Metals and Non-metals: Chemical Properties, Reactivity, Ionic Compounds, and Metallurgy

Metals and non-metals form the structural and chemical bedrock of our physical universe. From the iron that forms Earth's core and carries oxygen in our blood to the silicon that powers our microprocessors, understanding how these elements interact chemically is central to inorganic chemistry.

While physical properties such as malleability, ductility, and electrical conductivity allow us to classify elements broadly, their chemical behavior—specifically how they lose or gain electrons—defines their true identity.

This guide covers the core concepts of the NCERT Class 10 syllabus: chemical properties of metals and non-metals, the reactivity series, the mechanics of ionic bonding, and the metallurgical processes used to extract pure metals from natural ores.

1. Chemical Properties of Metals

Metals are electropositive elements. They tend to lose electrons from their valence shell to attain a stable octet, forming positively charged ions (cations):

$M \rightarrow M^{n+} + n e^{-}$

Because of this electron-donating capability, metals act as powerful reducing agents.

                           METALLIC REACTIVITY WITH OXYGEN
  
  [ Na, K ]  --------------> React vigorously at room temperature (Stored in Kerosene)
  [ Mg, Al, Zn, Pb ] ------> Form protective surface oxide layers (Passivation)
  [ Fe, Cu ] --------------> Do not burn easily; Fe filings burn brightly; Cu forms black CuO layer
  [ Ag, Au ] --------------> Completely unreactive with oxygen even at high temperatures

A. Reaction of Metals with Oxygen (Air)

Almost all metals combine with oxygen to form metal oxides.

$\text{Metal} + \text{Oxygen} \rightarrow \text{Metal Oxide}$

Basic Nature of Metal Oxides

Most metal oxides are basic in nature. When dissolved in water, soluble metal oxides (known as alkalis) produce hydroxide ions ($OH^-$):

$4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$

$Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$

$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

Amphoteric Oxides

Certain metal oxides exhibit both basic and acidic behavior. These are called amphoteric oxides. They react with both acids and bases to produce salt and water. Aluminium oxide ($Al_2O_3$) and zinc oxide ($ZnO$) are classic examples.

• Reaction of $Al_2O_3$ with an acid (HCl): $Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)$

• Reaction of $Al_2O_3$ with a base (NaOH): $Al_2O_3(s) + 2NaOH(aq) \rightarrow 2NaAlO_2(aq) + H_2O(l)$ (where $NaAlO_2$ is Sodium Aluminate)

• Reaction of $ZnO$ with an acid (HCl): $ZnO(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2O(l)$

• Reaction of $ZnO$ with a base (NaOH): $ZnO(s) + 2NaOH(aq) \rightarrow Na_2ZnO_2(aq) + H_2O(l)$ (where $Na_2ZnO_2$ is Sodium Zincate)

B. Reaction of Metals with Water

Metals react with water to form metal oxides or metal hydroxides, releasing hydrogen gas.

$\text{Metal} + \text{Water} \rightarrow \text{Metal Hydroxide / Metal Oxide} + \text{Hydrogen}$

                               REACTION WITH WATER
  
  Cold Water   --------> Sodium (Na), Potassium (K), Calcium (Ca)
  Hot Water    --------> Magnesium (Mg)
  Steam Only   --------> Aluminium (Al), Iron (Fe), Zinc (Zn)
  No Reaction  --------> Lead (Pb), Copper (Cu), Silver (Ag), Gold (Au)

1. Vigorous reaction with cold water (Sodium and Potassium): $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g) + \text{Heat energy}$ $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g) + \text{Heat energy}$ Note: The reaction is so exothermic that the evolved hydrogen gas immediately catches fire.

2. Less violent reaction with cold water (Calcium): $Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$ Note: Calcium starts floating because bubbles of hydrogen gas stick to its surface.

3. Reaction with hot water (Magnesium): Magnesium does not react with cold water; it reacts with hot water to form magnesium hydroxide and hydrogen. Like calcium, it also floats due to attached hydrogen bubbles. $Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g)$

4. Reaction with steam (Aluminium, Iron, Zinc): These metals do not react with cold or hot water, but react with steam to form the respective metal oxide and hydrogen gas. $2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$ $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$ (where $Fe_3O_4$ is Iron(II,III) oxide or ferroso-ferric oxide).

C. Reaction of Metals with Dilute Acids

Metals situated above hydrogen in the reactivity series displace hydrogen from dilute acids to form metal salts and hydrogen gas.

$\text{Metal} + \text{Dilute Acid} \rightarrow \text{Salt} + \text{Hydrogen gas}$

$Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$

$Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$

The Exceptional Case of Nitric Acid ($HNO_3$)

Hydrogen gas is not usually evolved when a metal reacts with dilute nitric acid ($HNO_3$).

This occurs because $HNO_3$ is a powerful oxidizing agent. It oxidizes the evolved $H_2$ gas into water ($H_2O$) and itself gets reduced to any of the nitrogen oxides ($NO_2$, $NO$, or $N_2O$).

Exceptions: Very dilute nitric acid ($1\%$) reacts with Magnesium ($Mg$) and Manganese ($Mn$) to liberate hydrogen gas:

$Mg(s) + 2HNO_3(\text{very dilute}) \rightarrow Mg(NO_3)_2(aq) + H_2(g)$

$Mn(s) + 2HNO_3(\text{very dilute}) \rightarrow Mn(NO_3)_2(aq) + H_2(g)$

D. Reaction of Metals with Solutions of Other Metal Salts

A more reactive metal displaces a less reactive metal from its aqueous salt solution. This is a single displacement reaction.

$\text{Metal } A + \text{Salt solution of Metal } B \rightarrow \text{Salt solution of Metal } A + \text{Metal } B$

$Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$ Observation: The blue color of $CuSO_4$ fades to light green ($FeSO_4$), and reddish-brown copper deposits on the iron nail.

2. Chemical Properties of Non-Metals

Non-metals are electronegative elements. They tend to gain electrons to complete their outer octet, forming negatively charged ions (anions):

$X + n e^{-} \rightarrow X^{n-}$

Because non-metals act as electron acceptors, they function as oxidizing agents.

A. Reaction of Non-Metals with Oxygen

Non-metals react with oxygen to form non-metallic oxides which are either acidic or neutral in nature. Non-metal oxides are covalent compounds.

Acidic Oxides

They dissolve in water to produce acids, turning blue litmus red.

$C(s) + O_2(g) \rightarrow CO_2(g)$

$CO_2(g) + H_2O(l) \rightarrow H_2CO_3(aq) \quad \text{(Carbonic Acid)}$

$S(s) + O_2(g) \rightarrow SO_2(g)$

$SO_2(g) + H_2O(l) \rightarrow H_2SO_3(aq) \quad \text{(Sulfurous Acid)}$

Neutral Oxides

These oxides do not show any acidic or basic properties and do not change the color of litmus paper. Examples include Carbon Monoxide ($CO$), Water ($H_2O$), Nitric Oxide ($NO$), and Dinitrogen Oxide ($N_2O$).

$2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$

$2C(s) + O_2(g) \; (\text{insufficient } O_2) \rightarrow 2CO(g)$

B. Summary of Chemical Differences: Metals vs. Non-Metals

3. The Reactivity Series of Metals

The Reactivity Series (or Activity Series) is an arrangement of metals in the order of their decreasing chemical reactivity. Hydrogen is included in this series because, like metals, it can lose an electron to form a positive ion ($H^+$).

   Metal      Symbol    Reactivity Level
  ─────────────────────────────────────────────────────────────
   Potassium    K       Most Reactive  ▲
   Sodium       Na                     │
   Calcium      Ca                     │  Extracted by Electrolysis
   Magnesium    Mg                     │  (High affinity for oxygen)
   Aluminium    Al                     │
  ─────────────────────────────────────────────────────────────
   Zinc         Zn                     │
   Iron         Fe                     │  Extracted by Reduction using
   Lead         Pb                     │  Carbon/Monoxide
   [Hydrogen]  [H]                     │
   Copper       Cu                     │
  ─────────────────────────────────────────────────────────────
   Mercury      Hg                     │  Extracted by Thermal Decomposition
   Silver       Ag                     │  Found in Native/Free State
   Gold         Au      Least Reactive ▼

Key Principles of the Reactivity Series

1. Displacement Power: Any metal placed higher in the series can displace any metal placed below it from its aqueous salt solution. $\text{Example: } Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s) \quad (\text{Occurs because } Zn > Cu)$ $Cu(s) + ZnSO_4(aq) \rightarrow \text{No Reaction} \quad (\text{Because } Cu < Zn)$
2. Reaction with Hydrogen Ions: Metals above Hydrogen displace $H_2$ gas from dilute acids ($HCl, H_2SO_4$). Metals below Hydrogen ($Cu, Hg, Ag, Au$) do not react with dilute acids to release $H_2$ gas.

4. Formation and Properties of Ionic Compounds

When metals react with non-metals, electrons are transferred completely from the valence shell of the metal atoms to the valence shell of the non-metal atoms. This transfer forms stable ionic (or electrovalent) bonds.

A. Step-by-Step Formation of Ionic Compounds

1. Formation of Sodium Chloride ($NaCl$)

• Sodium ($Na$): Atomic number = 11. Electronic configuration = $(2, 8, 1)$. It loses 1 electron to achieve the noble gas configuration of Neon $(2, 8)$: $Na (2, 8, 1) \rightarrow Na^+ (2, 8) + e^{-}$

• Chlorine ($Cl$): Atomic number = 17. Electronic configuration = $(2, 8, 7)$. It needs 1 electron to complete its octet and achieve the configuration of Argon $(2, 8, 8)$: $Cl (2, 8, 7) + e^{-} \rightarrow Cl^- (2, 8, 8)$

• Bond Formation: $Na^{\bullet} + \cdot\underset{\bullet\bullet}{\overset{\bullet\bullet}{\text{Cl}}}\bullet \longrightarrow [Na]^+ \left[ :\underset{\bullet\bullet}{\overset{\bullet\bullet}{\text{Cl}}}: \right]^- \longrightarrow NaCl$

2. Formation of Magnesium Chloride ($MgCl_2$)

• Magnesium ($Mg$): Atomic number = 12. Electronic configuration = $(2, 8, 2)$. $Mg (2, 8, 2) \rightarrow Mg^{2+} (2, 8) + 2e^{-}$

• Chlorine ($Cl$): Two chlorine atoms accept one electron each: $2Cl (2, 8, 7) + 2e^{-} \rightarrow 2Cl^- (2, 8, 8)$

• Bond Formation: $Mg^{\bullet\bullet} + 2 \cdot\underset{\bullet\bullet}{\overset{\bullet\bullet}{\text{Cl}}}\bullet \longrightarrow [Mg]^{2+} \left[ :\underset{\bullet\bullet}{\overset{\bullet\bullet}{\text{Cl}}}: \right]^-_2 \longrightarrow MgCl_2$

B. Properties of Ionic Compounds

                             PROPERTIES OF IONIC COMPOUNDS
  ┌───────────────────────┬───────────────────────┬───────────────────────┐
  │ Physical Nature &     │ Melting & Boiling     │ Electrical            │
  │ Hardness              │ Points                │ Conductivity          │
  ├───────────────────────┼───────────────────────┼───────────────────────┤
  │ Solid & Brittle due   │ High due to strong    │ Conducts in Molten &  │
  │ to strong inter-ionic │ electrostatic forces  │ Aqueous states only;  │
  │ attractive forces.    │ between ions.         │ NOT in solid state.   │
  └───────────────────────┴───────────────────────┴───────────────────────┘

1. Physical State and Brittleness: Ionic compounds are solids and hard due to strong forces of attraction between positive and negative ions. They are generally brittle and break into pieces when pressure is applied.
2. High Melting and Boiling Points: A significant amount of thermal energy is required to break the strong inter-ionic electrostatic forces.
   • Example: $NaCl$ has a melting point of $1074\text{ K } (801^\circ\text{C})$.
3. Solubility: Electrovalent compounds are soluble in polar solvents like water, but insoluble in non-polar organic solvents such as kerosene, petrol, and benzene.
4. Conduction of Electricity:
   • Solid State: Ionic compounds do not conduct electricity because ions are fixed in a rigid crystal lattice structure and cannot move.
   • Molten or Dissolved State: Electrostatic forces are broken by heat or water, allowing ions to move freely towards electrodes when electric current passes through them.

5. Basic Metallurgical Processes

Metallurgy is the complete scientific and technological sequence used to extract pure metals from their naturally occurring ores.

                                  STEPS IN METALLURGY
                                           │
                                     [ ORE DEPOSIT ]
                                           │
                                  1. Enrichment / Concentration
                                           │
                   ┌───────────────────────┼───────────────────────┐
                   ▼                       ▼                       ▼
            Metals of High          Metals of Medium         Metals of Low
              Reactivity               Reactivity              Reactivity
            (Na, Ca, Mg, Al)          (Zn, Fe, Pb)              (Hg, Cu)
                   │                       │                       │
           Electrolysis of            ┌────┴────┐           Thermal Reduction
             Molten Ore               ▼         ▼             (Direct Heating)
                   │              Carbonate  Sulfide               │
              Pure Metal             Ore       Ore                 │
                                      │         │                  │
                                 Calcination Roasting              │
                                      └────┬────┘                  │
                                           ▼                       │
                                      Metal Oxide                  │
                                           │                       │
                                      Reduction via                │
                                       Carbon/Al                   │
                                           │                       │
                                      Crude Metal                  │
                                           └───────────┬───────────┘
                                                       ▼
                                               3. Refining (Electrolytic)
                                                       │
                                                   PURE METAL

Essential Terminology

• Minerals: Naturally occurring inorganic elements or compounds found in the Earth's crust.
• Ores: Minerals that contain a sufficiently high percentage of a specific metal, allowing the metal to be extracted profitably. (Note: All ores are minerals, but not all minerals are ores).
• Gangue: Earthy impurities such as sand, soil, clay, and silica associated with mined ores.

Step 1: Concentration or Enrichment of Ore

The physical process of removing gangue from mined ore prior to chemical reduction. Methods include hydraulic washing, magnetic separation, and froth flotation (depending on differences in physical/chemical properties).

Step 2: Extraction of Crude Metal from Concentrated Ore

A. Extracting Metals Low in the Reactivity Series

Metals low in the series are unreactive. Their oxides can be reduced to metal by heating alone (thermal reduction).

1. Extraction of Mercury from Cinnabar ($HgS$):

   • Heating in air: $2HgS(s) + 3O_2(g) \xrightarrow{\Delta} 2HgO(s) + 2SO_2(g)$
   • Further heating reduces mercuric oxide to mercury: $2HgO(s) \xrightarrow{\Delta} 2Hg(l) + O_2(g)$

2. Extraction of Copper from Copper Pyrites/Glance ($Cu_2S$): $2Cu_2S(s) + 3O_2(g) \xrightarrow{\Delta} 2Cu_2O(s) + 2SO_2(g)$ $2Cu_2O(s) + Cu_2S(s) \xrightarrow{\Delta} 6Cu(s) + SO_2(g)$

B. Extracting Metals in the Middle of the Reactivity Series

Metals such as $Zn, Fe, Pb$ occur as sulfides or carbonates. It is easier to extract a metal from its oxide than from its sulfide or carbonate. Therefore, sulfide or carbonate ores must first be converted into metal oxides.

                           CONVERSION OF ORES TO OXIDES
  
      ROASTING                                  CALCINATION
  • Sulfide ores                           • Carbonate ores
  • Heated in EXCESS AIR                   • Heated in LIMITED AIR / NO AIR
  • E.g., 2ZnS + 3O₂ ──Δ──> 2ZnO + 2SO₂    • E.g., ZnCO₃ ──Δ──> ZnO + CO₂

• Roasting: Heating a sulfide ore strongly in the presence of excess air. $2ZnS(s) + 3O_2(g) \xrightarrow{\Delta} 2ZnO(s) + 2SO_2(g)$

• Calcination: Heating a carbonate ore strongly in limited or restricted air. $ZnCO_3(s) \xrightarrow{\Delta} ZnO(s) + CO_2(g)$

Reduction of Metal Oxides to Metals

1. Using Carbon (Coke): $ZnO(s) + C(s) \rightarrow Zn(s) + CO(g)$ $Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(s) + 3CO_2(g)$

2. Using Displacement Reactions (Thermite Process): Highly reactive metals like $Al, Na, Ca$ can be used as reducing agents because they displace metals of lower reactivity from their oxides. These reactions are highly exothermic, yielding metals in their molten state.

   • Thermite Reaction: The reaction of Iron(III) oxide with Aluminium powder: $Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$ Practical Application: Used to weld broken railway tracks or cracked machine parts in-situ.

C. Extracting Metals High in the Reactivity Series

High-reactivity metals ($Na, Ca, Mg, Al$) have a strong chemical affinity for oxygen and cannot be reduced using carbon. Carbon cannot split the strong metal-oxygen bonds in these oxides.

These metals are extracted by Electrolytic Reduction (Electrolysis of their molten salts).

• Extraction of Sodium from Molten $NaCl$: During electrolysis of molten Sodium Chloride:
  • At Cathode (Negative electrode - Reduction occurs): $Na^+ + e^- \rightarrow Na(s)$
  • At Anode (Positive electrode - Oxidation occurs): $2Cl^- \rightarrow Cl_2(g) + 2e^-$

Step 3: Refining of Metals (Electrolytic Refining)

Metals obtained by reduction processes contain impurities. Electrolytic refining is the most widely used method to obtain pure metals (such as $Cu, Zn, Ni, Ag, Au$).

                         ELECTROLYTIC REFINING CELL
  
                  Anode (+)                     Cathode (-)
             [Impure Metal Strip]          [Pure Thin Metal Strip]
                      │                              │
                      └──────────────┬───────────────┘
                                     │
                             Electrolyte Solution
                           (Acidified Metal Salt, e.g., CuSO₄)
  
  • Anode dissolves:   Cu (impure) ──────> Cu²⁺ + 2e⁻
  • Cathode deposits:  Cu²⁺ + 2e⁻  ──────> Cu (pure)
  • Bottom of Anode:   Anode Mud (insoluble impurities like Au, Ag)

Electrolytic Refining of Copper:

• Anode: Thick block of impure copper metal.
• Cathode: Thin sheet of pure copper metal.
• Electrolyte: Acidified solution of Copper Sulfate ($CuSO_4 + H_2SO_4$).
• Mechanism: On passing electric current:
  1. Pure copper from the anode dissolves into the electrolyte: $\text{At Anode (+): } Cu(s) \rightarrow Cu^{2+}(aq) + 2e^{-}$
  2. An equivalent amount of pure copper from the electrolyte deposits onto the cathode: $\text{At Cathode (-): } Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$
  3. Soluble impurities go into the solution, while insoluble impurities settle at the bottom of the anode as Anode Mud (contains valuable metals like gold and silver).

6. Real-World Applications & Analogies

1. The Railway Thermite Welding Process

• Real-World Application: Rail tracks expand and contract with temperature, leading to stress fractures. Instead of replacing entire rail lines, railway technicians place a crucible filled with Iron Oxide ($Fe_2O_3$) and Aluminium powder ($Al$) over the crack. Once ignited, the reaction produces liquid iron at temperatures over $2500^\circ\text{C}$. This molten iron flows directly into the gap, welding the rail joint seamless upon cooling.

2. The Sacrificial Anode (Galvanization & Corrosion Protection)

• Real-World Application: Underground steel pipes and massive iron ship hulls corrode in moist environments. To prevent this, blocks of zinc or magnesium are physically attached to the steel structure. Because zinc and magnesium are higher in the reactivity series than iron, they donate electrons preferentially, oxidizing first. The zinc acts as a "sacrificial anode," protecting the iron structural core from rusting.

3. The Salt Bridge Analogy for Ionic Bonding

• Analogy: Consider two individuals: one holding a heavy package they do not want (Sodium with 1 valence electron), and another desperately needing a package to complete a set (Chlorine with 7 valence electrons). Once the electron package is handed over, both achieve structural stability. However, this transfer leaves one positively charged and the other negatively charged. The resulting strong mutual electrostatic attraction keeps them bonded closely together, similar to two opposite magnetic poles.

7. Step-by-Step Solved Examples

Example 1: Identifying Amphoteric Oxides and Writing Reaction Equations

Question: You are given two unlabeled metal oxides, Oxide $X$ and Oxide $Y$. Oxide $X$ reacts with dilute $HCl$ but shows no reaction with aqueous $NaOH$. Oxide $Y$ reacts with both dilute $HCl$ and aqueous $NaOH$. Identify the chemical nature of Oxide $X$ and Oxide $Y$. Write balanced chemical equations for the reactions of Oxide $Y$ with both $HCl$ and $NaOH$, assuming $Y$ is Aluminium Oxide.

Solution:

• Step 1: Determine the nature of Oxide X. Oxide $X$ reacts only with an acid ($HCl$). Acids react with basic oxides to form salt and water. Therefore, Oxide $X$ is a Basic Oxide.

• Step 2: Determine the nature of Oxide Y. Oxide $Y$ reacts with both an acid ($HCl$) and a base ($NaOH$). Oxides that show both acidic and basic character are amphoteric. Therefore, Oxide $Y$ is an Amphoteric Oxide.

• Step 3: Write the balanced chemical reaction of $Al_2O_3$ with dilute $HCl$. $Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)$

• Step 4: Write the balanced chemical reaction of $Al_2O_3$ with aqueous $NaOH$. $Al_2O_3(s) + 2NaOH(aq) \rightarrow 2NaAlO_2(aq) + H_2O(l)$

Final Answer: Oxide $X$ is a basic oxide; Oxide $Y$ is an amphoteric oxide. The balanced equations yield Aluminium Chloride ($AlCl_3$) and Sodium Aluminate ($NaAlO_2$), respectively.

Example 2: Differentiating Metallurgy of Sulfide vs. Carbonate Ores

Question: Zinc is a metal in the middle of the reactivity series. It is found in nature as Sphalerite ($ZnS$) and Calamine ($ZnCO_3$). Detail the chemical conversion processes used to convert these two ores into Zinc Oxide, including balanced chemical equations with appropriate heat symbols.

Solution:

• Step 1: Process for Sulfide Ore ($ZnS$ - Sphalerite): Sulfide ores require Roasting—heating strongly in the presence of excess oxygen. $\text{Chemical Equation: } 2ZnS(s) + 3O_2(g) \xrightarrow{\Delta} 2ZnO(s) + 2SO_2(g)$

• Step 2: Process for Carbonate Ore ($ZnCO_3$ - Calamine): Carbonate ores require Calcination—heating strongly in limited oxygen or in the absence of air. $\text{Chemical Equation: } ZnCO_3(s) \xrightarrow{\Delta} ZnO(s) + CO_2(g)$

• Step 3: Reduction to Pure Metal: The resulting Zinc Oxide ($ZnO$) from either process is reduced using Carbon (coke) as a reducing agent: $ZnO(s) + C(s) \xrightarrow{\Delta} Zn(s) + CO(g)$

Example 3: Predicting Displacement Reactions

Question: A student placed clean iron nails into four separate test tubes containing solutions of $ZnSO_4$, $CuSO_4$, $Al_2(SO_4)_3$, and $AgNO_3$. In which test tubes will a chemical reaction occur? Write balanced chemical equations and order the metals involved ($Fe, Zn, Cu, Al, Ag$) by increasing reactivity based on the Reactivity Series.

Solution:

• Step 1: Recall the Reactivity Series order for the metals: $Al > Zn > Fe > Cu > Ag$

• Step 2: Evaluate each test tube interaction:

  1. $Fe(s) + ZnSO_4(aq)$: $Fe$ is less reactive than $Zn$. No Reaction.
  2. $Fe(s) + CuSO_4(aq)$: $Fe$ is more reactive than $Cu$. Reaction Occurs. $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$
  3. $Fe(s) + Al_2(SO_4)_3(aq)$: $Fe$ is less reactive than $Al$. No Reaction.
  4. $Fe(s) + AgNO_3(aq)$: $Fe$ is more reactive than $Ag$. Reaction Occurs. $Fe(s) + 2AgNO_3(aq) \rightarrow Fe(NO_3)_2(aq) + 2Ag(s)$

Final Answer: Reactions occur in the test tubes containing $CuSO_4$ and $AgNO_3$. Order of increasing reactivity: $Ag < Cu < Fe < Zn < Al$.

8. Common Student Mistakes to Avoid

  ┌─────────────────────────────────────┬─────────────────────────────────────┐
  │ Common Misconception / Error        │ Correct Chemical Concept            │
  ├─────────────────────────────────────┼─────────────────────────────────────┤
  │ 1. Writing that HNO₃ liberates H₂   │ HNO₃ is a strong oxidizer; it       │
  │    gas with all metals like HCl.    │ oxidizes H₂ to H₂O except with      │
  │                                     │ very dilute HNO₃ + Mg/Mn.           │
  ├─────────────────────────────────────┼─────────────────────────────────────┤
  │ 2. Confusing Roasting and           │ Roasting = Sulfide ore + Excess O₂. │
  │    Calcination conditions.          │ Calcination = Carbonate ore +       │
  │                                     │ Limited/No O₂.                      │
  ├─────────────────────────────────────┼─────────────────────────────────────┤
  │ 3. Stating solid NaCl conducts      │ Solid ionic solids have fixed ions.  │
  │    electricity.                     │ They only conduct when molten or    │
  │                                     │ dissolved in water.                 │
  ├─────────────────────────────────────┼─────────────────────────────────────┤
  │ 4. Thinking Carbon can reduce       │ Metals like Na, Ca, Mg, Al have a   │
  │    highly reactive oxides (Al₂O₃).   │ stronger affinity for O₂ than C does.│
  │                                     │ Electrolytic reduction is required. │
  └─────────────────────────────────────┴─────────────────────────────────────┘

9. Practice Questions for Self-Assessment

Question 1

An element $A$ burns in air with a dazzling white flame to form a white powder $B$. When water is added to $B$, it dissolves partially to form solution $C$, which turns red litmus paper blue.

1. Identify $A$, $B$, and $C$.
2. Write balanced chemical equations for both steps.
3. State whether element $A$ is a metal or a non-metal, giving a chemical reason.

Solution:

1. Identification:
   • Element $A$ is Magnesium ($Mg$).
   • Powder $B$ is Magnesium Oxide ($MgO$).
   • Solution $C$ is Magnesium Hydroxide ($Mg(OH)_2$).
2. Balanced Equations:
   • Step 1 (Burning): $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$
   • Step 2 (Dissolving in water): $MgO(s) + H_2O(l) \rightarrow Mg(OH)_2(aq)$
3. Classification:
   • Element $A$ ($Mg$) is a metal because its oxide ($MgO$) forms a basic hydroxide in water, turning red litmus blue.

Question 2

Using electron-dot structures, show the formation of Calcium Oxide ($CaO$). (Atomic numbers: $Ca = 20$, $O = 8$). Mention the cation and anion present in this compound.

Solution:

• Step 1: Write electronic configurations:

  • $Ca (Z=20): (2, 8, 8, 2) \rightarrow$ Needs to lose $2e^-$ to complete octet.
  • $O (Z=8): (2, 6) \rightarrow$ Needs to gain $2e^-$ to complete octet.

• Step 2: Electron transfer equation: $Ca \rightarrow Ca^{2+} + 2e^{-}$ $O + 2e^{-} \rightarrow O^{2-}$

• Step 3: Electron-dot representation: $Ca^{\bullet\bullet} + \cdot\cdot\underset{\bullet\bullet}{\text{O}}\bullet \longrightarrow [Ca]^{2+} \left[ :\underset{\bullet\bullet}{\overset{\bullet\bullet}{\text{O}}}: \right]^{2-} \longrightarrow CaO$

• Ions Present:

  • Cation: Calcium ion ($Ca^{2+}$)
  • Anion: Oxide ion ($O^{2-}$)

Question 3

Explain the electrolytic refining of crude Copper. Include:

1. Materials used for Anode, Cathode, and Electrolyte.
2. Equations for reactions at the Anode and Cathode.
3. Define 'Anode Mud'.

Solution:

1. Cell Components:
   • Anode: Impure block of Copper metal.
   • Cathode: Pure, thin sheet of Copper metal.
   • Electrolyte: Acidified Copper Sulfate ($CuSO_4$) solution.
2. Electrode Equations:
   • At Anode (+): $Cu(s) \text{ (impure)} \rightarrow Cu^{2+}(aq) + 2e^{-}$
   • At Cathode (-): $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s) \text{ (pure)}$
3. Anode Mud: The insoluble impurities (containing noble metals like silver and gold) that drop down and collect directly below the anode during electrolytic refining.

Question 4

A metal $M$ is found in nature as its carbonate ore $MCO_3$. Outline the metallurgical steps required to obtain the pure metal $M$ from this ore. (Assume $M$ is of medium reactivity).

Solution:

1. Concentration: Remove gangue impurities using physical separation methods.
2. Calcination: Heat the carbonate ore $MCO_3$ strongly in limited air to yield the metal oxide and carbon dioxide: $MCO_3(s) \xrightarrow{\Delta} MO(s) + CO_2(g)$
3. Reduction: Reduce the metal oxide $MO$ using carbon (coke) as a reducing agent: $MO(s) + C(s) \xrightarrow{\Delta} M(s) + CO(g)$
4. Refining: Purify the resulting crude metal $M$ using electrolytic refining to yield the pure metal.

10. Exam Revision & Frequently Asked Questions (FAQs)

Q1: Why are Sodium and Potassium stored under kerosene oil?

Answer: Sodium and Potassium are highly reactive metals at the top of the reactivity series. They react vigorously with oxygen, moisture, and carbon dioxide present in air at room temperature. This reaction generates significant heat, causing evolved hydrogen gas to ignite. Storing them under kerosene isolates them from air and atmospheric moisture, preventing accidental fires.

Q2: What is the Thermite reaction? Write its equation and state one industrial use.

Answer: The Thermite reaction is a highly exothermic single displacement reaction where Iron(III) oxide ($Fe_2O_3$) is reduced by Aluminium powder ($Al$). The heat released is sufficient to produce iron in its liquid molten state.

• Equation: $Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$
• Industrial Application: Used for welding broken railway tracks and cracked structural iron equipment.

Q3: Why do ionic compounds have high melting points, and why do they not conduct electricity in the solid state?

Answer:

1. High Melting Point: Ionic compounds consist of positively and negatively charged ions held together by strong inter-ionic electrostatic forces. A large amount of heat energy is needed to break these strong bonds.
2. Electrical Conductivity: Electrical conduction requires mobile charged particles. In the solid state, ions are locked into fixed positions within a rigid crystal lattice and cannot move. When melted or dissolved in water, the crystal structure breaks down, allowing ions to move freely towards electrodes to conduct electricity.

Q4: Explain why Aluminium cannot be extracted from Aluminium Oxide ($Al_2O_3$) using Carbon reduction.

Answer: Aluminium has a much stronger chemical affinity for oxygen than carbon does. Carbon is unable to break the strong ionic bonds between aluminium and oxygen in $Al_2O_3$. Therefore, Aluminium cannot be reduced by carbon; it must be extracted using Electrolytic Reduction (Hall-Héroult process) by passing an electric current through a molten mixture of alumina ($Al_2O_3$) and cryolite ($Na_3AlF_6$).