.
2. Increases electrical conductivity: Provides mobile $\text{Na}^+$ and fluoroaluminate ions that allow easy flow of electric current through the molten bath.

Combustion of Carbon Anodes:

During electrolysis, oxygen gas is evolved at the positive graphite anodes at temperatures around $950^\circ\text{C}$:
$$2\text{O}^{2-} \longrightarrow \text{O}_2 + 4e^-$$
At this elevated temperature, the evolved oxygen readily reacts with the carbon of the anode rods, burning them into carbon dioxide gas:
$$\text{C}(s) + \text{O}_2(g) \longrightarrow \text{CO}_2(g)$$
Consequently, the carbon anodes are progressively eaten away and must be continually replaced.

1. Charges and Ion Symbols:
• Aluminium ion: $\mathbf{\text{Al}^{3+}}$ (cation)
• Oxide ion: $\mathbf{\text{O}^{2-}}$ (anion)

2. Electron Transfer Mechanism:
• Electronic configurations: $\text{Al}: 2, 8, 3$ and $\text{O}: 2, 6$.
• Each aluminium atom loses 3 valence electrons to achieve the stable neon octet configuration ($2, 8$):
$$\text{Al} \longrightarrow \text{Al}^{3+} + 3e^-$$
• Each oxygen atom requires 2 electrons to complete its octet ($2, 8$):
$$\text{O} + 2e^- \longrightarrow \text{O}^{2-}$$
• To balance total electrons transferred, the lowest common multiple of 3 and 2 is 6 electrons:
$$2\text{Al} \; (2 \times 3e^- = 6e^-) + 3\text{O} \; (3 \times 2e^- = 6e^-) \longrightarrow 2\text{Al}^{3+} + 3\text{O}^{2-} \longrightarrow \mathbf{\text{Al}_2\text{O}_3}$$