Thermodynamics and Gas Laws

Boyle's Law states that for a fixed amount of gas at constant temperature, pressure and volume are inversely proportional (PV = constant). As pressure increases, volume decreases proportionally, and vice versa. This fundamental gas law was formulated by Robert Boyle in 1662.

Van der Waals forces (intermolecular attractions and repulsions) cause real gases to deviate from ideal gas behavior, especially at high pressure and low temperature. Ideal gas assumptions neglect molecular size and intermolecular forces.

Boyle's Law (1662) states that for a fixed amount of ideal gas at constant temperature, pressure and volume are inversely proportional: P₁V₁ = P₂V₂, or PV = constant. Option B confuses volume with temperature, while C incorrectly specifies real gases.

The ideal gas law relates pressure (P), volume (V), amount of gas in moles (n), the universal gas constant (R), and temperature (T). PV = nRT is the correct and complete form.

Boyle's law states that for an ideal gas at constant temperature, the pressure and volume are inversely proportional: P₁V₁ = P₂V₂. This means when pressure increases, volume decreases proportionally, and vice versa, as long as temperature remains unchanged. Option A correctly represents this relationship.

Charles's law states that for an ideal gas at constant pressure, volume is directly proportional to absolute temperature: V ∝ T, or V₁/T₁ = V₂/T₂. Cross-multiplying gives V₁T₂ = V₂T₁, which is option A. This means that as temperature increases, volume increases proportionally, and vice versa, when pressure is held constant.

An ideal gas perfectly follows Charles's law (V ∝ T at constant pressure) and Boyle's law (PV = constant at constant temperature). These two laws, along with Gay-Lussac's law, are fundamental gas laws that describe the behavior of ideal gases. Option B correctly identifies Charles's law and Boyle's law as essential for ideal gas behavior.

The ideal gas law is PV = nRT, where P is pressure, V is volume, n is moles of gas, R is the gas constant, and T is temperature. This equation combines Boyle's law, Charles's law, and Avogadro's law into one unified relationship. Option B shows the correct ideal gas law equation.

Dispersion forces (London forces) are weak attractive forces between molecules caused by temporary dipoles, not repulsive forces. Additionally, ideal gases by definition have no intermolecular forces at all. The statement is wrong on both counts.

Hund's rule deals with electron configuration in atoms - it states that electrons will occupy degenerate orbitals singly before pairing up. It has nothing to do with gas pressures, which are covered by Dalton's law of partial pressures.

At the microscopic level, gas molecules collide randomly with container walls causing local pressure variations, and temperature directly relates to the average kinetic energy (velocity) of molecules. Both statements are physically accurate.

 K.E. of the electron is 5 eV = 5 x 1.6 x 10^-19 J But K.E. = (3/2)KT ∴ 5 x 1.6 x 10^-19 = (3/2)(1.38 x 10^-23) x T T =  2 x 5 x 1.6 x 10^-19/3 x 1.38 x 10^-23   = 38.65  x 10³ K