Write short notes on any two. [2x5=10] i) Characteristics of homologous series ii) Laboratory preparation of ethyne. iii) Mendeleev's periodic table and its anomalies. iv) Law of mass action and…

i) Characteristics of Homologous Series

A homologous series is a group of organic compounds that share the same general formula, have similar chemical properties, and differ by a –CH₂– unit in their molecular structure. The following are the key characteristics of a homologous series:

1. Same General Formula

• All members of a homologous series follow a common molecular formula.
• Example:
  • Alkanes: $\mathrm{CnH}{}_{2n+2}$ (e.g., methane $\mathrm{CH}{}_4$, ethane $\mathrm{C}{}_2\mathrm{H}{}_6$, propane $\mathrm{C}{}_3\mathrm{H}{}_8$)
  • Alkenes: $\mathrm{CnH}{}_{2n}$ (e.g., ethene $\mathrm{C}{}_2\mathrm{H}{}_4$, propene $\mathrm{C}{}_3\mathrm{H}{}_6$)
  • Alkynes: $\mathrm{CnH}{}_{2n-2}$ (e.g., ethyne $\mathrm{C}{}_2\mathrm{H}{}_2$, propyne $\mathrm{C}{}_3\mathrm{H}{}_4$)

Methane (CH₄)Methane (CH₄)Ethane (C₂H₆)Ethane (C₂H₆)Propane (C₃H₈)Propane (C₃H₈)Ethene (C₂H₄)Ethene (C₂H₄)Propene (C₃H₆)Propene (C₃H₆)Ethyne (C₂H₂)Ethyne (C₂H₂)

2. Similar Chemical Properties

• Compounds in the same series undergo similar chemical reactions due to the presence of the same functional group.
• Example:
  • All alkanes undergo combustion and substitution reactions.
  • All alkenes undergo addition reactions (e.g., hydrogenation, halogenation).

3. Gradual Change in Physical Properties

• While chemical properties remain similar, physical properties (like melting point, boiling point, density, and solubility) change gradually with increasing molecular mass.
• Example:
  • Boiling point increases as the chain length increases (e.g., methane $\mathrm{CH}{}_4$ boils at -161.5°C, while octane $\mathrm{C}{}_8\mathrm{H}{}_{18}$ boils at 125.6°C).
  • Solubility in water decreases as the hydrocarbon chain length increases.

4. Difference by a –CH₂– Unit

• Each successive member differs from the previous one by a –CH₂– group.
• Example:
  • Methane ($\mathrm{CH}{}_4$) → Ethane ($\mathrm{C}{}_2\mathrm{H}{}_6$) → Propane ($\mathrm{C}{}_3\mathrm{H}{}_8$) → Butane ($\mathrm{C}{}_4\mathrm{H}{}_{10}$)

5. Uniform Increase in Molecular Mass

• The molecular mass increases by 14 u (mass of –CH₂–) as we move from one member to the next.
• Example:
  • Methane (16 u) → Ethane (30 u) → Propane (44 u)

6. Same Type of Bonding

• All members have the same type of bonding (e.g., only single bonds in alkanes, one double bond in alkenes, etc.).

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ii) Laboratory Preparation of Ethyne (Acetylene)

Ethyne ($\mathrm{C}{}_2\mathrm{H}{}_2$) is prepared in the laboratory by the hydrolysis of calcium carbide ($\mathrm{CaC}{}_2$). The reaction is as follows:

Chemical Reaction:

$$\mathrm{CaC}{}_2+2\mathrm{H}{}_2\mathrm{O}\to \mathrm{C}{}_2\mathrm{H}{}_2+\mathrm{Ca}(\mathrm{OH}){}_2$$

Procedure:

1. Apparatus Setup:

   • A round-bottom flask or test tube is used.
   • A dropping funnel is attached to the flask to add water slowly.
   • A delivery tube is connected to the flask, leading to an inverted gas jar filled with water (for collecting the gas).

2. Reagents:

   • Calcium carbide ($\mathrm{CaC}{}_2$): A grayish-white solid.
   • Water ($\mathrm{H}{}_2\mathrm{O}$): Added dropwise to control the reaction.

3. Process:

   • A small amount of calcium carbide is placed in the flask.
   • Water is added dropwise from the dropping funnel.
   • Ethyne gas is evolved, which is collected by the displacement of water (since ethyne is less dense than water and slightly soluble).

4. Purification (if needed):

   • The gas may be passed through concentrated sulfuric acid to remove moisture.
   • Impurities like hydrogen sulfide ($\mathrm{H}{}_2\mathrm{S}$) and phosphine ($\mathrm{PH}{}_3$) can be removed by passing the gas through acidified copper sulfate solution or lead nitrate solution.

Chemical Test for Ethyne:

• Ethyne burns with a sooty flame due to incomplete combustion.
• It decolorizes bromine water (addition reaction).
• It forms a white precipitate of silver acetylide when treated with ammoniacal silver nitrate solution: $$\mathrm{C}{}_2\mathrm{H}{}_2+2[\mathrm{Ag}(\mathrm{NH}{}_3){}_2]\mathrm{OH}\to \mathrm{Ag}{}_2\mathrm{C}{}_2(\mathrm{white}\mathrm{ppt})+4\mathrm{NH}{}_3+2\mathrm{H}{}_2\mathrm{O}$$

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iii) Mendeleev’s Periodic Table and Its Anomalies

Mendeleev’s Periodic Table (1869):

• Arrangement: Elements were arranged in increasing atomic mass and grouped based on similar chemical properties.
• Key Features:
  • Periods (Rows): 7 horizontal rows (representing electron shells).
  • Groups (Columns): 8 vertical columns (elements with similar properties).
  • Gaps Left for Undiscovered Elements: Mendeleev predicted the existence of gallium (Ga), germanium (Ge), and scandium (Sc) based on trends.
  • Corrected Atomic Masses: Some elements (e.g., tellurium (Te) and iodine (I)) were placed in order of properties rather than mass.

Anomalies in Mendeleev’s Periodic Table:

Despite its success, Mendeleev’s table had some limitations and anomalies:

1. Position of Hydrogen:

   • Hydrogen has properties of both alkali metals (Group 1) and halogens (Group 17).
   • Mendeleev placed it in Group 1, but it does not fit perfectly.

2. Increasing Atomic Mass Order:

   • Some elements with higher atomic masses were placed before those with lower masses if their properties matched.
   • Example:
     • Tellurium (Te, 127.6 u) was placed before iodine (I, 126.9 u) because Te resembles oxygen (Group 16), while I resembles fluorine (Group 17).

3. Isotopes Were Unknown:

   • Mendeleev did not account for isotopes (atoms of the same element with different masses).
   • Example: Argon (Ar, 40 u) has a higher mass than potassium (K, 39.1 u), but Mendeleev placed K before Ar based on properties.

4. No Explanation for Periodicity:

   • Mendeleev’s table was empirical (based on observations) and did not explain why properties repeated periodically.

5. Lanthanides and Actinides:

   • The f-block elements (lanthanides and actinides) were not properly accommodated in Mendeleev’s original table.

6. Noble Gases Were Missing:

   • Noble gases (e.g., helium, neon, argon) were not discovered at the time (discovered later by Ramsay, 1894).

Later Modifications:

• Henry Moseley (1913) rearranged the table based on atomic number (Z) instead of atomic mass, resolving most anomalies.
• Modern Periodic Law: Elements are arranged in increasing atomic number, and properties are periodic functions of atomic number.

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iv) Law of Mass Action and Equilibrium Constant

N₂+3H₂High pressure, catalyst (Fe)2NH₃

Law of Mass Action:

• Proposed by Norwegian chemists Cato Guldberg and Peter Waage (1864).

• States that:

  "The rate of a chemical reaction is directly proportional to the product of the active masses (concentrations) of the reacting substances."

• For a general reaction: $$aA+bB\rightleftharpoons cC+dD$$

  • Forward rate ($r_f$) ∝ $[A{]}^a[B{]}^b$
  • Backward rate ($r_b$) ∝ $[C{]}^c[D{]}^d$

• At equilibrium, the forward and backward rates are equal: $$r_f=r_b⟹k_f[A{]}^a[B{]}^b=k_b[C{]}^c[D{]}^d$$ $$\frac{k_f}{k_b}=\frac{[C{]}^c[D{]}^d}{[A{]}^a[B{]}^b}=K_{eq}$$ Where:

  • $K_{eq}$ = Equilibrium constant
  • $k_f$ = Forward rate constant
  • $k_b$ = Backward rate constant

Equilibrium Constant ($K_{eq}$):

• Definition: A constant value at a given temperature that relates the concentrations of reactants and products at equilibrium.
• Expression for the reaction $aA+bB\rightleftharpoons cC+dD$: $$K_{eq}=\frac{[C{]}^c[D{]}^d}{[A{]}^a[B{]}^b}$$
• Characteristics:
  1. Dependent on Temperature: Changes with temperature.
  2. Dimensionless (if concentrations are in mol/L): Often expressed in terms of molar concentrations (for gases, partial pressures can also be used).
  3. Does Not Depend on Initial Concentrations: Only depends on the stoichiometry of the reaction.
  4. Predicts Direction of Reaction:
     • If $Q<K_{eq}$, reaction proceeds forward (toward products).
     • If $Q>K_{eq}$, reaction proceeds backward (toward reactants).
     • If $Q=K_{eq}$, the system is at equilibrium.

Types of Equilibrium Constants:

1. Concentration Equilibrium Constant ($K_c$):

   • Used for reactions in solution phase.
   • Example: $$\mathrm{N}{}_2(\mathrm{g})+3\mathrm{H}{}_2(\mathrm{g})\rightleftharpoons 2\mathrm{NH}{}_3(\mathrm{g})$$ $$K_c=\frac{[\mathrm{NH}{}_3{]}^2}{[\mathrm{N}{}_2][\mathrm{H}{}_2{]}^3}$$

2. Pressure (or Gas) Equilibrium Constant ($K_p$):

   • Used for gaseous reactions, where partial pressures are used instead of concentrations.
   • Relation between $K_p$ and $K_c$: $$K_p=K_c(RT{)}^{\Delta n}$$ Where:
     • $R$ = Gas constant (0.0821 L·atm·K⁻¹·mol⁻¹)
     • $T$ = Temperature (K)
     • $\Delta n$ = Change in moles of gas ($\Delta n=\text{moles of gaseous products}-\text{moles of gaseous reactants}$)

3. Solubility Product Constant ($K_{sp}$):

   • For slightly soluble salts (e.g., $\mathrm{AgCl}$): $$\mathrm{AgCl}(\mathrm{s})\rightleftharpoons \mathrm{Ag}{}^{+}(\mathrm{aq})+\mathrm{Cl}{}^{-}(\mathrm{aq})$$ $$K_{sp}=[\mathrm{Ag}{}^{+}][\mathrm{Cl}{}^{-}]$$

4. Acid Dissociation Constant ($K_a$) and Base Dissociation Constant ($K_b$):

   • For weak acids and bases: $$\mathrm{HA}\rightleftharpoons \mathrm{H}{}^{+}+\mathrm{A}{}^{-}$$ $$K_a=\frac{[\mathrm{H}{}^{+}][\mathrm{A}{}^{-}]}{[\mathrm{HA}]}$$

Applications of Law of Mass Action:

1. Predicting Reaction Direction: Helps determine whether a reaction will proceed forward or backward.
2. Calculating Equilibrium Concentrations: Used to find unknown concentrations at equilibrium.
3. Industrial Processes: Optimizes conditions (e.g., Habit’s process for ammonia synthesis).
4. Solubility and Precipitation: Explains why some salts dissolve and others precipitate.

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CaC₂+2H₂OLaboratory setup: CaC₂ in waterC₂H₂+Ca(OH)₂