Introduction
Solutions are homogeneous mixtures composed of two or more substances. In the field of chemistry, understanding solutions is crucial as they are involved in many chemical reactions and processes. This study note will cover the key concepts related to solutions, including types of solutions, concentration terms, colligative properties, and various laws governing solutions.
Types of Solutions
Based on the Physical State of Solvent and Solute
- Solid Solutions
- Example: Alloys like brass (copper and zinc)
- Liquid Solutions
- Example: Salt in water, alcohol in water
- Gaseous Solutions
- Example: Air (mixture of gases)
Based on the Nature of Solvent
- Aqueous Solutions: Solvent is water.
- Example: NaCl in water
- Non-Aqueous Solutions: Solvent is not water.
- Example: Iodine in alcohol (tincture of iodine)
Water is the most common solvent and is often referred to as the "universal solvent" due to its ability to dissolve a wide range of substances.
Concentration Terms
Molarity (M)
Molarity is defined as the number of moles of solute per liter of solution.
$$ M = \frac{n}{V} $$
Where:
- $M$ = Molarity
- $n$ = Number of moles of solute
- $V$ = Volume of solution in liters
To prepare 1 M NaCl solution, dissolve 58.44 g of NaCl (1 mole) in enough water to make 1 liter of solution.
Molality (m)
Molality is defined as the number of moles of solute per kilogram of solvent.
$$ m = \frac{n}{m_{solvent}} $$
Where:
- $m$ = Molality
- $n$ = Number of moles of solute
- $m_{solvent}$ = Mass of solvent in kilograms
Normality (N)
Normality is the number of gram equivalents of solute per liter of solution.
$$ N = \frac{n_{eq}}{V} $$
Where:
- $N$ = Normality
- $n_{eq}$ = Number of gram equivalents of solute
- $V$ = Volume of solution in liters
Mole Fraction ($X$)
Mole fraction is the ratio of the number of moles of a component to the total number of moles of all components in the solution.
$$ X_A = \frac{n_A}{n_{total}} $$
Where:
- $X_A$ = Mole fraction of component A
- $n_A$ = Number of moles of component A
- $n_{total}$ = Total number of moles of all components
Parts Per Million (ppm)
Parts per million is used to express very dilute concentrations of substances.
$$ ppm = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 10^6 $$
TipFor very dilute solutions, 1 ppm is approximately equal to 1 mg/L.
Solubility
Factors Affecting Solubility
- Nature of Solute and Solvent
- "Like dissolves like": Polar solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents.
- Temperature
- Solubility of solids generally increases with temperature.
- Solubility of gases decreases with temperature.
- Pressure
- Solubility of gases increases with pressure (Henry's Law).
Henry's Law
Henry’s Law states that the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid.
$$ C = k_H \cdot P $$
Where:
- $C$ = Concentration of gas
- $k_H$ = Henry's Law constant
- $P$ = Partial pressure of gas
Students often confuse molarity and molality. Remember, molarity is volume-based, while molality is mass-based.
Colligative Properties
Colligative properties depend on the number of solute particles in a solution, not on their nature.
Relative Lowering of Vapor Pressure
$$ \frac{\Delta P}{P^0} = X_B $$
Where:
- $\Delta P$ = Lowering of vapor pressure
- $P^0$ = Vapor pressure of pure solvent
- $X_B$ = Mole fraction of solute
Elevation of Boiling Point
$$ \Delta T_b = K_b \cdot m $$
Where:
- $\Delta T_b$ = Elevation in boiling point
- $K_b$ = Ebullioscopic constant
- $m$ = Molality
Depression of Freezing Point
$$ \Delta T_f = K_f \cdot m $$
Where:
- $\Delta T_f$ = Depression in freezing point
- $K_f$ = Cryoscopic constant
- $m$ = Molality
Osmotic Pressure
$$ \Pi = MRT $$
Where:
- $\Pi$ = Osmotic pressure
- $M$ = Molarity
- $R$ = Gas constant
- $T$ = Temperature in Kelvin
Calculate the osmotic pressure of a 0.1 M NaCl solution at 298 K. Assume $R = 0.0821 , \text{L atm K}^{-1} \text{mol}^{-1}$.
$$ \Pi = (0.1) \times (0.0821) \times (298) = 2.45 , \text{atm} $$
Raoult's Law
Raoult's Law states that the partial vapor pressure of each volatile component in a solution is equal to the vapor pressure of the pure component multiplied by its mole fraction in the solution.
$$ P_A = X_A \cdot P_A^0 $$
Where:
- $P_A$ = Partial vapor pressure of component A
- $X_A$ = Mole fraction of component A
- $P_A^0$ = Vapor pressure of pure component A
Raoult's Law is applicable to ideal solutions where the intermolecular forces between unlike molecules are similar to those between like molecules.
Conclusion
Understanding solutions and their properties is fundamental in chemistry. This study note has covered the essential concepts, including types of solutions, concentration terms, factors affecting solubility, colligative properties, and Raoult's Law. Mastery of these concepts is crucial for solving problems related to solutions in the NEET Chemistry syllabus.
