Chemistry Cheatsheet

Cheatsheet Content

### Iodine Extraction - **Main Source of Iodine:** The primary commercial source of iodine is naturally occurring brines, particularly those found in oil and gas fields in Japan and Chile, where iodine exists as iodide ions. However, historically and for educational context, **seaweeds** (especially brown seaweeds like kelp, Fucus, and Laminaria) were a significant source. These seaweeds accumulate iodine from seawater. - **Extraction from Seaweeds (Historical Method):** 1. **Drying and Burning:** Large quantities of seaweeds are collected, dried in the sun, and then burnt in shallow pits. The organic matter burns away, leaving behind a gray ash known as "kelp." 2. **Dissolution:** This kelp contains about 0.5-1.5% iodine, primarily in the form of water-soluble metal iodides (e.g., sodium iodide, NaI, and potassium iodide, KI), along with other salts. Hot water is used to extract these soluble iodides from the ash. The solution is then concentrated by evaporation. 3. **Oxidation:** The concentrated solution (often called "mother liquor" or "iodine liquor") is then treated to oxidize the iodide ions ($I^-$) to elemental iodine ($I_2$). Two common methods for this oxidation are: * **Using Manganese Dioxide and Sulfuric Acid:** The solution is heated with a mixture of manganese dioxide ($MnO_2$) and concentrated sulfuric acid ($H_2SO_4$). The manganese dioxide acts as an oxidizing agent: $$2NaI + MnO_2 + 2H_2SO_4 \rightarrow Na_2SO_4 + MnSO_4 + 2H_2O + I_2$$ * **Using Chlorine Gas:** Chlorine gas ($Cl_2$) is bubbled through the solution. Since chlorine is a stronger oxidizing agent than iodine, it displaces iodine from its salts: $$2NaI + Cl_2 \rightarrow 2NaCl + I_2$$ 4. **Purification:** The liberated iodine precipitates as a solid. It is then separated by filtration, dried, and purified by sublimation, yielding dark purple-black crystals of pure iodine. ### First Order Reaction - **Definition:** A first-order reaction is a chemical reaction in which the rate of reaction is directly proportional to the concentration of a single reactant raised to the power of one. This means that if you double the concentration of that reactant, the reaction rate will also double. - **General Rate Law:** For a reaction $A \rightarrow Products$, the rate law for a first-order reaction is expressed as: $$Rate = k[A]^1$$ where $k$ is the rate constant (units: time⁻¹), and $[A]$ is the molar concentration of reactant A. - **Two Examples:** 1. **Radioactive Decay:** All natural and artificial radioactive decay processes follow first-order kinetics. For instance, the decay of carbon-14 ($^{14}C$) used in carbon dating: $$^{14}C \rightarrow ^{14}N + e^-$$ The rate of decay depends only on the amount of $^{14}C$ present. 2. **Decomposition of Dinitrogen Pentoxide:** The thermal decomposition of gaseous dinitrogen pentoxide ($N_2O_5$) is a classic example: $$2N_2O_5(g) \rightarrow 4NO_2(g) + O_2(g)$$ The rate of this reaction is found experimentally to be directly proportional to the concentration of $N_2O_5$. - **Obtain an Expression for the Rate Constant of a First-Order Reaction:** Consider a first-order reaction: $R \rightarrow P$ (Reactant to Products). The differential rate law for this reaction is: $$- \frac{d[R]}{dt} = k[R]$$ Rearranging the equation to separate variables: $$\frac{d[R]}{[R]} = -k dt$$ Now, we integrate both sides. Let $[R]_0$ be the initial concentration of reactant R at time $t=0$, and $[R]_t$ be the concentration of reactant R at any time $t$. $$\int_{[R]_0}^{[R]_t} \frac{d[R]}{[R]} = \int_0^t -k dt$$ Performing the integration: $$[\ln[R]]_{[R]_0}^{[R]_t} = -k[t]_0^t$$ $$\ln[R]_t - \ln[R]_0 = -k(t - 0)$$ This gives the integrated rate law in its logarithmic form: $$\ln \left(\frac{[R]_t}{[R]_0}\right) = -kt$$ To solve for the rate constant $k$: $$k = - \frac{1}{t} \ln \left(\frac{[R]_t}{[R]_0}\right)$$ Alternatively, we can write it as: $$k = \frac{1}{t} \ln \left(\frac{[R]_0}{[R]_t}\right)$$ If we convert from natural logarithm (ln) to common logarithm (log₁₀) by using $\ln x = 2.303 \log_{10} x$: $$k = \frac{2.303}{t} \log_{10} \left(\frac{[R]_0}{[R]_t}\right)$$ This equation allows us to calculate the rate constant $k$ for a first-order reaction if we know the initial concentration, the concentration at a specific time, and the time elapsed. ### Sulphuric Acid (Contact Process) - **Principle of Manufacture of Sulphuric Acid by Contact Process:** The Contact Process is the modern industrial method for the production of sulfuric acid. Its principle relies on the catalytic oxidation of sulfur dioxide ($SO_2$) to sulfur trioxide ($SO_3$) using a vanadium pentoxide ($V_2O_5$) catalyst, followed by the absorption of $SO_3$ into concentrated sulfuric acid to form oleum, which is then diluted with water to produce sulfuric acid. The key steps are: 1. **Production of Sulfur Dioxide ($SO_2$):** Sulfur dioxide is produced by burning sulfur in air or by roasting sulfide ores (e.g., iron pyrites). $$S(s) + O_2(g) \rightarrow SO_2(g)$$ $$4FeS_2(s) + 11O_2(g) \rightarrow 2Fe_2O_3(s) + 8SO_2(g)$$ 2. **Catalytic Oxidation of Sulfur Dioxide to Sulfur Trioxide ($SO_3$):** This is the most crucial step. Purified $SO_2$ is mixed with air (or oxygen) and passed over a catalyst, typically vanadium pentoxide ($V_2O_5$), at an optimal temperature of 400-450°C and pressure of 1-2 atm. This reaction is reversible and exothermic, so careful control of temperature is essential to maximize yield. $$2SO_2(g) + O_2(g) \xrightarrow{V_2O_5, 400-450^\circ C} 2SO_3(g) \quad \Delta H = -196.6 \text{ kJ/mol}$$ 3. **Absorption of Sulfur Trioxide ($SO_3$):** $SO_3$ gas is not absorbed directly in water to produce sulfuric acid because it forms a fine mist of $H_2SO_4$ that is difficult to condense. Instead, $SO_3$ is absorbed in concentrated sulfuric acid (98%) to produce pyrosulfuric acid, also known as oleum ($H_2S_2O_7$). $$SO_3(g) + H_2SO_4(conc.) \rightarrow H_2S_2O_7(l)$$ 4. **Dilution of Oleum:** The oleum is then diluted with a calculated amount of water to produce sulfuric acid of the desired concentration (usually 98%). $$H_2S_2O_7(l) + H_2O(l) \rightarrow 2H_2SO_4(aq)$$ - **Reaction of Concentrated Sulphuric Acid with Oxalic Acid:** Concentrated sulfuric acid is a powerful dehydrating agent. When it reacts with oxalic acid (ethanedioic acid), it removes water molecules from the oxalic acid, leading to its decomposition. $$(COOH)_2 \xrightarrow{Conc. H_2SO_4, \Delta} CO(g) + CO_2(g) + H_2O(l)$$ The water molecule formed remains associated with the sulfuric acid, effectively diluting it. The gases produced are carbon monoxide ($CO$) and carbon dioxide ($CO_2$). ### Sulphuric Acid (Lead Chamber Process) - **How Sulphuric Acid is Prepared by Lead Chamber Process:** The Lead Chamber Process is an older industrial method for producing sulfuric acid, largely superseded by the Contact Process due to its lower efficiency and production of less concentrated acid. It involves the oxidation of sulfur dioxide ($SO_2$) using oxides of nitrogen (nitric oxide, $NO$, and nitrogen dioxide, $NO_2$) as catalysts, all occurring within large lead-lined chambers. - **Principle of the Lead Chamber Process:** The process relies on the catalytic action of nitrogen oxides. Sulfur dioxide, air, and steam are introduced into lead chambers where nitrogen oxides facilitate the oxidation of $SO_2$ to $SO_3$. The $SO_3$ then combines with water vapor to form sulfuric acid. - **Key Steps and Reactions:** 1. **Production of Sulfur Dioxide ($SO_2$):** Similar to the Contact Process, $SO_2$ is produced by burning sulfur or roasting sulfide ores. $$S(s) + O_2(g) \rightarrow SO_2(g)$$ 2. **Production and Regeneration of Nitrogen Oxides:** Nitric oxide ($NO$) is typically produced by the catalytic oxidation of ammonia (Ostwald process) or by the decomposition of nitric acid. $$4NH_3(g) + 5O_2(g) \xrightarrow{Pt, \Delta} 4NO(g) + 6H_2O(g)$$ In the chambers, $NO$ reacts with atmospheric oxygen to form nitrogen dioxide ($NO_2$). $$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$$ 3. **Oxidation of Sulfur Dioxide:** Nitrogen dioxide acts as an oxidizing agent, transferring oxygen to $SO_2$ to form $SO_3$, and in the process, $NO_2$ is reduced back to $NO$. $$SO_2(g) + NO_2(g) \rightarrow SO_3(g) + NO(g)$$ 4. **Formation of Sulfuric Acid:** The $SO_3$ then reacts with water vapor (steam) present in the chambers to form sulfuric acid. $$SO_3(g) + H_2O(g) \rightarrow H_2SO_4(aq)$$ 5. **Recycling of Nitrogen Oxides:** The $NO$ produced in step 3 is re-oxidized by atmospheric oxygen (step 2) to $NO_2$, allowing the cycle to continue. This demonstrates the catalytic role of nitrogen oxides. - **Overall:** The process produces sulfuric acid of about 60-78% concentration, which is less pure and less concentrated than that produced by the Contact Process. ### Nitric Acid (Ostwald's Process) - **Principle of Manufacture of Nitric Acid from Ammonia (Ostwald's Process):** The Ostwald process is the industrial method for producing nitric acid ($HNO_3$). It involves three main stages: the catalytic oxidation of ammonia ($NH_3$) to nitric oxide ($NO$), the oxidation of nitric oxide to nitrogen dioxide ($NO_2$), and finally, the absorption of nitrogen dioxide in water to form nitric acid. - **Key Steps and Reactions:** 1. **Catalytic Oxidation of Ammonia:** Ammonia gas, mixed with air (oxygen), is passed over a heated platinum-rhodium gauze catalyst (typically 90% Pt, 10% Rh) at about 800-900°C and 4-10 atm pressure. This highly exothermic reaction produces nitric oxide ($NO$) and water vapor. $$4NH_3(g) + 5O_2(g) \xrightarrow{Pt/Rh \text{ gauze}, 800-900^\circ C} 4NO(g) + 6H_2O(g) \quad \Delta H = -905 \text{ kJ/mol}$$ This step must be carried out quickly to prevent the further oxidation of NO to N₂. 2. **Oxidation of Nitric Oxide:** The nitric oxide ($NO$) produced is then rapidly cooled and reacted with excess oxygen (from the air) to form nitrogen dioxide ($NO_2$). This reaction is also exothermic and occurs at lower temperatures (around 50°C). $$2NO(g) + O_2(g) \rightarrow 2NO_2(g) \quad \Delta H = -114 \text{ kJ/mol}$$ 3. **Absorption of Nitrogen Dioxide:** The nitrogen dioxide ($NO_2$) gas is then absorbed in water, typically in absorption towers, where it reacts to form nitric acid ($HNO_3$) and regenerates nitric oxide ($NO$). The regenerated $NO$ is then recycled back to step 2. $$3NO_2(g) + H_2O(l) \rightarrow 2HNO_3(aq) + NO(g)$$ The nitric acid produced is typically about 68% by weight. For higher concentrations, it can be dehydrated by distillation with concentrated sulfuric acid. - **Reaction between White Phosphorus and Concentrated Nitric Acid:** Concentrated nitric acid is a strong oxidizing agent. White phosphorus ($P_4$) is highly reactive and readily oxidized. When white phosphorus reacts with hot concentrated nitric acid, it is oxidized to phosphoric acid ($H_3PO_4$), while nitric acid is reduced, typically to nitrogen dioxide ($NO_2$). $$P_4(s) + 20HNO_3(conc.) \rightarrow 4H_3PO_4(aq) + 20NO_2(g) + 4H_2O(l)$$ In this reaction, phosphorus's oxidation state changes from 0 to +5, and nitrogen's oxidation state changes from +5 in $HNO_3$ to +4 in $NO_2$. ### Batteries - **What is a Battery?** A battery (or more accurately, an electrochemical cell or galvanic cell) is a device that converts stored chemical energy directly into electrical energy through spontaneous redox (reduction-oxidation) reactions. It consists of one or more electrochemical cells, each comprising two half-cells: an anode (where oxidation occurs) and a cathode (where reduction occurs), separated by an electrolyte. - **Explanation of its Types:** Batteries are broadly classified into two main types based on their ability to be recharged: 1. **Primary Cells (or Primary Batteries):** * **Definition:** These are non-rechargeable batteries designed for single use. The chemical reactions that produce electrical energy are irreversible or practically irreversible. Once the reactants are consumed or the electrodes are significantly altered, the battery stops producing current and cannot be restored to its original state by recharging. * **Characteristics:** * Disposable after use. * Lower initial cost. * Often have a long shelf life. * Suitable for low-drain applications or where recharging is impractical. * **Examples:** * **Dry Cell (Leclanché Cell):** Common in flashlights and remote controls. It uses a carbon rod (cathode), a zinc container (anode), and an electrolyte paste of ammonium chloride, zinc chloride, and manganese dioxide. * **Alkaline Batteries:** Improved version of the dry cell, using potassium hydroxide as the electrolyte. Provides higher energy density and longer shelf life. * **Mercury Cells:** Historically used in small devices like hearing aids; now less common due to environmental concerns. 2. **Secondary Cells (or Secondary Batteries):** * **Definition:** These are rechargeable batteries. The chemical reactions that occur during discharge (producing electricity) can be reversed by applying an external electrical current, which forces the reaction to proceed in the opposite direction (charging). This allows the battery to be used multiple times. * **Characteristics:** * Rechargeable, allowing for multiple cycles of use. * Higher initial cost but lower cost per use over time. * Suitable for high-drain applications and where portability is key. * **Examples:** * **Lead-Acid Batteries:** Commonly used in automobiles. They consist of lead electrodes and a sulfuric acid electrolyte. * **Lithium-ion (Li-ion) Batteries:** Widely used in mobile phones, laptops, and electric vehicles due to their high energy density and lightweight nature. * **Nickel-Cadmium (Ni-Cd) Batteries:** Older rechargeable type, being phased out due to cadmium toxicity. * **Nickel-Metal Hydride (NiMH) Batteries:** Replaced Ni-Cd in many applications, offering higher capacity and less toxicity. - **What do you mean by Primary and Secondary Cell?** * **Primary Cell:** A primary cell is an electrochemical cell that converts chemical energy into electrical energy through an irreversible chemical reaction. Once the active chemicals are consumed, the cell cannot be recharged and must be discarded. It is a "use-and-throw" type of battery. * **Secondary Cell:** A secondary cell is an electrochemical cell in which the chemical reactions are reversible. This means that after discharge, electrical energy can be supplied from an external source to reverse the chemical reactions, restoring the cell to its charged state. It is a "rechargeable" type of battery. ### Emulsion - **What is an Emulsion?** An emulsion is a heterogeneous mixture of two or more immiscible liquids, where one liquid is dispersed in the other in the form of very fine droplets. For an emulsion to be stable, a third component called an emulsifying agent (or emulsifier) is usually required. The emulsifying agent forms an interfacial film between the dispersed droplets and the continuous phase, preventing the droplets from coalescing. - **How many types of its are there? Give an example of each type.** Emulsions are primarily classified into two main types based on which liquid forms the dispersed phase and which forms the continuous phase: 1. **Oil in Water (O/W) Emulsion:** * **Description:** In this type of emulsion, oil (or any non-polar liquid) is the dispersed phase, and water (or any polar liquid) is the continuous (dispersion) medium. The oil droplets are suspended in water. * **Characteristics:** These emulsions are typically diluted with water, conduct electricity, and feel less greasy. * **Emulsifying Agents:** Proteins, gums, natural soaps (e.g., sodium stearate), detergents. * **Example:** * **Milk:** Fat (oil) globules are dispersed in water. Casein, a protein, acts as the natural emulsifying agent. * **Vanishing Cream:** Oil droplets dispersed in water. * **Mayonnaise:** Oil dispersed in vinegar (mostly water), with egg yolk proteins acting as emulsifiers. 2. **Water in Oil (W/O) Emulsion:** * **Description:** In this type of emulsion, water (or any polar liquid) is the dispersed phase, and oil (or any non-polar liquid) is the continuous (dispersion) medium. The water droplets are suspended in oil. * **Characteristics:** These emulsions are typically diluted with oil, do not conduct electricity (or conduct poorly), and feel greasy. * **Emulsifying Agents:** Heavy metal salts of fatty acids, long-chain alcohols, lampblack. * **Example:** * **Butter:** Water droplets are dispersed in fat (oil). * **Cold Cream:** Water droplets dispersed in oil. * **Cod Liver Oil:** Water dispersed in oil. - **Additional Considerations:** * **Microemulsions:** These are thermodynamically stable, transparent mixtures of oil, water, and surfactant, often with a cosurfactant. Their droplet sizes are much smaller than traditional emulsions. * **Multiple Emulsions:** These are emulsions where the dispersed phase itself is an emulsion (e.g., W/O/W or O/W/O). ### Elevation of Boiling Point - **What do you mean by Elevation of Boiling Point?** The elevation of boiling point is a colligative property observed when a non-volatile solute is dissolved in a pure solvent. It refers to the phenomenon where the boiling point of the resulting solution ($T_b$) is higher than the boiling point of the pure solvent ($T_b^0$) at the same atmospheric pressure. The presence of non-volatile solute particles lowers the vapor pressure of the solvent, meaning that a higher temperature is required to reach the point where the solution's vapor pressure equals the external atmospheric pressure, thus causing the boiling point to elevate. The elevation in boiling point ($\Delta T_b$) is defined as: $$\Delta T_b = T_b - T_b^0$$ - **Obtain a Mathematical Expression to Determine the Molecular Weight of Solute with the Help of Elevation of Boiling Point:** The elevation of boiling point ($\Delta T_b$) is experimentally found to be directly proportional to the molal concentration (molality, $m$) of the solute in a dilute solution. $$\Delta T_b \propto m$$ Introducing a proportionality constant, we get: $$\Delta T_b = K_b \cdot m$$ Where: * $\Delta T_b$ = Elevation in boiling point (in °C or K) * $K_b$ = Ebullioscopic constant or molal elevation constant. It is a characteristic constant for a particular solvent (units: °C kg mol⁻¹ or K kg mol⁻¹). It represents the elevation in boiling point when 1 mole of a non-volatile solute is dissolved in 1 kg of the solvent. * $m$ = Molality of the solution (moles of solute per kilogram of solvent). Now, let's express molality in terms of the mass and molecular weight of the solute and solvent: Molality ($m$) is defined as: $$m = \frac{\text{Moles of solute (n}_2)}{\text{Mass of solvent in kg (W}_1)}$$ If $w_2$ is the mass of the solute in grams, and $M_2$ is the molar mass (molecular weight) of the solute in g/mol, then: $$\text{Moles of solute (n}_2) = \frac{w_2}{M_2}$$ If $w_1$ is the mass of the solvent in grams, then the mass of solvent in kilograms is $\frac{w_1}{1000}$. Substituting these into the molality expression: $$m = \frac{w_2/M_2}{w_1/1000} = \frac{w_2 \times 1000}{M_2 \times w_1}$$ Now, substitute this expression for $m$ into the boiling point elevation equation: $$\Delta T_b = K_b \cdot \frac{w_2 \times 1000}{M_2 \times w_1}$$ To determine the molecular weight of the solute ($M_2$), we can rearrange the equation: $$M_2 = \frac{K_b \times w_2 \times 1000}{\Delta T_b \times w_1}$$ This mathematical expression allows us to calculate the molar mass ($M_2$) of a non-volatile solute by experimentally measuring the elevation in boiling point ($\Delta T_b$) and knowing the masses of solute ($w_2$) and solvent ($w_1$), and the ebullioscopic constant ($K_b$) for the solvent. ### Osmotic Pressure - **Define Osmotic Pressure:** Osmotic pressure ($\Pi$) is defined as the minimum pressure that must be applied to a solution to prevent the inward flow of its pure solvent across a semipermeable membrane. Alternatively, it can be defined as the excess pressure that builds up on the solution side of a semipermeable membrane at equilibrium, due to the net movement of solvent molecules from the pure solvent side to the solution side (osmosis). Osmosis is the spontaneous net movement of solvent molecules through a selectively permeable membrane into a region of higher solute concentration, aiming to equalize solute concentrations on the two sides. - **Prove that Osmotic Pressure is a Colligative Property/Molecular Property:** Colligative properties are those properties of solutions that depend solely on the number of solute particles present in the solution, irrespective of their nature (identity). To prove that osmotic pressure is a colligative property, we need to show that its magnitude depends only on the concentration of the solute. For dilute solutions, the osmotic pressure ($\Pi$) can be experimentally found to follow a relationship similar to the ideal gas law, known as the **van't Hoff equation**: $$\Pi V = n_2 RT$$ Where: * $\Pi$ = Osmotic pressure (in atm or Pa) * $V$ = Volume of the solution (in L or m³) * $n_2$ = Number of moles of solute * $R$ = Ideal gas constant (e.g., 0.0821 L atm mol⁻¹ K⁻¹ or 8.314 J mol⁻¹ K⁻¹) * $T$ = Absolute temperature (in K) Rearranging the equation to solve for osmotic pressure: $$\Pi = \frac{n_2}{V} RT$$ We know that the term $\frac{n_2}{V}$ represents the molar concentration (molarity, $C$) of the solute in the solution (moles of solute per liter of solution). Therefore, the van't Hoff equation can be written as: $$\Pi = CRT$$ **Analysis for Colligative Property:** From the equation $\Pi = CRT$, we can observe the following: 1. **Dependence on Solute Amount:** The osmotic pressure ($\Pi$) is directly proportional to the molar concentration ($C$) of the solute. Molar concentration, in turn, is directly proportional to the number of moles of solute ($n_2$) present in a given volume of solution. 2. **Independence of Solute Identity:** The equation does not include any terms related to the specific chemical nature, size, or type of the solute particles (e.g., whether it's glucose, urea, or a different non-electrolyte). The gas constant $R$ and temperature $T$ are independent of the solute's identity. **Conclusion:** Since the osmotic pressure ($\Pi$) depends only on the number of moles of solute ($n_2$) dissolved in a given volume of solution ($V$) (i.e., its concentration) and is independent of the specific chemical identity of the solute, it is unequivocally a **colligative property**. ### Roasting & Calcination These are two important metallurgical processes used in the preliminary treatment of ores, typically before reduction to extract the metal. Both involve heating the ore, but under different conditions and for different purposes. - **Roasting:** * **Definition:** Roasting is the process of heating a concentrated ore strongly in the presence of excess air or oxygen, usually below its melting point. * **Purpose:** 1. **Conversion of Sulfide Ores to Oxides:** This is the primary purpose. Sulfide ores are very common, but it is easier to reduce metal oxides than metal sulfides. * **Example:** Zinc blende ($ZnS$) is roasted to convert it into zinc oxide ($ZnO$). $$2ZnS(s) + 3O_2(g) \xrightarrow{\Delta} 2ZnO(s) + 2SO_2(g)$$ * **Example:** Galena ($PbS$) to lead oxide ($PbO$). $$2PbS(s) + 3O_2(g) \xrightarrow{\Delta} 2PbO(s) + 2SO_2(g)$$ 2. **Removal of Volatile Impurities:** Impurities like arsenic, antimony, and sulfur are oxidized and volatilized as their oxides (e.g., $As_2O_3$, $Sb_2O_3$, $SO_2$). 3. **Making the ore porous:** This facilitates subsequent reduction processes. * **Conditions:** Requires the presence of air/oxygen. * **Products:** Typically metal oxides and gaseous oxides of sulfur or other volatile impurities. - **Calcination:** * **Definition:** Calcination is the process of heating an ore strongly in the absence or limited supply of air, usually below its melting point. * **Purpose:** 1. **Decomposition of Carbonate Ores:** Carbonate ores decompose to form metal oxides, releasing carbon dioxide. * **Example:** Limestone ($CaCO_3$) is calcined to produce quicklime ($CaO$). $$CaCO_3(s) \xrightarrow{\Delta} CaO(s) + CO_2(g)$$ * **Example:** Calamine ($ZnCO_3$) to zinc oxide ($ZnO$). $$ZnCO_3(s) \xrightarrow{\Delta} ZnO(s) + CO_2(g)$$ 2. **Decomposition of Hydrated Ores:** Hydrated ores lose their water of crystallization. * **Example:** Bauxite ($Al_2O_3 \cdot xH_2O$) loses water upon calcination. $$Al_2O_3 \cdot xH_2O(s) \xrightarrow{\Delta} Al_2O_3(s) + xH_2O(g)$$ 3. **Decomposition of Hydroxide Ores:** Hydroxide ores decompose to form metal oxides and water. * **Example:** Ferric hydroxide ($Fe(OH)_3$) to ferric oxide ($Fe_2O_3$). $$2Fe(OH)_3(s) \xrightarrow{\Delta} Fe_2O_3(s) + 3H_2O(g)$$ 4. **Removal of Volatile Organic Matter:** Any organic impurities present in the ore are driven off. * **Conditions:** Absence or limited supply of air/oxygen. * **Products:** Typically metal oxides and gaseous products like $CO_2$ or $H_2O$. ### Mineral & Ore These terms are fundamental in geology and metallurgy, referring to naturally occurring substances containing metals. While related, they are not interchangeable. - **Mineral:** * **Definition:** A mineral is a naturally occurring, inorganic solid with a definite chemical composition and a characteristic ordered atomic structure (crystalline structure). Minerals are the basic building blocks of rocks. * **Characteristics:** * **Naturally Occurring:** Formed by natural geological processes. * **Inorganic:** Generally not derived from living organisms (though some exceptions exist, like calcite in shells). * **Solid:** Exists in a solid state at normal temperatures and pressures. * **Definite Chemical Composition:** Can be expressed by a chemical formula (e.g., quartz is $SiO_2$, pyrite is $FeS_2$). * **Ordered Internal Structure:** Atoms are arranged in a specific, repeating pattern, giving rise to crystalline forms. * **Examples:** Quartz ($SiO_2$), Feldspar (e.g., $KAlSi_3O_8$), Calcite ($CaCO_3$), Mica, Bauxite ($Al_2O_3 \cdot 2H_2O$), Clay ($Al_2O_3 \cdot 2SiO_2 \cdot 2H_2O$). - **Ore:** * **Definition:** An ore is a naturally occurring rock or mineral from which one or more valuable metals (or other valuable materials) can be extracted **profitably and economically**. * **Characteristics:** * **Economic Viability:** The crucial factor for an ore is that it must be economically feasible to extract the desired metal. This depends on the concentration of the metal in the mineral, the cost of mining and processing, and the market value of the metal. * **Contains a Metal:** It must contain a metal in a sufficient quantity. * **Relationship to Minerals:** * **All ores are minerals, but not all minerals are ores.** For example: * **Bauxite** is a mineral that contains aluminum. Because aluminum can be extracted from bauxite profitably, bauxite is an **ore of aluminum**. * **Clay** is also a mineral that contains aluminum (aluminum silicates). However, extracting aluminum from clay is currently not economically viable compared to bauxite. Therefore, clay is a mineral of aluminum but **not an ore of aluminum**. * **Hematite** ($Fe_2O_3$) is a mineral and a primary ore of iron. * **Magnetite** ($Fe_3O_4$) is also a mineral and an ore of iron. * **Examples of Ores:** Hematite (iron), Bauxite (aluminum), Galena (lead), Cinnabar (mercury), Zinc blende (zinc), Chalcopyrite (copper). In summary, "mineral" is a scientific classification based on composition and structure, while "ore" is an economic classification based on the feasibility of metal extraction. ### Absorption & Adsorption These two terms describe processes where one substance takes up another, but they differ fundamentally in *where* the uptake occurs. - **Absorption:** * **Definition:** Absorption is a bulk phenomenon where a substance (the **absorbate**) is uniformly distributed throughout the entire volume or bulk of another substance (the **absorbent**). The particles of the absorbate penetrate deep into the material of the absorbent. * **Characteristics:** * **Bulk Phenomenon:** Occurs throughout the entire material. * **Uniform Distribution:** The absorbate is evenly distributed within the absorbent. * **Endothermic Process:** Often involves the breaking of bonds or creation of new phases, which can be energy-consuming. * **Slower Process:** Typically slower as it involves penetration into the bulk. * **No Concentration Gradient:** After saturation, the concentration of the absorbate is uniform within the absorbent. * **Examples:** 1. **Water vapor absorbed by anhydrous calcium chloride ($CaCl_2$):** The water molecules are taken up into the entire crystal lattice of $CaCl_2$. 2. **Sponge soaking up water:** Water fills the pores and spaces throughout the sponge. 3. **Ammonia gas absorbed in water:** Ammonia dissolves and spreads throughout the water volume to form ammonium hydroxide. 4. **Dye absorbed by a fabric:** The dye molecules penetrate the fibers of the fabric. - **Adsorption:** * **Definition:** Adsorption is a surface phenomenon where a substance (the **adsorbate**) accumulates or concentrates only on the surface of another substance (the **adsorbent**). The particles of the adsorbate do not penetrate into the bulk of the adsorbent. * **Characteristics:** * **Surface Phenomenon:** Occurs only at the interface between the two phases. * **Uneven Distribution:** Concentration of adsorbate is higher on the surface than in the bulk. * **Exothermic Process:** Usually releases energy (heat of adsorption) because it involves the formation of new bonds or interactions. * **Faster Process:** Initial adsorption is typically rapid as it only involves surface interaction. * **Equilibrium:** A dynamic equilibrium is established between adsorption and desorption. * **Examples:** 1. **Ammonia gas adsorbed on activated charcoal:** Ammonia molecules stick to the surface of charcoal. 2. **Silica gel adsorbing water vapor:** Water molecules condense and adhere to the surface of silica gel. 3. **Dye adsorbed by animal charcoal:** The dye molecules are selectively removed from a solution and concentrate on the charcoal surface. 4. **Gases (like $N_2$, $O_2$) adsorbed on metal surfaces:** Crucial in heterogeneous catalysis. In essence, absorption is like a sponge filling with water, while adsorption is like dust settling on a surface. ### Physical Adsorption & Chemical Adsorption Adsorption, a surface phenomenon, can be further classified into two main types based on the nature of the forces of attraction between the adsorbate and the adsorbent: Physical Adsorption (Physisorption) and Chemical Adsorption (Chemisorption). | Feature | Physical Adsorption (Physisorption) | Chemical Adsorption (Chemisorption) | | :-------------------------- | :---------------------------------------------------------------- | :---------------------------------------------------------------- | | **Nature of Forces** | Weak van der Waals forces (dispersion forces, dipole-dipole, etc.) exist between adsorbate and adsorbent. | Strong chemical bonds (covalent or ionic) are formed between adsorbate and adsorbent. | | **Heat of Adsorption** | Low (typically 20-40 kJ/mol). Similar to the enthalpy of liquefaction of the adsorbate. | High (typically 80-240 kJ/mol). Comparable to the enthalpy of chemical reactions. | | **Reversibility** | Highly reversible. Adsorbate can be easily removed by heating or reducing pressure. | Generally irreversible. Desorption requires much higher temperatures, often leading to decomposition of adsorbate. | | **Layers (Adsorbate Film)** | Multimolecular layer. Adsorbate molecules can stack up in multiple layers on the adsorbent surface. | Unimolecular layer. Adsorbate molecules form only a single layer on the adsorbent surface. | | **Specificity** | Not specific. Any gas can physisorb on any solid surface, provided conditions are favorable (low temp). | Highly specific. Chemisorption occurs only if there is a possibility of chemical bond formation between adsorbate and adsorbent. | | **Temperature Conditions** | Favored by low temperature. Adsorption decreases with increasing temperature. | Favored by high temperature (initially, to overcome activation energy), but eventually decreases at very high temperatures due to bond breaking. | | **Pressure Conditions** | Increases with increasing pressure. | Increases with pressure, but often saturates at high pressures. | | **Activation Energy** | Very low or negligible activation energy. | High activation energy is often required. | | **Effect of Surface Area** | Increases with increasing surface area of the adsorbent. | Increases with increasing surface area of the adsorbent. | | **Examples** | Adsorption of $N_2$ on mica at low temperature, adsorption of noble gases on charcoal. | Adsorption of $H_2$ on Ni, adsorption of $O_2$ on W, adsorption of $CO$ on Pt. | ### Schottky Defect & Frenkel Defect These are two primary types of point defects (imperfections) found in ionic crystals, which affect their physical and electrical properties. | Feature | Schottky Defect | Frenkel Defect | | :-------------------------- | :---------------------------------------------------------------- | :------------------------------------------------------------------ | | **Definition** | It is a type of vacancy defect where an equal number of cations and anions are missing from their regular lattice sites, creating vacancies. | It is a type of interstitial defect where an ion (usually a cation due to smaller size) leaves its normal lattice site and occupies an interstitial position within the crystal. | | **Cause** | Thermal vibrations causing ions to leave lattice sites, forming vacant pairs. | An ion displaces from its lattice site to an interstitial site, creating a vacancy and an interstitial defect. | | **Effect on Density** | Decreases the density of the crystal because ions are missing from the lattice, reducing the overall mass without a significant change in volume. | Does not change the density of the crystal because the ions are merely displaced within the crystal lattice; no ions are missing from the crystal. | | **Electrical Neutrality** | Maintained. An equal number of oppositely charged ions are missing, so the overall charge of the crystal remains neutral. | Maintained. The displaced ion retains its charge, and the vacancy has an equal but opposite charge, preserving neutrality. | | **Ionic Size Requirement** | Occurs predominantly in ionic compounds where the cation and anion are of nearly similar sizes. | Occurs predominantly in ionic compounds where there is a large difference in the sizes of the cation and anion (cations are typically much smaller). | | **Coordination Number** | Favored in compounds with high coordination numbers. | Favored in compounds with low coordination numbers. | | **Mobility** | Defects can move, contributing to electrical conductivity (ionic conduction by vacancy mechanism). | Ions can move through interstitial sites, contributing to electrical conductivity. | | **Examples** | NaCl, KCl, CsCl, AgBr (AgBr shows both Schottky and Frenkel defects). | AgCl, AgBr, AgI, ZnS. | ### Colligative Property - **What is a Colligative Property?** Colligative properties are a set of physical properties of solutions that depend solely on the number of solute particles present in a given amount of solvent, and not on the chemical identity or nature of these solute particles. The term "colligative" comes from the Latin word "colligatus," meaning "bound together," implying that these properties are bound to the collective number of particles. These properties are particularly important for dilute solutions of non-volatile solutes. - **Key Characteristics:** * **Dependence on Number of Particles:** The magnitude of a colligative property is directly proportional to the concentration of solute particles (e.g., moles, molecules, or ions) in the solution. * **Independence of Solute Identity:** The chemical nature (size, mass, charge, etc.) of the solute does not influence these properties, as long as the solute is non-volatile and does not associate or dissociate anomalously in the solvent. * **Effect on Solvent Properties:** Colligative properties are essentially changes in the solvent's properties caused by the presence of a solute. - **The Four Main Colligative Properties are:** 1. **Relative Lowering of Vapor Pressure (RLVP):** The vapor pressure of a solvent is lowered when a non-volatile solute is dissolved in it. The relative lowering of vapor pressure is proportional to the mole fraction of the solute. $$\frac{P^0 - P_s}{P^0} = X_2 = \frac{n_2}{n_1 + n_2}$$ Where $P^0$ is the vapor pressure of pure solvent, $P_s$ is the vapor pressure of solution, and $X_2$ is the mole fraction of solute. 2. **Elevation of Boiling Point ($\Delta T_b$):** The boiling point of a solvent increases when a non-volatile solute is added. This is because the lowered vapor pressure requires a higher temperature to reach the external atmospheric pressure. $$\Delta T_b = K_b \cdot m$$ Where $K_b$ is the molal elevation constant and $m$ is the molality of the solute. 3. **Depression of Freezing Point ($\Delta T_f$):** The freezing point of a solvent decreases when a non-volatile solute is added. The solute particles interfere with the formation of the ordered crystal structure of the solvent. $$\Delta T_f = K_f \cdot m$$ Where $K_f$ is the molal depression constant and $m$ is the molality of the solute. 4. **Osmotic Pressure ($\Pi$):** The pressure that must be applied to a solution to prevent the inward flow of solvent across a semipermeable membrane. It is directly proportional to the molar concentration of the solute. $$\Pi = CRT$$ Where $C$ is the molar concentration, $R$ is the gas constant, and $T$ is the absolute temperature. - **Importance:** Colligative properties are crucial for: * Determining the molar mass of unknown non-volatile solutes. * Understanding biological processes (e.g., cell membrane transport, kidney function). * Applications like antifreeze in car radiators (freezing point depression) or desalination of water (reverse osmosis). ### Lyophobic & Lyophilic Colloids Colloids are heterogeneous mixtures in which one substance is dispersed uniformly throughout another substance. The dispersed particles are larger than molecules but small enough not to settle out. Based on the affinity of the dispersed phase for the dispersion medium (solvent), colloids are broadly classified as lyophilic or lyophobic. The term "lyo" refers to the solvent; if the solvent is water, these terms become "hydrophilic" and "hydrophobic." | Feature | Lyophilic Colloids (Solvent-Loving) | Lyophobic Colloids (Solvent-Hating) | | :-------------------------- | :---------------------------------------------------------------- | :---------------------------------------------------------------- | | **Meaning** | "Lyo" (solvent) + "philic" (loving). Strong affinity between dispersed phase and dispersion medium. | "Lyo" (solvent) + "phobic" (hating). Little or no affinity between dispersed phase and dispersion medium. | | **Preparation** | Easily prepared by directly mixing the dispersed phase with the dispersion medium. The substances spontaneously form colloidal solutions. | Requires special methods for preparation, as they do not form spontaneously. Methods include dispersion (e.g., mechanical, electrical) and condensation (e.g., chemical reactions). | | **Stability** | Highly stable. The strong interaction (solvation) between the dispersed phase and the dispersion medium protects the particles from aggregation and coagulation. | Less stable. They are easily coagulated or precipitated by adding small amounts of electrolytes, heating, or shaking. Protection by charge and small solvation layer is minimal. | | **Reversibility** | Reversible. If the dispersion medium is removed (e.g., by evaporation), the dispersed phase can be redispersed by simply adding the dispersion medium back. | Irreversible. Once precipitated, the dispersed phase cannot be easily reformed into a colloidal solution by simply adding the dispersion medium. | | **Viscosity** | Viscosity of the colloidal solution is much higher than that of the pure dispersion medium. This is due to extensive solvation. | Viscosity of the colloidal solution is nearly the same as that of the pure dispersion medium. | | **Surface Tension** | Surface tension of the colloidal solution is generally lower than that of the pure dispersion medium. | Surface tension of the colloidal solution is nearly the same as that of the pure dispersion medium. | | **Charge on Particles** | Particles may or may not carry a charge. The stability is primarily due to solvation. | Particles invariably carry a charge (either positive or negative), which contributes to their stability by repelling each other. | | **Tendency to Coagulate** | Do not coagulate easily. Require large amounts of electrolyte to cause precipitation (salting out). | Easily coagulated by small amounts of electrolytes. | | **Examples** | Starch in water, Gum, Gelatin, Proteins, Rubber (in organic solvents). | Metal sols (e.g., Gold sol, Silver sol), Metal sulfides (e.g., $As_2S_3$ sol), Ferric hydroxide sol ($Fe(OH)_3$). | ### Methanoic Acid & Ethanoic Acid Methanoic acid (formic acid) and ethanoic acid (acetic acid) are the first two members of the carboxylic acid homologous series. While both are carboxylic acids, they exhibit distinct chemical properties due to their structural differences. | Feature | Methanoic Acid (Formic Acid) | Ethanoic Acid (Acetic Acid) | | :-------------------------- | :---------------------------------------------------------------- | :---------------------------------------------------------------- | | **Chemical Formula** | HCOOH | CH₃COOH | | **Structural Formula** | Contains both a carboxyl group (-COOH) and an aldehyde group (-CHO) within the same molecule. This unique structure gives it dual functionality. | Contains a carboxyl group (-COOH) and a methyl group (-CH₃). It lacks the aldehyde functional group. | | **IUPAC Name** | Methanoic Acid | Ethanoic Acid | | **Common Name** | Formic Acid | Acetic Acid | | **Reducing Property** | **Strong reducing agent.** Due to the presence of the -CHO group, it can be easily oxidized. It reduces Tollen's reagent (to silver mirror) and Fehling's solution (to red precipitate of $Cu_2O$). | **Does not act as a reducing agent.** It lacks the aldehyde group and cannot reduce Tollen's or Fehling's reagents. | | **Oxidation** | Easily oxidized to carbon dioxide ($CO_2$) and water ($H_2O$), often by mild oxidizing agents. | Not easily oxidized. Requires strong oxidizing agents and harsher conditions. | | **Acidity** | Stronger acid than ethanoic acid (pKa ≈ 3.75). The absence of an electron-donating alkyl group (like -CH₃) makes the carboxyl proton more acidic. | Weaker acid than methanoic acid (pKa ≈ 4.76). The electron-donating methyl group slightly destabilizes the carboxylate ion, making it a weaker acid. | | **Preparation Methods** | Industrially, by heating carbon monoxide (CO) with sodium hydroxide (NaOH) under pressure, followed by acidification with $H_2SO_4$. Also synthesized from methanol. | Industrially, by carbonylation of methanol ($CH_3OH + CO \rightarrow CH_3COOH$) or by oxidation of ethanol ($CH_3CH_2OH \rightarrow CH_3COOH$). Biologically, through fermentation. | | **Uses** | Used as a preservative, in leather tanning, textiles, and as a reducing agent in organic synthesis. | Used as vinegar (dilute solution), in the production of polymers (e.g., PVC), solvents, and pharmaceuticals. | ### Transition Elements (Complex Compounds & Coloured Compounds) Transition elements (d-block elements) are known for forming a wide variety of complex compounds and for their compounds often being highly colored. These properties arise primarily from their electronic configuration, particularly the presence of partially filled d-orbitals. - **Why do transition elements form complex compounds?** A complex compound (or coordination compound) consists of a central metal atom or ion (often a transition metal) bonded to a number of surrounding molecules or ions, called ligands, via coordinate covalent bonds (dative bonds). Transition metals exhibit a strong tendency to form such complexes due to the following reasons: 1. **Small Size and High Effective Nuclear Charge:** Transition metal ions typically have a relatively small size and a high positive charge (high effective nuclear charge). This allows them to strongly attract and polarize electron clouds of ligands, facilitating the formation of coordinate bonds. 2. **Presence of Vacant d-orbitals of Appropriate Energy:** Transition metal ions possess vacant (n-1)d orbitals that are available to accept electron pairs donated by ligands. Each ligand acts as a Lewis base, donating a lone pair of electrons to the central metal ion (Lewis acid), forming a coordinate covalent bond. The number of available d-orbitals allows for various coordination numbers and geometries. 3. **Variable Oxidation States:** Transition elements exhibit multiple oxidation states. This variability allows them to form complexes with different types of ligands and in different chemical environments. The oxidation state of the metal ion influences the number of vacant d-orbitals and their energy, impacting complex formation. 4. **High Tendency to Form Coordinate Bonds:** The combination of available vacant d-orbitals, small size, and high charge density makes transition metals excellent electron-pair acceptors, leading to strong coordinate bond formation with Lewis base ligands. - **Why do transition elements form coloured compounds?** The vibrant colors exhibited by most transition metal compounds, both in solid state and in solution, are primarily due to the presence of partially filled (n-1)d orbitals and the phenomenon of d-d electronic transitions. 1. **Partially Filled (n-1)d Orbitals:** Transition metal ions typically have incompletely filled d-orbitals (i.e., d¹ to d⁹ configurations). Ions with completely empty (d⁰, e.g., $Sc^{3+}$) or completely filled (d¹⁰, e.g., $Zn^{2+}$) d-orbitals are usually colorless because they cannot undergo d-d transitions. 2. **Ligand Field Splitting (Crystal Field Theory):** In an isolated transition metal atom or ion, all five d-orbitals are degenerate (have the same energy). However, when ligands approach the central metal ion to form a complex, the electrostatic field created by the ligands causes the degeneracy of the d-orbitals to be lifted. The d-orbitals split into two or more sets of different energy levels (e.g., $t_{2g}$ and $e_g$ in octahedral complexes). This energy difference ($\Delta_0$ or $\Delta_t$) typically corresponds to the energy of photons in the visible region of the electromagnetic spectrum. 3. **d-d Transitions:** When white light (which contains all colors of the visible spectrum) falls on a transition metal complex, electrons from the lower energy d-orbitals can absorb specific wavelengths (colors) of light. This absorbed energy promotes them to the higher energy d-orbitals. This process is called a **d-d transition**. 4. **Complementary Color:** The color that we perceive for the complex is not the color that was absorbed, but rather the **complementary color** (the color that is left over after the absorption). For example, if a complex absorbs red light, it will appear green; if it absorbs blue light, it will appear yellow. The specific color observed depends on the nature of the metal ion, its oxidation state, and the type of ligands surrounding it (which influence the magnitude of the ligand field splitting, $\Delta$). ### Power Alcohol & van't Hoff Factor - **Power Alcohol:** * **Definition:** Power alcohol refers to a mixture of petrol (gasoline) and ethanol (ethyl alcohol) used as a fuel for internal combustion engines. Typically, it contains about 5% to 25% ethanol, with the remainder being gasoline. Anhydrous (water-free) ethanol is preferred for blending with gasoline. * **Purpose and Advantages:** 1. **Renewable Resource:** Ethanol can be produced from biomass (e.g., corn, sugarcane, cellulosic materials) through fermentation, making it a renewable and potentially sustainable alternative to fossil fuels. 2. **Reduced Emissions:** Ethanol burns more cleanly than pure gasoline, producing fewer harmful pollutants such as carbon monoxide and unburnt hydrocarbons. It also reduces net carbon dioxide emissions if produced sustainably, as the CO₂ released during combustion is offset by the CO₂ absorbed by the plants during their growth. 3. **Higher Octane Rating:** Ethanol has a higher octane rating than gasoline, which can improve engine performance and reduce engine knocking. 4. **Oxygenate:** Ethanol acts as an oxygenate, improving the combustion efficiency of the fuel. * **Disadvantages:** 1. **Lower Energy Content:** Ethanol has a lower energy density than gasoline, meaning a vehicle running on power alcohol might travel a shorter distance per liter of fuel compared to pure gasoline. 2. **Corrosion:** Ethanol can be corrosive to certain engine components (plastics, rubber, metals) not designed for alcohol fuels. 3. **Hygroscopic Nature:** Ethanol is hygroscopic, meaning it absorbs water. Water contamination in fuel can lead to phase separation and engine problems. 4. **Food vs. Fuel Debate:** The use of food crops (like corn) for ethanol production can raise ethical concerns about food security and prices. - **van't Hoff Factor ($i$):** * **Definition:** The van't Hoff factor, denoted by $i$, is a dimensionless quantity that represents the ratio of the actual number of particles (ions or molecules) in solution after dissociation or association to the number of formula units initially dissolved in the solution. It quantifies the extent to which a solute dissociates or associates in a solvent, thereby influencing colligative properties. * **Mathematical Expression:** $$i = \frac{\text{Observed Colligative Property}}{\text{Calculated Colligative Property (assuming no dissociation/association)}}$$ Alternatively, it can be expressed as: $$i = \frac{\text{Total number of moles of particles after dissociation/association}}{\text{Number of moles of formula units initially dissolved}}$$ * **Significance:** Colligative properties (like boiling point elevation, freezing point depression, osmotic pressure, and relative lowering of vapor pressure) depend on the total number of solute particles. The van't Hoff factor modifies the colligative property equations to account for the actual number of particles present: * $\Delta T_b = i K_b m$ * $\Delta T_f = i K_f m$ * $\Pi = i CRT$ * **Values of $i$:** 1. **For Non-electrolytes:** For substances that do not dissociate or associate in solution (e.g., glucose, urea, sucrose), the van't Hoff factor $i = 1$. The number of particles in solution is the same as the number of molecules dissolved. 2. **For Electrolytes (Dissociation):** For substances that dissociate into ions in solution (e.g., salts, strong acids/bases), $i > 1$. The value of $i$ depends on the number of ions produced per formula unit and the degree of dissociation. * Example: For NaCl, $i \approx 2$ (dissociates into $Na^+$ and $Cl^-$). * Example: For $CaCl_2$, $i \approx 3$ (dissociates into $Ca^{2+}$ and $2Cl^-$). 3. **For Solutes that Associate:** For substances that associate (combine) in solution (e.g., ethanoic acid in benzene forming dimers), $i ### Double Salt & Complex Salt Both double salts and complex salts (or coordination compounds) are types of addition compounds, which are formed by the combination of two or more stable simple compounds. However, they differ significantly in their stability and behavior when dissolved in a solvent. | Feature | Double Salt | Complex Salt (Coordination Compound) | | :-------------------------- | :---------------------------------------------------------------- | :---------------------------------------------------------------- | | **Definition** | An addition compound that is stable only in the solid state. When dissolved in water, it completely dissociates into its constituent simple ions. | An addition compound that retains its identity and stability even when dissolved in water (or other solvent). It contains a complex ion which does not dissociate into its constituent simple ions. | | **Stability in Solution** | Loses its identity in solution. It breaks down into all the simple ions from which it was formed. | Retains its identity in solution. The complex ion remains intact, though the counter-ions may dissociate. | | **Tests for Ions** | Gives positive tests for all the simple ions from which it is constituted. For example, Mohr's salt gives tests for $Fe^{2+}$, $NH_4^+$, and $SO_4^{2-}$. | Does not give tests for all the simple ions present within the complex ion. For example, $K_4[Fe(CN)_6]$ gives tests for $K^+$ but not for $Fe^{2+}$ or $CN^-$. | | **Type of Bonding** | Contains only ionic bonds between the constituent simple ions. | Contains both ionic bonds (between the complex ion and counter-ions) and coordinate covalent bonds (between the central metal atom/ion and ligands within the complex ion). | | **Structure** | Lattice structure of two simple salts combined in stoichiometric ratio. | Central metal atom/ion surrounded by ligands in a specific geometry, forming a distinct complex ion. | | **Examples** | 1. **Mohr's Salt** (Ferrous ammonium sulfate): $FeSO_4 \cdot (NH_4)_2SO_4 \cdot 6H_2O$. In solution, it gives $Fe^{2+}$, $NH_4^+$, and $SO_4^{2-}$ ions. 2. **Carnallite:** $KCl \cdot MgCl_2 \cdot 6H_2O$. In solution, it gives $K^+$, $Mg^{2+}$, and $Cl^-$ ions. 3. **Potash Alum:** $K_2SO_4 \cdot Al_2(SO_4)_3 \cdot 24H_2O$. In solution, it gives $K^+$, $Al^{3+}$, and $SO_4^{2-}$ ions. | 1. **Potassium Ferrocyanide:** $K_4[Fe(CN)_6]$. In solution, it dissociates as $4K^+ + [Fe(CN)_6]^{4-}$. The complex ion $[Fe(CN)_6]^{4-}$ remains stable. 2. **Tetraamminecopper(II) Sulfate:** $[Cu(NH_3)_4]SO_4$. In solution, it dissociates as $[Cu(NH_3)_4]^{2+} + SO_4^{2-}$. The complex ion $[Cu(NH_3)_4]^{2+}$ remains stable. 3. **Silver Diammine Chloride:** $[Ag(NH_3)_2]Cl$. In solution, it dissociates as $[Ag(NH_3)_2]^+ + Cl^-$. | ### Ideal & Non-Ideal Solutions Solutions are mixtures of two or more components. Their behavior can be categorized as ideal or non-ideal, based on how closely they follow Raoult's law and the nature of interactions between their components. Raoult's law states that for a solution of volatile components, the partial vapor pressure of each component is equal to the vapor pressure of the pure component multiplied by its mole fraction in the solution. | Feature | Ideal Solutions | Non-Ideal Solutions | | :-------------------------- | :---------------------------------------------------------------- | :---------------------------------------------------------------- | | **Raoult's Law** | **Obeys Raoult's law** over the entire range of concentrations and temperatures. The partial vapor pressure of each component is directly proportional to its mole fraction. | **Does not obey Raoult's law.** Shows deviations from Raoult's law. | | **Intermolecular Forces (IMFs)** | The intermolecular forces of attraction between the solute-solvent molecules (A-B) are almost identical to those between solute-solute (A-A) and solvent-solvent (B-B) molecules. | The intermolecular forces of attraction between solute-solvent molecules (A-B) are either stronger or weaker than those between solute-solute (A-A) and solvent-solvent (B-B) molecules. | | **Enthalpy of Mixing ($\Delta H_{mix}$)** | **$\Delta H_{mix} = 0$**. No heat is absorbed or evolved when the components are mixed, as the energy required to break A-A and B-B interactions is equal to the energy released upon formation of A-B interactions. | **$\Delta H_{mix} \ne 0$**. **Positive Deviation:** $\Delta H_{mix} > 0$ (endothermic, heat absorbed). **Negative Deviation:** $\Delta H_{mix} **Positive Deviation:** $\Delta V_{mix} > 0$ (volume increases). **Negative Deviation:** $\Delta V_{mix} 1. Benzene and Toluene 2. n-Hexane and n-Heptane 3. Chloroethane and Bromoethane | **Positive Deviation:** A-B interactions are weaker than A-A and B-B. Components escape more easily, higher vapor pressure. 1. Ethanol and Acetone 2. Carbon disulfide and Acetone 3. Water and Ethanol **Negative Deviation:** A-B interactions are stronger than A-A and B-B. Components escape less easily, lower vapor pressure. 1. Chloroform and Acetone 2. Nitric acid and Water 3. Phenol and Aniline | | **Free Energy of Mixing ($\Delta G_{mix}$)** | $\Delta G_{mix} 0$ (increased randomness). | $\Delta S_{mix} > 0$ (increased randomness). | ### What is Soap? How does it work in cleaning clothes? - **What is Soap?** Soap is a salt of a long-chain fatty acid. Chemically, soaps are typically sodium or potassium salts of carboxylic acids with long aliphatic (hydrocarbon) chains, usually containing 12 to 18 carbon atoms. Common examples include sodium stearate ($C_{17}H_{35}COONa$), sodium palmitate ($C_{15}H_{31}COONa$), and sodium oleate ($C_{17}H_{33}COONa$). Soaps are produced by the saponification (hydrolysis) of fats and oils (triglycerides) with a strong base like sodium hydroxide (for hard soap) or potassium hydroxide (for soft soap). The structure of a soap molecule is amphiphilic, meaning it has two distinct parts: 1. **A long non-polar hydrocarbon tail:** This part is hydrophobic (water-hating) and lipophilic (oil-loving). It consists of a long chain of carbon and hydrogen atoms. 2. **A short polar ionic head:** This part is hydrophilic (water-loving) and lipophobic (oil-hating). It consists of the carboxylate group ($-COO^-$) and the metal cation (e.g., $Na^+$ or $K^+$). - **How does it work in cleaning clothes?** The cleaning action of soap is a result of its unique amphiphilic structure, which allows it to interact with both oil/grease (dirt) and water. The process involves several steps: 1. **Wetting:** When soap is added to water, it lowers the surface tension of the water. This allows the water to spread more easily and penetrate the fabric, wetting the dirt more effectively. 2. **Emulsification of Grease/Oil (Dirt):** * When clothes with oily or greasy dirt are put into soapy water, the non-polar hydrocarbon tails of the soap molecules orient themselves towards the oil/grease particles, as they are attracted to the non-polar dirt. * Simultaneously, the polar ionic heads of the soap molecules remain in contact with the surrounding water, as they are attracted to the polar water molecules. 3. **Micelle Formation:** As more soap molecules surround the oil/grease particle, they form a spherical structure called a **micelle**. In a micelle, the hydrophobic tails are embedded in the oil/grease droplet, while the hydrophilic heads form the outer surface, facing the water. 4. **Suspension and Repulsion:** * The micelle effectively encapsulates the dirt particle. Because the outer surface of the micelle is negatively charged (due to the carboxylate heads), the micelles repel each other. This prevents the dirt particles from clumping together and redepositing on the fabric. * The dirt-containing micelles are now soluble or suspended in water, forming a stable emulsion. 5. **Rinsing Away:** Mechanical agitation (scrubbing, washing machine action) helps to dislodge the dirt-laden micelles from the fabric. When the clothes are rinsed with fresh water, these suspended micelles, along with the encapsulated dirt, are washed away, leaving the fabric clean. In essence, soap acts as an emulsifying agent, breaking down large oil/grease stains into tiny droplets that can be dispersed and carried away by water. ### Principal Ores and Their Compositions for Several Important Metals Ores are naturally occurring rocks or minerals from which a metal can be extracted economically. Here is a table of principal ores and their compositions for some important metals: | Metal | Ore Name | Chemical Composition | Type of Ore | | :----------- | :------------------- | :--------------------------------------------------------- | :----------------- | | **Aluminum** | Bauxite | $Al_2O_3 \cdot xH_2O$ (Hydrated Aluminum Oxide) | Oxide | | | Corundum | $Al_2O_3$ (Aluminum Oxide) | Oxide | | **Iron** | Hematite | $Fe_2O_3$ (Ferric Oxide) | Oxide | | | Magnetite | $Fe_3O_4$ (Ferrous-Ferric Oxide) | Oxide | | | Siderite | $FeCO_3$ (Iron Carbonate) | Carbonate | | | Iron Pyrites (Fool's Gold) | $FeS_2$ (Iron(II) Disulfide) - *not an important ore due to S impurity* | Sulfide | | **Copper** | Copper Pyrites (Chalcopyrite) | $CuFeS_2$ (Copper Iron Sulfide) | Sulfide | | | Cuprite | $Cu_2O$ (Cuprous Oxide) | Oxide | | | Malachite | $CuCO_3 \cdot Cu(OH)_2$ (Basic Copper Carbonate) | Carbonate/Hydroxide| | | Chalcocite | $Cu_2S$ (Copper(I) Sulfide) | Sulfide | | **Zinc** | Zinc Blende (Sphalerite) | $ZnS$ (Zinc Sulfide) | Sulfide | | | Calamine (Smithsonite) | $ZnCO_3$ (Zinc Carbonate) | Carbonate | | | Zincite | $ZnO$ (Zinc Oxide) | Oxide | | **Lead** | Galena | $PbS$ (Lead Sulfide) | Sulfide | | | Cerussite | $PbCO_3$ (Lead Carbonate) | Carbonate | | | Anglesite | $PbSO_4$ (Lead Sulfate) | Sulfate | | **Tin** | Cassiterite | $SnO_2$ (Tin Dioxide) | Oxide | | **Gold** | Native Gold | $Au$ (often alloyed with Silver or Copper) | Native Metal | | | Calaverite | $AuTe_2$ (Gold Telluride) | Telluride | | **Silver** | Argentite (Silver Glance) | $Ag_2S$ (Silver Sulfide) | Sulfide | | | Horn Silver (Cerargyrite) | $AgCl$ (Silver Chloride) | Halide | | | Native Silver | $Ag$ | Native Metal | | **Mercury** | Cinnabar | $HgS$ (Mercury Sulfide) | Sulfide | | **Magnesium**| Magnesite | $MgCO_3$ (Magnesium Carbonate) | Carbonate | | | Dolomite | $CaMg(CO_3)_2$ (Calcium Magnesium Carbonate) | Carbonate | | | Carnallite | $KCl \cdot MgCl_2 \cdot 6H_2O$ (Potassium Magnesium Chloride Hexahydrate) | Chloride | ### Table of Particles of Unit Cell In crystallography, a unit cell is the smallest repeating unit that shows the full symmetry of a crystal lattice. The number of atoms or particles effectively belonging to a unit cell depends on their position within the cell (corner, face, edge, or body center), as different positions are shared with a different number of adjacent unit cells. | Unit Cell Type | Contribution from Particles at Corners | Contribution from Particles at Body Center | Contribution from Particles at Face Center | Total Number of Atoms (Z) per Unit Cell | | :------------------------- | :------------------------------------- | :----------------------------------------- | :----------------------------------------- | :-------------------------------------- | | **Simple Cubic (SC)** | $8 \text{ corners} \times \frac{1}{8} \text{ atom/corner} = 1$ | $0$ (no atom at body center) | $0$ (no atoms on faces) | **1** | | **Body-Centered Cubic (BCC)** | $8 \text{ corners} \times \frac{1}{8} \text{ atom/corner} = 1$ | $1 \text{ atom} \times 1 \text{ atom/body} = 1$ | $0$ (no atoms on faces) | **2** | | **Face-Centered Cubic (FCC)** | $8 \text{ corners} \times \frac{1}{8} \text{ atom/corner} = 1$ | $0$ (no atom at body center) | $6 \text{ faces} \times \frac{1}{2} \text{ atom/face} = 3$ | **4** | | **End-Centered Cubic (ECC)** | $8 \text{ corners} \times \frac{1}{8} \text{ atom/corner} = 1$ | $0$ | $2 \text{ opposite faces} \times \frac{1}{2} \text{ atom/face} = 1$ | **2** | **Explanation of Contributions:** * **Corner Particle:** An atom located at a corner of a cubic unit cell is shared by 8 adjacent unit cells. Therefore, only $\frac{1}{8}$ of that atom belongs to any single unit cell. * **Body-Centered Particle:** An atom located entirely within the body of a unit cell (not shared with any other cell) contributes 1 whole atom to that unit cell. * **Face-Centered Particle:** An atom located on the center of a face of a cubic unit cell is shared by 2 adjacent unit cells. Therefore, $\frac{1}{2}$ of that atom belongs to any single unit cell. * **Edge-Centered Particle:** (Not common in simple cubic systems, but relevant for other crystal systems). An atom on an edge is shared by 4 adjacent unit cells. Therefore, $\frac{1}{4}$ of that atom belongs to any single unit cell. The total number of atoms (Z) per unit cell is crucial for calculating the density of a crystalline solid and understanding its packing efficiency. ### Difference between Crystalline and Amorphous Solids Solids are characterized by their rigid structure and definite shape. They can be broadly classified into two main categories: crystalline solids and amorphous solids, based on the arrangement of their constituent particles (atoms, ions, or molecules). | Feature | Crystalline Solids | Amorphous Solids | | :-------------------------- | :---------------------------------------------------------------- | :---------------------------------------------------------------- | | **Arrangement of Particles**| **Ordered, long-range arrangement.** Particles are arranged in a definite, repeating three-dimensional pattern throughout the entire volume of the solid. This gives rise to a crystal lattice. | **Disordered, short-range arrangement (random).** Particles are arranged randomly, like in a liquid, but "frozen" in position. There might be some local order, but no long-range periodicity. | | **Melting Point** | **Sharp and definite melting point.** They melt abruptly at a specific temperature, as all bonds are broken simultaneously due to uniform arrangement. | **Melt over a range of temperature.** They gradually soften over a range of temperatures, becoming increasingly fluid before they melt. This is because their bonds are not uniformly strong. | | **Cleavage Properties** | When cut with a sharp-edged tool, they cleave into two pieces with **smooth, planar surfaces** and definite angles. | When cut with a sharp-edged tool, they break into two pieces with **irregular, rough surfaces**. | | **Anisotropy/Isotropy** | **Anisotropic.** Their physical properties (e.g., electrical conductivity, thermal conductivity, refractive index, mechanical strength) are different when measured along different directions within the crystal. This is due to the ordered arrangement of particles. | **Isotropic.** Their physical properties are the same in all directions because of the random arrangement of particles. The properties average out in all directions. | | **Nature** | **True solids.** They have a definite shape and volume and exhibit characteristic properties of solids. | **Pseudo-solids or supercooled liquids.** They have a tendency to flow, albeit very slowly, over long periods (e.g., old glass panes are thicker at the bottom). They lack the true solid-state characteristics of long-range order. | | **Heat of Fusion** | **Definite heat of fusion.** They have a characteristic and fixed amount of heat absorbed during melting. | **No definite heat of fusion.** They do not have a sharp melting point, so there is no specific heat of fusion. | | **Cooling Curve** | Shows a sharp break (plateau) at the melting point, indicating a phase transition at a constant temperature. | Shows a continuous change in slope (no sharp break) as they gradually solidify over a temperature range. | | **Examples** | Sodium chloride (NaCl), Quartz ($SiO_2$), Diamond, Sugar, Ice, all metals. | Glass, Rubber, Plastics (polyethylene, PVC), Tar, Amorphous silicon. |