Steam
Steam is the gaseous phase of water, formed when liquid water reaches its boiling point and molecules gain sufficient energy to enter the vapor state. Technically, pure steam is an invisible gas consisting of water molecules (H2O), distinct from the visible white clouds often referred to as "steam," which are actually aerosols of microscopic liquid water droplets formed by condensation in cooler air.[1][2] In thermodynamics, steam is valued for its high latent heat of vaporization, which allows it to absorb and release large amounts of energy during phase changes, making it an efficient working fluid. It can exist as saturated steam (in equilibrium with liquid water) or superheated steam (above the boiling point at a given pressure), influencing its properties like density, temperature, and energy content, as detailed in steam tables. Historically, steam powered the Industrial Revolution through engines and turbines, and today it remains essential in energy transfer processes.[1][3] Steam is produced industrially by boiling water in boilers using heat from fuels, electricity, or nuclear sources, and classified by quality (dry or wet) and pressure (low, medium, high). Its applications span power generation in steam turbines for electricity, industrial processes like chemical manufacturing and food processing, district heating systems, propulsion in ships and locomotives, and sterilization in medical and laboratory settings, underscoring its versatility and economic importance.[4][5]Fundamentals
Definition and Occurrence
Steam is the gaseous phase of water, specifically water vapor produced when liquid water reaches its boiling point of 100°C (212°F) at standard atmospheric pressure of 1 atm.[6] This form of water vapor is invisible to the naked eye, consisting of individual H₂O molecules dispersed in the air without condensing into droplets.[7] It must be distinguished from visible "steam," which is actually a suspension of tiny liquid water droplets or aerosols formed when superheated water vapor cools rapidly upon contact with cooler air, appearing as fog or mist.[6] In nature, steam plays a central role in the hydrologic cycle, primarily through evaporation from bodies of water such as oceans, lakes, and rivers, where solar energy provides the heat to convert liquid water into vapor.[8] This process accounts for about 90% of atmospheric water vapor, with the remainder contributed by transpiration, in which plants absorb groundwater and release water vapor through stomata in their leaves.[9] Atmospheric water vapor, though invisible, influences weather phenomena by rising, cooling, and condensing around particles like dust or salt crystals to form clouds, which are composed of liquid droplets rather than vapor itself.[10] Artificially, steam is generated in controlled environments by heating water to or above its boiling point, such as in household kitchens via boiling kettles or pots on stoves, where the resulting vapor can be used for cooking or humidification.[8] In laboratory settings, steam is produced using devices like autoclaves or steam generators, often for sterilization, experimentation, or demonstration of thermodynamic principles, with electric heaters or external boilers ensuring precise temperature control.[11] Early historical observations and uses of steam date back to ancient civilizations, notably in the 1st century CE when Hero of Alexandria, a Greco-Egyptian engineer, described the aeolipile—a rudimentary steam-powered device that demonstrated rotational motion from escaping vapor—though it was primarily a curiosity rather than a practical tool.[12]Physical Properties
Steam exhibits low density as a gas phase of water, significantly less than that of liquid water. At 100°C and 1 atm (saturation conditions for dry steam), its density is approximately 0.598 kg/m³, calculated from the specific volume of 1.673 m³/kg. This density decreases with increasing temperature at constant pressure due to the expansion of the gas molecules, following the ideal gas law at low pressures.[13] The dynamic viscosity of steam, which measures its resistance to flow, is around 1.26 × 10⁻⁵ Pa·s at 100°C and atmospheric pressure. This value increases slightly with temperature in the low-pressure regime, reflecting the enhanced molecular motion. Steam's thermal conductivity, indicating its ability to conduct heat, is about 0.025 W/(m·K) at 100°C, which is notably higher than that of liquid water (approximately 0.68 W/(m·K) at the same temperature) due to the gaseous structure enabling heat transfer primarily through molecular collisions and diffusion.[14][15] The isobaric specific heat capacity (Cp) of dry saturated steam is approximately 2.010 kJ/(kg·K) near 100°C, representing the heat required to raise the temperature of 1 kg of steam by 1 K at constant pressure. This value varies modestly with temperature but remains key for heat transfer calculations in steam systems. Regarding compressibility, steam approximates ideal gas behavior at low pressures (below 10 bar), where the compressibility factor Z is nearly 1, meaning volume changes linearly with pressure; however, at higher pressures, real gas effects cause deviations, with Z > 1 due to repulsive intermolecular forces dominating.[16] Optically, pure dry steam is transparent to visible light, indistinguishable from air in clarity, as water vapor molecules do not significantly scatter or absorb wavelengths in the visible spectrum. In contrast, wet steam appears opaque or milky white because of Mie scattering by microscopic water droplets suspended in the vapor.Thermodynamics
Phase Transitions
Phase transitions involving steam refer to the changes between the liquid (water) and gaseous (vapor or steam) states of H₂O, with boiling representing the liquid-to-vapor transition and condensation the reverse. These processes are governed by thermodynamic equilibrium conditions where the vapor pressure of water equals the surrounding pressure at the saturation temperature. Boiling occurs throughout the liquid volume once the saturation temperature is reached, requiring sufficient energy input to overcome intermolecular forces and achieve the latent heat of vaporization.[17] The normal boiling point of water, defined at 1 atm (101.325 kPa) pressure, is 100 °C. This temperature varies with external pressure, increasing at higher pressures and decreasing at lower ones, as described by the Clausius-Clapeyron equation: \frac{dP}{dT} = \frac{L}{T \Delta V} where L is the molar latent heat of vaporization, T is the absolute temperature, and \Delta V is the change in molar volume between the vapor and liquid phases. This relation quantifies how pressure influences the energy barrier for the phase change, with practical implications for processes like pressure cooking or high-altitude boiling.[17] Evaporation differs from boiling as a surface-limited process that can occur at any temperature below the boiling point, driven by the escape of high-energy molecules from the liquid-vapor interface into the atmosphere. In contrast, boiling involves bulk nucleation of vapor bubbles throughout the liquid at the saturation temperature, leading to vigorous phase conversion. Condensation, the reverse of vaporization, happens when steam is cooled below its saturation temperature, causing supersaturation and the formation of liquid droplets; this manifests as dew on cool surfaces or fog in the atmosphere when water vapor condenses on aerosols.[18][19] Key equilibrium points define the boundaries of these transitions: the triple point of water, at which solid, liquid, and vapor phases coexist, occurs at 0.01 °C and 611.657 Pa. The critical point marks the end of the liquid-vapor coexistence curve, at 374 °C and 22.064 MPa; above this, the distinction between liquid and gas phases disappears, yielding a supercritical fluid with properties intermediate between the two. Under rapid heating or cooling conditions, phase transitions can exhibit hysteresis due to kinetic barriers in nucleation, such as superheating where pure water exceeds 100 °C without boiling (as seen in microwave-heated containers lacking nucleation sites) or supercooling where vapor persists below the dew point before condensing.[20][21][22]Thermodynamic Properties and Steam Tables
Steam's thermodynamic properties, such as enthalpy, entropy, and latent heat of vaporization, are fundamental for analyzing energy transfers in phase change and expansion processes. Enthalpy (h) represents the total heat content, including internal energy and flow work, and is particularly important for saturated and superheated steam. For saturated steam at 100°C, the specific enthalpy of the vapor (h_g) is approximately 2676 kJ/kg. In superheated steam, enthalpy increases with temperature at constant pressure, reflecting additional sensible heat input. Entropy (s) measures the disorder or unavailable energy, with the specific entropy of saturated steam at 100°C being about 7.355 kJ/(kg·K). The latent heat of vaporization (L_v or h_fg), the energy required to change liquid water to steam at constant temperature, is approximately 2257 kJ/kg at 100°C and progressively decreases, reaching zero at the critical point where the distinction between liquid and vapor phases vanishes.[23] These properties are systematically documented in steam tables, which provide values as functions of temperature and pressure for practical engineering calculations. The standard for modern steam tables is the IAPWS-IF97 formulation, developed by the International Association for the Properties of Water and Steam (IAPWS) for industrial applications, particularly in the steam power sector. IAPWS-IF97 uses a region-based approach with the Helmholtz free energy as the fundamental thermodynamic property, enabling accurate computation of derived properties like enthalpy, entropy, and specific volume across a wide range: from the triple point (0.01°C, 0.611 kPa) to 800°C and 100 MPa, with typical uncertainties of 0.1–0.5% for enthalpy and entropy in common operating ranges. This formulation replaced earlier models for better computational efficiency and precision in tabulations and software implementations.[24][25] In applications like steam turbines, these properties facilitate analysis of isentropic processes, where expansion occurs without entropy change under ideal reversible adiabatic conditions. The isentropic efficiency (η) quantifies real performance relative to this ideal: \eta = \frac{h_1 - h_2}{h_1 - h_{2s}} Here, h_1 is the inlet enthalpy, h_2 is the actual outlet enthalpy, and h_{2s} is the outlet enthalpy for an isentropic process at the same entropy as the inlet and the actual outlet pressure; values are interpolated from IAPWS-IF97 steam tables.[26] At high pressures, steam exhibits real gas behavior, deviating from the ideal gas law (PV = RT), as captured by the compressibility factor Z = PV/(RT), which can drop below 0.3 near the critical point (374.15°C, 22.064 MPa) due to intermolecular forces. IAPWS-IF97 explicitly accounts for these deviations in property predictions, ensuring reliability in supercritical and high-pressure systems.[24] For illustration, the following table summarizes key saturated steam properties at 100°C based on IAPWS-IF97:| Property | Symbol | Value | Unit |
|---|---|---|---|
| Specific enthalpy (vapor) | h_g | 2675.5 | kJ/kg |
| Specific entropy (vapor) | s_g | 7.3549 | kJ/(kg·K) |
| Latent heat of vaporization | h_fg | 2257.0 | kJ/kg |