Chemical Thermodynamics MCQs

Matter can leave or enter the system, so the system is not isolated or closed.

Energy can also be transferred as heat or work across the boundary.

Therefore a system that exchanges both matter and energy with surroundings is an Open system.

Work done by a gas is w = ∫ p_ext dV, where p_ext is the pressure exerted by the surroundings.

In a free expansion the gas expands against a vacuum, so the external pressure is zero throughout the process.

Therefore the integral gives zero work, so the work done is Zero.

First law: ΔU = q + w.

At constant volume, the system does no expansion work, so w = 0.

Hence the heat absorbed q equals ΔU, the change in internal energy. deltaU

Specific heat (c) is the heat needed to raise 1 g of a substance by 1 °C.

For water, c = 4.18 J g

ΔH° = [–393.5 + 2(–285.8)] – (–74.8) = –965.1 + 74.8 = **–890 kJ mol⁻¹** (combustion of CH₄).

Enthalpy change for a reaction depends only on initial and final states, not on the path taken.

By constructing a hypothetical series of steps whose individual ΔH values are known, we can sum them to obtain the overall ΔH.

This allows us to find ΔH for reactions that are hard to measure in a single experiment.

Thus Hess’s law is mainly used to calculate enthalpy changes that are difficult to measure directly.

Work = Pₑₓₜ ΔV = 2 atm × (5 L – 1 L) = 8 L·atm

8 L·atm × 101.3 J / L·atm ≈ 8.1 × 10² J

So the gas does about **8 L·atm (≈ 8.1 × 10² J)** of work.

At constant pressure the heat transferred to a system is defined as the heat of reaction, qₚ.

From the definition of enthalpy, H = U + pV, its differential at constant pressure gives dH = dU + p dV = δqₚ.

Thus the heat absorbed equals the change in enthalpy, i.e., ΔH. Change in enthalpy.

In a bomb calorimeter the reacting mixture is sealed in a rigid container, so its volume does not change during the reaction.

Because the volume is fixed, the heat measured corresponds to the change in internal energy (ΔU) at constant volume.

Thus the heat of reaction is recorded under constant‑volume conditions. Volume

Enthalpy is a state function, meaning the total change in enthalpy depends solely on the initial state (reactants) and final state (products), independent of the reaction pathway (Hess's Law).

To transform reactants into products:

1-Decomposing the reactants into their constituent standard elements requires the reverse of formation

2- Synthesizing the products from those standard elements releases or absorbs their enthalpy of formation

Adding both paths together yields products minus reactants.

ΔH for a reaction is the heat change when reactants convert to products.

Reversing the reaction changes the sign of the enthalpy change because the heat absorbed becomes heat released (and vice‑versa).

Therefore ΔH for B → A = –20 kJ.

Entropy S = k ln W, where W is the number of accessible micro‑states.

In a gas molecules move freely in translation, rotation and vibration, giving the largest W compared to liquids (restricted movement) and solids (vibrations about fixed positions).

Hence the entropy is highest for the gas.

A system that can exchange matter but not energy with surroundings is a closed system; one that can exchange both is open.

If neither matter nor energy can cross the boundary, the system is isolated.

Hence the ideal thermos flask is an isolated system.

Internal energy (U) of a system is defined as the total microscopic energy possessed by its particles.

It includes the translational, rotational and vibrational kinetic energies of the molecules as well as the intermolecular potential energies.

Therefore it is the kinetic plus potential energy of constituents.

For a reaction at constant pressure, enthalpy change ΔH = ΔU + Δ(PV).

For ideal gases PV = nRT, so the change in PV equals Δ(n_g)·R·T.

Substituting gives ΔH = ΔU + Δ(n_g)RT.

Hence ΔH = ΔU + Δ(ng)RT.

q = m c ΔT

Substitute m = 100 g, c = 4.18 J g⁻¹ °C⁻¹, ΔT = 10 °C → q = 100 × 4.18 × 10

q = 4180 J, which matches the given answer. 4180 J

Enthalpy of formation is for H₂(g)+½O₂(g)→H₂O(l) and is –286 kJ per mole of water formed.

The decomposition reaction is the exact reverse: H₂O(l)→H₂(g)+½O₂(g).

Reversing a reaction changes the sign of ΔH while the magnitude stays the same.

Therefore the enthalpy change for decomposition is +286 kJ/mol. +286 kJ/mol

Entropy (S) is defined by the statistical relation S = k ln W, where W is the number of possible microstates of the system.

A larger number of microstates means the particles can be arranged in many ways, indicating higher randomness.

Therefore entropy measures the randomness or disorder of a system. Randomness or disorder of a system

A sealed steel container is rigid, so no mass can pass through its walls.

Heat can be transferred through the steel by conduction or radiation.

Thus only energy (as heat) can be exchanged while matter remains confined.

First law states that the energy of an isolated system is constant, so any change in its internal energy (ΔU) must be accounted for by energy transferred as heat (q) and as work (w).

Using the sign convention of chemistry, heat absorbed by the system is taken as positive (q > 0) and work done on the system is also positive (w > 0).

Hence the deltaU = q + w