Internal pressure is a thermodynamic quantity that describes the change in a system's internal energy with respect to its volume at constant temperature, mathematically defined as \pi_T = \left( \frac{\partial U}{\partial V} \right)_T, where U is the internal energy, V is the volume, and T is the temperature.[1] This property has units of pressure and provides a direct measure of the strength and nature of intermolecular forces within the substance, being positive when attractive forces predominate over repulsive ones.[1] For an ideal gas, where intermolecular interactions are absent, the internal pressure is zero.[1]A fundamental thermodynamic relation connects internal pressure to the equation of state of the system: \pi_T = T \left( \frac{\partial P}{\partial T} \right)_V - P, where P is the pressure.[1] This expression enables the calculation of \pi_T from experimental data on pressure, temperature, and volume, without directly measuring energy changes.[1] In real gases and liquids, internal pressure arises from the potential energy associated with molecular attractions, contrasting with the kinetic contributions to pressure in ideal systems.[1]For condensed phases like liquids, internal pressure can be approximated using empirical coefficients: \pi_T \approx \frac{\alpha}{\kappa_T}, where \alpha = \frac{1}{V} \left( \frac{\partial V}{\partial T} \right)_P is the thermal expansion coefficient and \kappa_T = -\frac{1}{V} \left( \frac{\partial V}{\partial P} \right)_T is the isothermal compressibility.[1] Typical values for liquids are on the order of thousands of atmospheres, far exceeding ordinary external pressures, which underscores the significant role of cohesive forces in maintaining liquid structure.[1][2] This approximation highlights internal pressure's utility in correlating thermodynamic behavior with molecular properties in non-ideal fluids.[1]Internal pressure is essential for analyzing deviations from ideal behavior in equations of state, such as the van der Waals model, where it equals \frac{a n^2}{V^2} for n moles, reflecting attractive interactions parameterized by a.[1][3] Internal pressure contributes to understanding thermodynamic properties in various systems.[1]
Thermodynamic Foundations
Definition and Basic Relations
In thermodynamics, the internal pressure, often denoted as \pi_T, is defined as the partial derivative of the internal energy U with respect to the volume V at constant temperature T:\pi_T = \left( \frac{\partial U}{\partial V} \right)_T.This quantity characterizes the dependence of the internal energy on volume under isothermal conditions and reflects the contribution of intermolecular forces to the energy variation during expansion or compression.The concept is rooted in the foundational thermodynamic relations developed in the 19th century.The internal pressure connects to other thermodynamic state functions through the fundamental identity for the differential of internal energy: dU = T \, dS - P \, dV, where T is the absolute temperature, S is the entropy, and P is the mechanical pressure.[4] Taking the partial derivative with respect to V at constant T yields \left( \frac{\partial U}{\partial V} \right)_T = T \left( \frac{\partial S}{\partial V} \right)_T - P.[5]The term \left( \frac{\partial S}{\partial V} \right)_T is linked to pressure via the Maxwell relation \left( \frac{\partial S}{\partial V} \right)_T = \left( \frac{\partial P}{\partial T} \right)_V, which arises from the equality of mixed second partial derivatives of the Helmholtz free energy A = U - T S. The differential form dA = -S \, dT - P \, dV ensures this symmetry, as \left( \frac{\partial^2 A}{\partial T \partial V} \right) = -\left( \frac{\partial S}{\partial V} \right)_T = -\left( \frac{\partial P}{\partial T} \right)_V.[6] Substituting this relation gives the basic equation for internal pressure:\pi_T = T \left( \frac{\partial P}{\partial T} \right)_V - P.This expression allows \pi_T to be evaluated using properties derivable from an equation of state.[5][7]The internal pressure shares the dimensions of pressure, with SI units of pascals (Pa) or, in other systems, atmospheres (atm) or bars. In condensed phases like liquids, its magnitude often exceeds the external pressure by orders of magnitude, typically ranging from hundreds to thousands of bars, underscoring the role of attractive forces in maintaining cohesion.
Derivation of the Equation of State
The internal pressure \pi_T, defined as \left( \frac{\partial U}{\partial V} \right)_T for a closed system containing a single substance, arises from the fundamental thermodynamic relation dU = T\, dS - P\, dV, which holds for reversible processes in equilibrium.[1]To derive an explicit expression, consider the internal energy as a function of temperature and volume, U = U(T, V). The total differential is dU = \left( \frac{\partial U}{\partial T} \right)_V dT + \left( \frac{\partial U}{\partial V} \right)_T dV. At constant temperature, dT = 0, so \left( \frac{\partial U}{\partial V} \right)_T = \frac{1}{dV} (dU)_{T}. Substituting the fundamental relation at constant T yields \left( \frac{\partial U}{\partial V} \right)_T = T \left( \frac{\partial S}{\partial V} \right)_T - P, where the division by dV is implicit in the partial derivatives.[1]Applying the Maxwell relation from the equality of mixed second partial derivatives of the Helmholtz energy A(T, V), \left( \frac{\partial S}{\partial V} \right)_T = -\left( \frac{\partial^2 A}{\partial V \partial T} \right) = \left( \frac{\partial P}{\partial T} \right)_V, since P = -\left( \frac{\partial A}{\partial V} \right)_T. Thus, \pi_T = T \left( \frac{\partial P}{\partial T} \right)_V - P. This equation of state connects the internal pressure directly to the equation of state P = P(T, V) and represents a general thermodynamic identity for fluids in reversible, isothermal processes.[1]To express \left( \frac{\partial P}{\partial T} \right)_V in terms of measurable coefficients, use the cyclic relation for the variables P, T, and V: \left( \frac{\partial P}{\partial T} \right)_V \left( \frac{\partial T}{\partial V} \right)_P \left( \frac{\partial V}{\partial P} \right)_T = -1. Rearranging gives \left( \frac{\partial P}{\partial T} \right)_V = -\frac{\left( \frac{\partial V}{\partial T} \right)_P}{\left( \frac{\partial V}{\partial P} \right)_T}. The thermal expansion coefficient is \alpha = \frac{1}{V} \left( \frac{\partial V}{\partial T} \right)_P, so \left( \frac{\partial V}{\partial T} \right)_P = \alpha V. The isothermal compressibility is \kappa_T = -\frac{1}{V} \left( \frac{\partial V}{\partial P} \right)_T, so \left( \frac{\partial V}{\partial P} \right)_T = -\kappa_T V. Substituting these yields \left( \frac{\partial P}{\partial T} \right)_V = -\frac{\alpha V}{-\kappa_T V} = \frac{\alpha}{\kappa_T}.[8]Therefore, the general form of the equation of state for internal pressure is \pi_T = T \frac{\alpha}{\kappa_T} - P, valid under the assumptions of reversible thermodynamics for single-phase fluids. This derivation relies on isentropic and isothermal conditions through the partial derivatives, ensuring applicability to equilibrium states.[1]
Behavior in Gases
Ideal Gases
For ideal gases, the internal pressure \pi_T = T \left( \frac{\partial P}{\partial T} \right)_V - P evaluates to zero, as derived from the equation of state PV = nRT. Differentiating the ideal gas law with respect to temperature at constant volume yields \left( \frac{\partial P}{\partial T} \right)_V = \frac{P}{T}, so substituting gives \pi_T = T \cdot \frac{P}{T} - P = 0. This result holds for any state point, indicating that the thermodynamic tendency for expansion due to thermal effects exactly balances the direct pressure contribution.Physically, the vanishing internal pressure reflects that the internal energy U of an ideal gas depends solely on temperature, as established by Joule's law from free expansion experiments./03%3A_First_Law_of_Thermodynamics/3.01%3A_Calculation_of_Internal_Energy_Changes) With U = U(T) only, there is no volume dependence at constant temperature, implying the absence of intermolecular forces that could alter potential energy during expansion.[9] Thus, changes in volume do not affect the internal energy beyond kinetic contributions tied to temperature.From the kinetic theory perspective, ideal gas molecules are modeled as point particles undergoing elastic collisions with no intermolecular attractions or repulsions, so the internal energy arises purely from translational kinetic energy with no potential energy terms dependent on volume or density./12%3A_Temperature_and_Kinetic_Theory/12.5%3A_Kinetic_Theory) This simplifies the energy to \frac{3}{2} nRT for monatomic gases, reinforcing that volume changes at fixed temperature do not contribute to energy variations.[10]In the limit of low densities or high temperatures, real gases approach ideal behavior, where internal pressure tends to zero as intermolecular interactions become negligible compared to thermal motion.To verify numerically, consider 1 mol of an ideal gas at standard temperature and pressure (STP: T = 273 K, P = 1 atm, V = 22.4 L). The partial derivative \left( \frac{\partial P}{\partial T} \right)_V = \frac{nR}{V} = \frac{P}{T} \approx 0.00366 atm/K, so \pi_T = 273 \cdot 0.00366 - 1 = 0 atm, confirming the exact cancellation.
Real Gases
In real gases, deviations from ideal behavior arise primarily from intermolecular attractive forces and the finite size of molecules, leading to a non-zero internal pressure \pi_T = \left( \frac{\partial U}{\partial V} \right)_T. The van der Waals equation of state accounts for these effects through the form \left( P + \frac{a}{V_m^2} \right) (V_m - b) = [RT](/page/RT), where [V_m](/page/Molar_volume) is the molar volume, a reflects the strength of attractive interactions that reduce the observed pressure, and b represents the excluded volume per mole due to molecular repulsion. Applying the thermodynamic identity \pi_T = T \left( \frac{\partial P}{\partial T} \right)_{V_m} - P to this equation yields \pi_T = \frac{a}{V_m^2}.[11]This expression highlights the physical role of internal pressure in real gases: the term \frac{a}{V_m^2} quantifies the cohesive effect of attractive forces, which generate a positive \pi_T that resists volume expansion by effectively pulling molecules inward, contrasting with the zero internal pressure of ideal gases where no such interactions exist. The excluded volume b influences the kinetic contribution to pressure but does not affect \pi_T directly, as the internal energy for a van der Waals gas takes the form U = U_\text{ideal}(T) - \frac{a n^2}{V}, with the negative potential energy term arising solely from attractions.[12][11]Alternative models, such as the virial expansion of the equation of state P = \frac{RT}{V_m} + \frac{B(T) RT}{V_m^2} + \cdots, provide a series representation where the second virial coefficient B(T) encodes temperature-dependent pairwise interactions. In the low-density limit dominated by the second virial term, the internal pressure is \pi_T \approx \frac{R T^2}{V_m^2} \frac{dB}{dT}. For gases with dominant attractions, B(T) is negative at low temperatures and increases with T (\frac{dB}{dT} > 0), yielding a positive \pi_T consistent with the van der Waals result, as the van der Waals B(T) = b - \frac{a}{RT} gives \frac{dB}{dT} = \frac{a}{R T^2} and thus \pi_T = \frac{a}{V_m^2}.[13]The temperature dependence of internal pressure reflects the competition between thermal motion and intermolecular forces: in the van der Waals model, \pi_T is independent of T due to the constant a, but more realistic descriptions show \pi_T decreasing with increasing T as kinetic energy diminishes the relative influence of attractions. Near the critical point, where liquefaction becomes possible, the van der Waals model predicts \pi_T = 3 P_c at T_c, with critical molar volume V_{m,c} = 3b and P_c = \frac{a}{27 b^2}; this elevated internal pressure signifies the balance enabling phase coexistence.[12][11]Representative calculations illustrate the magnitude: for CO_2 at 300 K with van der Waals constant a = 3.64 \, \mathrm{L}^2 \cdot \mathrm{bar} \cdot \mathrm{mol}^{-2}, at a compressed molar volume V_m = 0.5 \, \mathrm{L/mol} (approximating conditions of \sim 50 \, \mathrm{bar}), \pi_T = \frac{3.64}{(0.5)^2} = 14.6 \, \mathrm{bar} \approx 14 \, \mathrm{atm}. For N_2 with a = 1.41 \, \mathrm{L}^2 \cdot \mathrm{bar} \cdot \mathrm{mol}^{-2} under similar conditions, \pi_T \approx 5.6 \, \mathrm{bar}. These values, on the order of 10--100 atm at elevated densities, underscore the significance of internal pressure in non-ideal regimes at room temperature.[14]
Experimental Aspects
Measurement Techniques
Internal pressure, denoted as \pi_T = \left( \frac{\partial U}{\partial V} \right)_T, is typically determined through indirect experimental approaches that leverage related thermodynamic quantities, as direct measurement is challenging due to its microscopic nature. One common indirect method involves computing \pi_T from the thermal pressure coefficient \left( \frac{\partial P}{\partial T} \right)_V, which relates to internal pressure via the thermodynamic identity \pi_T = T \left( \frac{\partial P}{\partial T} \right)_V - P. This coefficient is measured by monitoring pressure variations with temperature while maintaining constant volume, often using piston-cylinder apparatuses or high-pressure optical cells capable of withstanding up to several gigapascals. In these setups, a sample is confined in a rigid or precisely controlled volume, and temperature is varied incrementally while pressure is recorded with high-precision transducers, allowing differentiation to obtain \left( \frac{\partial P}{\partial T} \right)_V.[15][16]Calorimetric techniques provide a direct route to assess \left( \frac{\partial U}{\partial V} \right)_T by conducting isothermal expansion experiments, where the heat absorbed or released during controlled volume changes at constant temperature reveals the internal energy variation. In such experiments, a sample in a calorimeter undergoes reversible expansion against a piston, with heat q measured to compute \Delta U = q - P \Delta V, yielding \pi_T as the slope of U versus V at fixed T. These methods are particularly effective for liquids, employing adiabatic or isothermal calorimeters equipped with sensitive thermistors to detect minute heat fluxes, often achieving resolutions on the order of 0.01 J/g. Advanced variants integrate simultaneous measurements of heat capacity to refine the assessment, especially in dense fluids.[15][2]A practical and widely adopted indirect approach calculates internal pressure using the isobaric thermal expansion coefficient \alpha_P = \frac{1}{V} \left( \frac{\partial V}{\partial T} \right)_P and the isothermal compressibility \kappa_T = -\frac{1}{V} \left( \frac{\partial V}{\partial P} \right)_T, via the relation \pi_T = T \frac{\alpha_P}{\kappa_T} - P (as derived from thermodynamic identities). \alpha_P is determined through dilatometry, where a sample's dimensional changes are tracked as a function of temperature using push-rod or optical dilatometers, providing linear expansion data convertible to volumetric \alpha_P with uncertainties below 1% over wide temperature ranges. For \kappa_T, ultrasonic techniques measure the speed of sound c in the medium, from which the adiabatic compressibility \kappa_S = 1/(\rho c^2) is obtained and converted to \kappa_T using \kappa_T = \kappa_S \cdot C_P / C_V, with pulse-echo or Brillouin scattering methods ensuring high precision in liquids and solids.[15][17]PVT apparatuses form the core instrumentation for these measurements, integrating pressure vessels, temperature-controlled ovens, and volume-displacement sensors to collect simultaneous P, V, and T data across phases. Modern systems, such as magnetic suspension densimeters or automated piston-cylinder units, achieve pressure accuracies of ±0.1% full scale (typically up to 200 MPa) and temperature stabilities within ±0.01 K, enabling reliable computation of derived properties like \pi_T. These devices often incorporate software for real-time data acquisition and error minimization.[16][18]Key challenges in these techniques include achieving true equilibrium conditions, particularly in viscous or multiphase systems where non-equilibrium states can skew volume or pressure readings, necessitating long stabilization times and hysteresis corrections. Additionally, for real systems, adjustments for non-ideal effects like viscous dissipation or phase boundaries are required, often involving iterative modeling to align experimental data with thermodynamic consistency. Near critical points, fluctuations amplify uncertainties, demanding specialized high-resolution setups.[15][2]
Historical and Key Experiments
One of the earliest landmark experiments providing insight into internal pressure was James Prescott Joule's free expansion experiment in the 1840s. In this setup, Joule allowed a gas, such as air, to expand freely into a vacuum within an insulated system, ensuring no heat exchange or work was performed, so the internal energy U remained constant. For ideal gases, the absence of temperature change upon expansion indicated that (∂U/∂V)_T = 0, implying zero internal pressure π_T at constant temperature. This indirect evidence established the foundational behavior for ideal gases, though Joule's apparatus had limited sensitivity to detect small deviations in real gases.[19]In the 1850s, Rudolf Clausius advanced the understanding of internal pressure through his development of thermodynamic equations of state, building on experimental data showing deviations from ideal gas behavior in real gases. Clausius incorporated PVT measurements from contemporaries like Regnault, which revealed non-zero contributions to internal energy from intermolecular forces at higher pressures and densities. These experiments highlighted how real gases exhibit π_T ≠ 0, as pressure-volume relations deviated from Boyle's law, necessitating corrections in the equation of state for accurate predictions of energy changes. Clausius' analysis in his 1850 memoir formalized the first law in the form dU = đQ - P dV, with the combined form dU = T dS - P dV incorporating entropy developed in his 1865 work "The Mechanical Theory of Heat".[20]Percy Williams Bridgman's high-pressure experiments in the early 1900s provided direct confirmation of non-zero internal pressure in gases and liquids. Using piston-cylinder apparatuses capable of generating pressures up to several thousand atmospheres, Bridgman measured compressibility and PVT data for various gases and liquids, enabling calculation of (∂U/∂V)_T via thermodynamic relations like π_T = T(∂P/∂T)_V - P. His work on gases such as hydrogen and nitrogen at elevated pressures demonstrated that internal pressure arises from molecular attractions, with values increasing significantly under compression. For liquids like water, Bridgman's measurements yielded internal pressures on the order of thousands of atmospheres at room temperature, underscoring cohesive forces in condensed phases. For instance, Bridgman's data showed π_T on the order of several atmospheres for compressed gases at room temperature, validating theoretical predictions for real gas behavior.[21]Twentieth-century PVT studies further quantified internal pressure for specific gases, building on Bridgman's methods with improved precision. For air at 300 K and near-atmospheric conditions, experimental determinations yielded π_T ≈ 0.002 atm, reflecting minor deviations from ideality, though values rose to around 1-2 atm under moderate compression from intermolecular forces. These results, derived from high-accuracy calorimetry and PVT measurements, underscored the small but measurable role of π_T in atmospheric and engineering applications.[22]Post-1950s, the measurement of internal pressure evolved from direct experimental calorimetry to computational methods, leveraging digital computers for molecular simulations. Techniques like molecular dynamics allowed simulation of particle interactions in gases, computing (∂U/∂V)_T from potential energy profiles without physical high-pressure setups. This shift enabled predictions for extreme conditions inaccessible experimentally, such as high-density gases, while validating historical data through ab initio calculations.[23]