Solutions Class 12 Chemistry Chapter 1 One Shot | New NCERT CBSE | Rationalised syllabus topics

A solution is a homogeneous mixture of two or more components. A homogeneous mixture has a uniform composition and properties throughout. Examples include salt dissolved in water, air, tea, coffee, and alloys like brass (copper and zinc) and bronze (copper and tin). Solutions can exist in solid, liquid, or gaseous states. A solution typically consists of a solvent (the component present in the largest quantity, determining the physical state of the solution) and one or more solutes (the components present in lesser quantities). A binary solution contains only one solute and one solvent.

Solutions are categorized into three types based on the physical state of the solvent: gaseous solutions, liquid solutions, and solid solutions. Gaseous solutions have a gas as the solvent, examples include solid in gas (camphor in nitrogen), liquid in gas (chloroform in nitrogen), and gas in gas (air). Liquid solutions have a liquid as the solvent. Examples include solid in liquid (sugar in water), liquid in liquid (ethanol in water), and gas in liquid (oxygen dissolved in water, crucial for aquatic life). Solid solutions have a solid as the solvent. Examples include solid in solid (copper in gold alloy), liquid in solid (mercury amalgam with sodium), and gas in solid (hydrogen in palladium).

Concentration describes the amount of solute and solvent in a solution. It can be expressed qualitatively (e.g., dilute or concentrated) or quantitatively (using specific numerical values). Quantitative methods are crucial for precise measurements and include mass percentage, volume percentage, mass by volume percentage, parts per million (PPM), mole fraction, molarity, and molality.

Mass percentage (w/w) calculates the mass of a component per 100 parts of the solution by mass. Volume percentage (v/v) calculates the volume of a component per 100 parts of the solution by volume, often used for liquid-liquid solutions. Mass by volume percentage (w/v) indicates the mass of a component present in 100 ml of the solution, commonly used in medicine and pharmacy. These methods provide quantitative ways to express how much of a component is present in a solution.

Parts per million (PPM) is used when a component is present in very small, trace quantities, expressing the number of parts of the component per million parts of the solution. Mole fraction (χ) expresses the ratio of the number of moles of one component to the total number of moles of all components in the solution. For a binary solution with components A and B, χA + χB = 1. A detailed example of calculating the mole fraction of ethylene glycol in a solution is provided.

Molarity (M) defines the number of moles of solute present per liter of solution. It is denoted by 'M' and its unit is molar (mol/L). Molarity is calculated as (moles of solute) / (volume of solution in liters). A numerical example demonstrates how to calculate molarity for a given solution of NaOH.

Molality (m) defines the number of moles of solute per kilogram of solvent. Unlike molarity, molality is independent of temperature because it involves mass, not volume. It is denoted by 'm' or 'moles per kg'. A numerical example illustrates the calculation of molality for a solution of ethanoic acid in benzene.

Molarity is temperature-dependent because volume changes with temperature, while molality is temperature-independent as it relies on mass. This makes molality a preferred concentration unit in some scientific contexts, particularly where temperature fluctuations are significant. Practical problems involving mass percentage, molarity, and mole fraction calculations are solved.

Solubility is the maximum amount of a substance that can be dissolved in a given amount of solvent at a specific temperature and pressure. The principle 'like dissolves like' explains that polar substances dissolve in polar solvents, and non-polar substances dissolve in non-polar solvents due to similar molecular interactions. Factors affecting solubility include the nature of the solute and solvent, temperature, and pressure.

When a solid dissolves in a liquid, the process is called dissolution, which increases the solution's concentration. Simultaneously, crystallization occurs, where solute particles aggregate and separate from the solution. A dynamic equilibrium is established when the rates of dissolution and crystallization become equal. A solution is saturated when no more solute can dissolve at a given temperature and pressure. Unsaturated solutions can still dissolve more solute. According to Le Chatelier's Principle, temperature affects solubility: if dissolution is endothermic, solubility increases with temperature; if exothermic, solubility decreases with temperature.

The solubility of gases in liquids is significant, as seen in aquatic life (oxygen in water) and carbonated drinks (CO2 in water). Gas solubility increases with increasing pressure due to the highly compressible nature of gases. Henry's Law states that the partial pressure of a gas above a liquid is directly proportional to its mole fraction in the solution (P = Khχ). A higher Henry's constant (Kh) indicates lower solubility. Practical applications include why aquatic life prefers cold water (higher oxygen solubility) and the effervescence in soda water (CO2 release upon pressure drop). Scuba diving, where divers breathe pressurized air, illustrates how high pressure increases nitrogen solubility in blood, leading to 'bends' upon ascent due to bubble formation—a problem mitigated by diluting breathing gas with helium. A numerical problem on Henry's Law is demonstrated.

Vapor pressure is the pressure exerted by the vapor of a liquid in a closed container at equilibrium. It depends on the volatility of the liquid; volatile liquids readily vaporize. Raoult's Law states that for a solution of volatile liquids, the partial vapor pressure of each component is directly proportional to its mole fraction in the liquid phase (P1 = P1°χ1). P1° is the vapor pressure of the pure component. Dalton's Law of Partial Pressures states that the total vapor pressure of the solution is the sum of the partial pressures of its components (Ptotal = P1 + P2). Raoult's Law can be seen as a special case of Henry's Law.

When a non-volatile solute is added to a pure solvent, the vapor pressure of the solution decreases. This is because the non-volatile solute occupies a portion of the solvent's surface, reducing the number of solvent molecules that can escape into the vapor phase. The lowering of vapor pressure is independent of whether the solute is volatile or non-volatile, as the solvent's surface area available for vaporization is reduced in either case. This leads to a relative lowering of vapor pressure.

Ideal solutions obey Raoult's Law across all concentrations and temperatures. For ideal solutions, the enthalpy of mixing (ΔHmix) and volume of mixing (ΔVmix) are zero. Perfect ideal solutions are rare, but examples include bromobenzene and chloroethane solutions. Non-ideal solutions do not obey Raoult's Law and exhibit deviations. Positive deviation occurs when the vapor pressure is higher than expected (weaker solute-solvent interactions), like ethanol and acetone. Negative deviation occurs when the vapor pressure is lower than expected (stronger solute-solvent interactions), like phenol and aniline. Azeotropes are special liquid mixtures that boil at a constant temperature and have the same composition in liquid and vapor phases (minimum boiling azeotropes for positive deviation, maximum boiling azeotropes for negative deviation).

Colligative properties depend only on the number of solute particles, not their nature. The relative lowering of vapor pressure is a colligative property. According to Raoult's Law, the relative lowering of vapor pressure (ΔP1/P1°) is equal to the mole fraction of the solute (χ2). For dilute solutions, this can be approximated as n2/n1 (moles of solute / moles of solvent). A numerical example demonstrates calculating the mass of a non-volatile solute based on the relative lowering of vapor pressure.

Elevation in boiling point (ΔTb) is the increase in a solution's boiling point compared to that of a pure solvent. A liquid boils when its vapor pressure equals the atmospheric pressure. When a non-volatile solute is added, the vapor pressure of the solvent decreases, requiring a higher temperature to reach the atmospheric pressure, thus elevating the boiling point. ΔTb is directly proportional to the molality (m) of the solution, expressed as ΔTb = Kb * m, where Kb is the ebullioscopic constant (or molal elevation constant). A numerical example is provided to calculate the molar mass of a solute using boiling point elevation.

Depression in freezing point (ΔTf) is the decrease in a solution's freezing point compared to that of a pure solvent. A substance freezes when its vapor pressure in the liquid phase equals its vapor pressure in the solid phase. Adding a non-volatile solute lowers the solvent's vapor pressure, which means a lower temperature is needed for the solution's vapor pressure to match that of the solid solvent, hence depressing the freezing point. ΔTf is directly proportional to the molality (m) of the solution, expressed as ΔTf = Kf * m, where Kf is the cryoscopic constant (or molal depression constant). An interesting real-world example is how sea water does not freeze at 0°C due to dissolved salts. A numerical problem demonstrates calculating the molar mass of a solute using freezing point depression.

Osmosis is the spontaneous net movement of solvent molecules through a selectively permeable membrane from a region of higher solvent concentration to a region of lower solvent concentration. This movement continues until equilibrium is reached. Osmotic pressure (π) is the minimum pressure that must be applied to the solution side to prevent the net flow of solvent across the semi-permeable membrane. Osmotic pressure is given by π = CRT, where C is the molarity, R is the gas constant, and T is the temperature. This property is advantageous for determining molar masses, especially for macromolecules, as measurements are typically done at room temperature.

Isotonic solutions have the same osmotic pressure as another solution (e.g., blood plasma). When placed in an isotonic solution, cells neither swell nor shrink because there is no net movement of solvent. Hypertonic solutions have a higher osmotic pressure (higher solute concentration) than another solution, causing cells to shrink as solvent moves out. Hypotonic solutions have a lower osmotic pressure (lower solute concentration), causing cells to swell as solvent moves in. A practical example is the use of intravenous saline solutions (0.9% NaCl) in medicine to avoid adverse effects on blood cells.

Reverse osmosis (RO) occurs when an external pressure greater than the osmotic pressure is applied to the solution side of a semi-permeable membrane. This forces the solvent molecules to move from the solution side (higher solute concentration) to the pure solvent side (lower solute concentration), against the natural direction of osmosis. RO is widely used for desalination of seawater and water purification in homes (RO water purifiers), effectively removing salts and impurities.

Abnormal molar mass refers to situations where the experimentally determined molar mass of a solute differs from its theoretically expected value. This often happens due to association (solute molecules combine, like acetic acid forming dimers in benzene) or dissociation (solute molecules break apart) in the solution. The Van't Hoff factor (i) is introduced to account for these deviations. It is defined as the ratio of the observed colligative property to the calculated colligative property, or the ratio of the normal molar mass to the abnormal molar mass. The Van't Hoff factor helps to correctly predict colligative properties in solutions where association or dissociation occurs. The expressions for colligative properties are modified by multiplying them with the Van't Hoff factor (i) to account for these abnormalities.