𝐁𝐢𝐡𝐚𝐫 𝐁𝐨𝐚𝐫𝐝 𝐂𝐥𝐚𝐬𝐬 𝟏𝟐 𝐂𝐡𝐞𝐦𝐢𝐬𝐭𝐫𝐲 𝐂𝐡 - 𝟐 𝐒𝐨𝐥𝐮𝐭𝐢𝐨𝐧 𝐎𝐧𝐞 𝐒𝐡𝐨𝐭 | 𝐎𝐛𝐣𝐞𝐜𝐭𝐢𝐯𝐞 + 𝐒𝐮𝐛𝐣𝐞𝐜𝐭𝐢𝐯𝐞 + 𝐏𝐘𝐐𝐬 𝐢𝐧 𝟏 𝐕𝐢𝐝𝐞𝐨

Class 12 Chemistry - Solution Chapter One Shot - PW Bihar Board English Medium - Notes

1. Summary

This video is a one-shot revision of the Class 12 Chemistry Solution chapter, specifically tailored for Bihar Board students. It covers objective and subjective questions, important previous year questions (PYQs), short notes for exam preparation, and is based on the 2026 board pattern. The video aims to help students achieve high scores in chemistry by providing a comprehensive understanding of the chapter.

2. Key Takeaways

* **Focus:** The video comprehensively covers the Solution chapter, including theory (A-Z), subjective questions, and tips for success.

* **Content Coverage:** The video includes objective questions (MCQs), subjective questions, important PYQs, exam-oriented short notes, and explanations based on the 2026 board pattern.

* **Target Audience:** It's specifically designed for Class 12 Science students in Bihar Board and is useful for other state boards as well.

* **Key Concepts:** The video emphasizes concepts like homogeneous and heterogeneous mixtures, different types of solutions (binary, etc.), concentration terms (mass %, volume %, molarity, molality, ppm, mole fraction), Raoult's Law, and colligative properties.

3. Detailed Notes

**I. Introduction and Overview (0:00:00 - 0:00:48)**

* The video begins with an introduction, audio and video check.

* The topic of the lecture is "Solutions," chapter 2 in the physical chemistry of class 12 chemistry.

* It aims to solve any doubts related to the chapter, particularly for Bihar Board students.

**II. Solutions and Homogeneous Mixtures (0:00:48 - 0:01:10)**

* **Definition of Solution:** A homogeneous mixture of two or more components.

* The emphasis on homogeneous mixtures, differentiating them from heterogeneous mixtures.

**III. Homogeneous vs. Heterogeneous Mixtures: Examples (0:01:10 - 0:01:42)**

* The video emphasizes that in solutions, we focus on homogeneous mixtures.

* Examples are given: Water with salt is a homogeneous mixture while water and oil will form heterogeneous mixtures.

**IV. Properties of Homogeneous and Heterogeneous Mixtures (0:01:42 - 0:02:25)**

* **Homogeneous Mixture Characteristics:** Uniform composition, single phase, components not visible.

* **Heterogeneous Mixture Characteristics:** Non-uniform composition, multiple phases, components often visible.

**V. Solution Components (0:02:25 - 0:02:40)**

* Solutions always need two components: A solvent and a solute

* **Solvent:** Present in major amount and is the one that "dissolves" other components.

* **Solute:** Present in minor amount and gets "dissolved" by the solvent.

* Example: Salt is the solute and water is the solvent in a saltwater solution.

* If the amount of two compounds are the same, then the one with the state of the solution is considered the solvent.

**VI. Types of Solutions (0:02:40 - 0:03:36)**

* **Binary Solution:** Consists of two components.

* **Ternary Solution:** Consists of three components.

* **Quaternary Solution:** Consists of four components.

* The focus is on binary solutions.

**VII. Solution Types Based on Physical State (0:03:36 - 0:04:11)**

* The video discusses different types of solutions based on the physical state of the solvent and provides examples:

* Gaseous solution

* Liquid solution

* Solid solution

* Emphasis that the phase of the solution is determined by the phase of the solvent.

**VIII. Concentration Terms (0:04:11 - 0:05:29)**

* **Purpose of Concentration Terms:** To quantify the amount of solute dissolved in a specific amount of solvent.

* **Mass Percentage:** Ratio of mass of solute to the mass of the solution (both in grams), multiplied by 100.

* **Volume Percentage:** Ratio of volume of solute to the volume of the solution (both in milliliters), multiplied by 100.

* **Molarity (M):** Number of moles of solute dissolved per liter of solution. (moles of solute/liter of solution)

* **Moles:** Calculated by the ratio of mass by molar mass.

* **Molality (m):** Number of moles of solute dissolved per kilogram of solvent. (moles of solute/kg of solvent)

* The difference between molality and molarity.

**IX. Question and Example (0:05:29 - 0:07:45)**

* Example:

* Mass percentage of sugar in sugar syrup given:

* Total weight. 234.2 g

* Weight of sugar: 34.2 g

* Solve the molarity.

* Molarity: moles of solute/liters of solution.

* Convert all units to one unit to solve the problem.

**X. Molarity vs. Molality (0:07:45 - 0:08:24)**

* **Key Differences**

* Defintion

* Unit

* The third difference is that is the value is not dependent on temperature.

**XI. Henry's Law (0:08:24 - 0:13:00)**

* **Henry's Law:** The solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid surface at a constant temperature.

* **Henry's Law Equation:** P = K * X. Where P is the pressure, K is Henry’s constant, and X is the mole fraction.

* **Implication:** The higher the value of KH, the lower the solubility of the gas.

**XII. Applications of Henry's Law (0:13:00 - 0:15:39)**

* **Soft Drinks:** CO2 is dissolved under high pressure to increase its solubility, and the bottles are sealed.

* **Scuba Diving:** To prevent "bends" or decompression sickness, divers use a mixture of gases (oxygen, nitrogen, and helium). Nitrogen's (N2) low solubility and the choice of helium gas.

* * If you are to go to higher altitudes, then oxygen is reduced because of the reduced pressure in the atmospheric. This causes the lack of oxygen in the body causing Anoxia.

**XIII. Factors Affecting Solubility (0:15:39 - 0:17:12)**

* **Nature of Solute and Solvent (0:15:39 - 0:16:08):** "Like dissolves like" (Polar solutes dissolve in polar solvents; nonpolar solutes in nonpolar solvents)

* **Temperature (0:16:08 - 0:16:18):** For solids in liquids, generally, solubility increases with increasing temperature.

* **Pressure (0:16:18 - 0:17:12):** The solubility of gases in liquids increases with increasing pressure (Henry's Law).

**XIV. Solution Formation Process (0:17:12 - 0:17:40)**

* **Dissolution:** The process where a solid solute dissolves in a liquid solvent, breaking it down.

* **Crystallization:** The reverse process, where solute particles come together to form a solid (occurs in a saturated solution).

* **Equilibrium:** Equilibrium is reached when the rate of dissolution equals the rate of crystallization.

**XV. Saturated and Unsaturated Solutions (0:17:40 - 0:18:19)**

* **Saturated Solution:** Contains the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature.

* **Unsaturated Solution:** Can dissolve more solute.

**XVI. Henry's Law (Further Explanation) (0:18:19 - 0:19:07)**

* **Relationship to Raoult's Law:** Henry's law is a special case of Raoult's law.

* P = KH * X: Relationship between pressure and solubility. P - partial pressure, KH - Henry's Constant, X - Mole Fraction.

* Higher KH means lower solubility.

* Example: CO2 in soft drinks (High pressure, high solubility)

**XVII. Henry's Law: Applications (0:19:07 - 0:20:01)**

* **Applications**

* Soft drinks (carbonated beverages): CO2 is dissolved under high pressure.

* Scuba diving

* Mountain sickness (altitude): Low oxygen pressure.

* Divers and mountains at high altitudes where there is a low presence of atmospheric pressure

* A higher atmosphere pressure results in a greater solubility.

**XVIII. Limitations of Henry's Law (0:20:01 - 0:20:24)**

* High pressure

* Very low temperatures

* Gases that react with the solvent

* Only valid for dilute solutions.

**XIX. Raoult's Law (0:20:24 - 0:20:34)**

* The main thing to note about Raoult’s Law is that for an ideal solution, Raoult's law applies.

* Raoult’s law is the foundation behind many of the colligative properties.

**XX. Raoult's Law: (Further Discussion) (0:20:34 - 0:21:42)**

* **Non Ideal Solutions:** (Liquid solute and solvent mix at any proportion). The properties deviate from Raoult's Law.

**XXI. Types of Non-Ideal Solutions (0:21:42 - 0:22:18)**

* There are two types of Non-Ideal Solutions:

* Positive Deviations

* If P total is greater than what’s expected

* Negative Deviations

* If P total is less than what’s expected

**XXII. Colligative Properties (0:22:18 - 0:22:27)**

* **Colligative Properties Definition:** Properties of a solution that depend on the number of solute particles in the solution, not their nature.

* **Four Colligative Properties:**

* Lowering of vapor pressure

* Elevation of boiling point

* Depression of freezing point

* Osmotic pressure.

**XXIII. RLVP (0:22:27 - 0:22:38)**

* **RLVP (Relative Lowering of Vapor Pressure):** The vapor pressure of a solution is always lower than the vapor pressure of the pure solvent.

* **Equation:** P0 - P / P0 = n2 / n1 + n2

**XXIV. Lowering of Vapor Pressure and Examples (0:22:38 - 0:23:08)**

* **Why does this happen?** (When a non-volatile solute is added)

* Example: The vapor pressure will decrease as you add solid solutes to make the solution.

* Solute molecules on the surface of the liquid occupy the surface area, reducing solvent evaporation.

**XXV. Elevation of Boiling Point (0:23:08 - 0:23:16)**

* **Definition:** The boiling point of a solution is higher than the boiling point of the pure solvent.

* **Why?** Because the vapor pressure of the solution is less than the vapor pressure of pure solvent.

**XXVI. Elevation of Boiling Point. Example (0:23:16 - 0:23:27)**

* As the heat increases, so does boiling

* Heat is released. As a result, this causes the boiling point for each compound to go to a different boiling point

* To find the boiling point you will need to use the formula of delta tb = kb *m.

* * Kb is boiling point constant * M is Molality.

**XXVII. Depression in Freezing Point (0:23:27 - 0:23:38)**

* **Definition:** The freezing point of a solution is lower than the freezing point of the pure solvent.

* Why? Because when you add a solid, the solute decreases because it is not soluble.

**XXVIII. Osmosis (0:23:38 - 0:23:46)**

* **Definition:** The spontaneous movement of solvent molecules across a semipermeable membrane from a region of low solute concentration to a region of higher solute concentration.

**XXIX. Osmotic Pressure (0:23:46 - 0:23:53)**

* **Definition:** The external pressure required to stop the flow of the solvent across the semipermeable membrane.

**XXX. Formula and Applications of Osmotic Pressure (0:23:53 - 0:24:01)**

* **Osmotic Pressure Formula:** π = CRT (where π is osmotic pressure, C is the concentration, R is the gas constant, and T is temperature in Kelvin.)

* Examples such as shrinking cells and or raisins

**XXXI. Van't Hoff Factor (0:24:01 - 0:25:36)**

* **Van't Hoff Factor (i):** A factor to account for the dissociation or association of solute particles.

* **Definition:** The ratio of the total number of moles of particles after association or dissociation to the initial number of moles of solute.

**XXXII. Applications of Van't Hoff Factor (0:25:36 - 0:25:48)**

* **For electrolytes:**

* Strong electrolytes: (e.g., NaCl -> Na+ + Cl-): i = 2 (because of the dissociation, breaking down to 2 moles).

* For a compound of Na2SO4: i = 3.

**XXXIII. Putting it all Together (0:25:48 - 0:26:02)**

* Use of i in the formulas

* Pi = iCRT

* Delta Tb = iKbm

* Delta Tf = iKfm

* P = iP0x2

**XXXIV. Tips and Tricks (0:26:02 - 0:26:59)**

* The presenter stresses understanding over rote memorization.

* Focus on the relationships and the "why" behind the concepts.

* Review the key definitions and concepts for the exam.

* Work through practice questions to reinforce the concepts.

**XXXV. Summary and Next Steps (0:27:00 - 0:27:54)**

* The presenter summarizes the content covered.

* Students are encouraged to review the material and attempt the problems

* Next lecture will be focused on questions and problems.