This MCQ module is based on: Thermodynamic Terms
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Thermodynamic Terms
🎓 Class 11
Chemistry
CBSE
Theory
Ch 5 – Thermodynamics
⏱ ~14 min
🌐 Language: [gtranslate]
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Thermodynamic Terms, System and Surroundings
5.1 Thermodynamics — An Introduction
Chemical reactions are accompanied by energy changes. Some release heat (a candle burning), some absorb heat (ice melting), some release light, and some store chemical energy. Thermodynamics is the branch of science that deals with the energetics of physical and chemical processes. It tells us:
- How much energy is exchanged when a reaction occurs (1st law)
- Whether the reaction will occur spontaneously (2nd law)
- How efficient the energy conversion is
Limitation: Thermodynamics tells us if a process can occur and how much energy is exchanged — but NOT how fast it occurs (that is chemical kinetics) and NOT what happens at the molecular level.
5.2 The System and the Surroundings
To study energy changes, we divide the universe into two parts:
System: The part of the universe under study, separated from the rest by a real or imaginary boundary.
Surroundings: Everything outside the system that can exchange energy or matter with it.
Universe = System + Surroundings.
Surroundings: Everything outside the system that can exchange energy or matter with it.
Universe = System + Surroundings.
For example, when we study a chemical reaction in a beaker, the contents of the beaker form the system, and the beaker, air, table, room — all are surroundings.
5.2.1 Types of Systems
Based on the nature of the boundary, systems are classified into three types:
| Type | Exchange of matter | Exchange of energy | Example |
|---|---|---|---|
| Open | Yes | Yes | Boiling water in an open pan |
| Closed | No | Yes | Water in a sealed metal flask |
| Isolated | No | No | Hot tea in a perfect thermos flask (idealised) |
5.3 The State of the System & State Functions
The state of a system is described by giving the values of measurable properties called state variables: pressure (P), volume (V), temperature (T), amount (n), composition.
State function (state property): A property whose value depends only on the present state of the system, and not on how that state was reached. Examples: P, V, T, internal energy U, enthalpy H, entropy S, Gibbs energy G.
Path functions: Quantities whose values depend on the route taken — e.g., heat (q) and work (w).
Path functions: Quantities whose values depend on the route taken — e.g., heat (q) and work (w).
Mountaineering analogy: your altitude at the summit is a state property — same regardless of which trail you took. But the distance walked depends on the route. Altitude ≡ state function; distance walked ≡ path function.
5.4 The Internal Energy (U)
Every substance possesses internal energy U — the sum of all kinetic and potential energies of its constituent molecules: translational, rotational, vibrational, electronic, nuclear and intermolecular interactions.
U is a state function. Only changes ΔU = U_final − U_initial can be measured, never the absolute value.
Internal energy can change by two distinct routes:
- By exchange of heat (q) — energy transfer due to a temperature difference
- By exchange of work (w) — energy transfer through any other mechanism (e.g., pushing a piston, electrical work)
5.4.1 Sign Conventions (IUPAC)
| Quantity | Positive (+) | Negative (−) |
|---|---|---|
| Heat q | Absorbed BY the system (endothermic) | Released BY the system (exothermic) |
| Work w | Work done ON the system (compression) | Work done BY the system (expansion) |
| ΔU | Energy increases | Energy decreases |
Memory aid: "Heat in is positive, work in is positive." Anything added to the system (whether heat or work) raises its internal energy.
5.5 Thermodynamic Processes
A process is the way a system changes from one state to another. Five common types:
| Process | What is constant | Mathematical condition | Example |
|---|---|---|---|
| Isothermal | Temperature (T) | ΔT = 0 | Slow expansion of gas in contact with a thermostat |
| Adiabatic | No heat exchange | q = 0 | Air rushing out of a punctured tyre (rapid) |
| Isobaric | Pressure (P) | ΔP = 0 | Open-vessel reactions (atmospheric pressure) |
| Isochoric | Volume (V) | ΔV = 0 | Reaction in a sealed bomb calorimeter |
| Cyclic | Returns to initial state | ΔU = 0 for one cycle | Refrigerator, heat-engine cycle |
5.5.1 Reversible vs Irreversible Processes
A reversible process proceeds infinitesimally slowly, with the system always in equilibrium with the surroundings. An irreversible process proceeds at finite speed and the system passes through non-equilibrium states (e.g., gas suddenly expanding into vacuum).
Real processes are always irreversible (finite-speed, dissipation present). Reversible processes are useful idealisations giving the maximum work output (or minimum work input).
🎯 Interactive: Identify the System Type
Pick a real-world example. The simulator tells you whether it is open, closed or isolated, and whether matter and/or energy can cross the boundary.
Type: Open
Matter exchange: Yes | Energy exchange: Yes
Steam (matter) escapes; heat from flame enters and rises into air.
🧪 Activity 5.1 — Path Dependence of Heat & Work
Setup: 1 mole of an ideal gas expands from (P₁=2 atm, V₁=10 L) to (P₂=1 atm, V₂=20 L). It does so by two paths:
Path A: Free expansion into vacuum (irreversible, no external pressure)
Path B: Reversible isothermal expansion at T = 244 K
Predict: How does the work done by the gas compare in the two paths? Is internal energy change ΔU the same?
Path A: P_ext = 0 → w = −P_ext·ΔV = 0 (no work done)
Path B: w_rev = −nRT ln(V₂/V₁) = −(1)(8.314)(244)(ln 2) = −1406 J ≈ −1.4 kJ
ΔU is the same in both paths (state function!) — for isothermal ideal gas, ΔU = 0. So q must adjust: q_A = 0, q_B = +1406 J.
Lesson: Heat and work are path functions; their values depend on the route. Internal energy ΔU is a state function and is route-independent.
Worked Example 5.1: Sign Conventions
A system absorbs 200 J of heat and does 80 J of work on the surroundings. Calculate ΔU.
Apply IUPAC convention:
Heat absorbed BY system: q = +200 J
Work done BY system (energy lost): w = −80 J
ΔU = q + w = 200 + (−80) = +120 J
The system's internal energy rose by 120 J.
Heat absorbed BY system: q = +200 J
Work done BY system (energy lost): w = −80 J
ΔU = q + w = 200 + (−80) = +120 J
The system's internal energy rose by 120 J.
Worked Example 5.2: Identify the Process
A gas in a sealed steel cylinder is heated. Identify the type of process and which thermodynamic variables are zero.
A sealed steel cylinder has fixed volume → isochoric process.
ΔV = 0 → w = −P_ext·ΔV = 0.
By 1st law: ΔU = q + w = q (all the heat goes into raising U; none is lost as work). This is why bomb calorimeters measure q_v = ΔU directly.
ΔV = 0 → w = −P_ext·ΔV = 0.
By 1st law: ΔU = q + w = q (all the heat goes into raising U; none is lost as work). This is why bomb calorimeters measure q_v = ΔU directly.
🎯 Competency-Based Questions
Q1. Which of the following is NOT a state function? L1 Remember
Answer: (c) Heat. Heat (q) and work (w) are path functions — their values depend on the route, not just the initial and final states.
Q2. A pressure cooker on a stove with the safety valve closed exemplifies which type of system? L2 Understand
Answer: A closed system. The walls prevent the steam (matter) from escaping but heat still flows from the stove (energy can cross). It is NOT isolated because energy is still exchanged.
Q3. State whether True or False: "ΔU depends on the path taken between initial and final states." L2 Understand
FALSE. Internal energy U is a state function — ΔU depends only on the initial and final states, not the path. However, q and w (which sum to ΔU) ARE path-dependent.
Q4. Predict the type of process: a balloon being slowly compressed inside a thermally insulated room. L3 Apply
Answer: Insulation prevents heat transfer → q = 0 → adiabatic process. The work done in compression goes entirely into raising the internal energy: ΔU = w. The gas heats up — this is why bicycle pumps get hot during rapid pumping!
Q5. HOT: Two identical gases are taken from state A to state B by two different paths. Path 1 is reversible and does 200 J of work; Path 2 is irreversible and does 150 J of work. Compare ΔU and the heat absorbed in each. L5 Evaluate
Answer: ΔU is identical in both paths (state function!). Let ΔU = X.
Path 1: w₁ = −200 J → q₁ = X − w₁ = X + 200 J.
Path 2: w₂ = −150 J → q₂ = X + 150 J.
So q₁ > q₂: the reversible path absorbs MORE heat. This is consistent with the principle that reversible processes give maximum work and absorb maximum heat for the same ΔU.
Path 1: w₁ = −200 J → q₁ = X − w₁ = X + 200 J.
Path 2: w₂ = −150 J → q₂ = X + 150 J.
So q₁ > q₂: the reversible path absorbs MORE heat. This is consistent with the principle that reversible processes give maximum work and absorb maximum heat for the same ΔU.
🧠 Assertion–Reason Questions
Choose: (A) Both true, R explains A. (B) Both true, R doesn't explain A. (C) A true, R false. (D) A false, R true.
A: Internal energy is a state function.
R: The change in internal energy depends only on the initial and final states of a system.
Answer: (A). Both true; R explains A. Path-independence is the defining characteristic of a state function.
A: Heat absorbed by a system is given a positive sign.
R: Heat raises the internal energy of the system, which is a positive change.
Answer: (A). Both true; R explains A. The IUPAC sign convention is: anything added to the system raises its internal energy and is taken positive.
A: A perfectly isolated system can exchange neither matter nor energy with the surroundings.
R: Isolated systems exist abundantly in nature.
Answer: (C). A is TRUE (definition of isolated). R is FALSE — perfect isolation is an idealisation; real systems always leak some energy.
Frequently Asked Questions — Thermodynamic Terms, System and Surroundings
What is a thermodynamic system?
A thermodynamic system is the specific part of the universe chosen for study in a thermodynamic analysis. Everything outside the system is called the surroundings, and the imaginary or real surface separating them is the boundary. In NCERT Class 11 Chemistry Chapter 5, systems are classified as open (exchange both matter and energy with surroundings, like a cup of hot coffee), closed (exchange only energy, like a sealed flask of gas), or isolated (exchange neither, like a perfect thermos flask). The choice of system shapes the analysis — the system is what we apply the laws of thermodynamics to.
What is the difference between open, closed and isolated systems?
An open system exchanges both matter and energy with its surroundings — example: water boiling in an open kettle, the human body, an open chemistry beaker. A closed system exchanges only energy (heat or work) but not matter — example: gas in a sealed cylinder, a closed reaction flask, water in a closed pot. An isolated system exchanges neither matter nor energy — example: a perfect thermos flask or the universe as a whole. NCERT Class 11 Chemistry Chapter 5 uses this classification to set up thermodynamic problems correctly, since the equations depend on the type of system.
What is a state function in thermodynamics?
A state function is a property of a thermodynamic system whose value depends only on the current state of the system (defined by P, V, T, n) and not on the path by which the system reached that state. Examples include internal energy (U), enthalpy (H), entropy (S) and Gibbs free energy (G). In contrast, heat (q) and work (w) are path functions — their values depend on the route taken. NCERT Class 11 Chemistry uses this distinction extensively: state-function changes (ΔU, ΔH) can be calculated using Hess's law regardless of mechanism.
What are intensive and extensive properties?
Intensive properties do not depend on the amount or size of the system — examples include temperature, pressure, density, viscosity, refractive index and concentration. Extensive properties depend on the amount of substance — examples include mass, volume, internal energy, enthalpy, entropy and heat capacity. The ratio of two extensive properties (e.g., molar volume = V/n, or density = m/V) gives an intensive property. NCERT Class 11 Chemistry Chapter 5 uses this classification to identify which properties of a system are size-independent for thermodynamic calculations.
What is thermodynamic equilibrium?
A system is in thermodynamic equilibrium when all its macroscopic properties (P, V, T, composition) remain constant with time and there is no net flow of matter or energy. Three conditions must be satisfied simultaneously: (1) thermal equilibrium — uniform temperature, no heat flow; (2) mechanical equilibrium — uniform pressure, no work done; (3) chemical equilibrium — no net chemical reaction or mass transfer. NCERT Class 11 Chemistry Chapter 5 uses thermodynamic equilibrium as the starting point for state-function analysis and as the condition for measuring P, V and T accurately.
Why is the universe considered an isolated system?
The universe is conventionally treated as an isolated system in thermodynamics because, by definition, nothing exists outside it that it can exchange matter or energy with. This idealisation allows the second law of thermodynamics to be stated as 'the entropy of the universe always increases for any real process.' For any system + its surroundings = universe: ΔS_total ≥ 0. NCERT Class 11 Chemistry Chapter 5 uses this framework to define spontaneity. The universe-as-isolated-system idea also underlies the first law (energy of the universe is constant).
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