Chemistry · Thermodynamics · NEET
Internal energy is the SUM of every energy inside the system at once: the kinetic energy of moving molecules, the potential energy in every chemical bond, the energy of electrons in their orbits, and even the energy stored in the nucleus. To get the absolute value of U you would have to count all of these exactly for every particle. We have no way to do that. So U has no fixed, knowable absolute value. This matters for NEET because every thermodynamics formula uses ΔU, never a raw U value.
We use the first law: ΔU = q + w. We put the system in a sealed, insulated (adiabatic) container so no heat escapes, meaning q = 0. Then ΔU = w, the work done on the system, which we CAN measure. In a bomb calorimeter the volume is also fixed, so no pressure-volume work is done and the heat we measure equals ΔU. We only ever need the difference between two states, so absolute U is not needed.
Internal energy is a STATE FUNCTION. Its value depends only on the current state of the system (temperature, pressure, amount), not on the path used to reach that state. That is why ΔU between two states is always the same no matter which route you take. Heat (q) and work (w) alone are path functions, but their sum q + w = ΔU is path-independent. NEET loves to test this: ΔU is fixed, q and w can differ for different paths.
Absolute internal energy would mean the exact total value of U measured from a true zero. That true zero cannot be reached or defined, so the absolute value is unknown. When your textbook just says 'internal energy', it usually means the value relative to some reference or, more often, the CHANGE ΔU. In NEET problems you never plug in an absolute U; you always work with ΔU.
The Greek letter Δ (delta) means 'change in'. We write ΔU because only the change has a real, measurable meaning. U_final and U_initial each have unknown absolute values, but their difference U_final − U_initial is measurable and is what all NEET formulas (ΔU = q + w, ΔH = ΔU + Δn_g RT) use.
Yes, exactly. Enthalpy is defined as H = U + pV. Since U has no known absolute value, H also has no known absolute value. That is why we only ever use ΔH (change in enthalpy) too. This is why NEET gives you enthalpy of reaction, formation, combustion etc. as CHANGES, never as one fixed number.
At a certain temperature T (K), during a process, 500 J is absorbed by the system and work of 200 J is done by the system. Then the change in internal energy of the system is:
A gas is allowed to expand in a well-insulated container against a constant external pressure of 2.5 atm from an initial volume of 2.50 L to a final volume of 4.50 L. The change in internal energy ΔU of the gas in joules will be:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
No. It is not a limit of instruments. Internal energy adds up every form of energy down to the nucleus, and there is no defined true zero of energy for matter, so an absolute value simply does not exist to measure.
The first law of thermodynamics: ΔU = q + w, where q is heat added to the system and w is work done on the system. In an insulated (adiabatic) system q = 0, so ΔU equals the measurable work.
Because internal energy is a state function. Its value depends only on the state, so the difference between two fixed states is fixed. Heat and work can each differ by path, but their sum ΔU cannot.
Yes. H = U + pV, and since U has no absolute value, neither does H. We always use ΔH, the change in enthalpy, in NEET thermochemistry.
Internal energy is extensive — it depends on the amount of substance. Double the moles and you double U (and ΔU). Temperature and pressure, by contrast, are intensive.
