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Water dripping from a slab of ice and then freezing, forming icicles

Freezing is a phase transition in which a liquid turns into a solid when its temperature is lowered below its freezing point.[1][2]

For most substances, the melting and freezing points are the same temperature; however, certain substances possess differing solid-liquid transition temperatures. For example, agar displays a hysteresis in its melting point and freezing point. It melts at 85 °C (185 °F) and solidifies from 32 to 40 °C (90 to 104 °F).[3]

Crystallization

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Main article: Crystallization

Most liquids freeze by crystallization, formation of crystalline solid from the uniform liquid. This is a first-order thermodynamic phase transition, which means that as long as solid and liquid coexist, the temperature of the whole system remains very nearly equal to the melting point due to the slow removal of heat when in contact with air, which is a poor heat conductor.[citation needed] Because of the latent heat of fusion, the freezing is greatly slowed and the temperature will not drop anymore once the freezing starts but will continue dropping once it finishes.[citation needed]

Crystallization consists of two major events, nucleation and crystal growth. "Nucleation" is the step wherein the molecules start to gather into clusters, on the nanometer scale, arranging in a defined and periodic manner that defines the crystal structure. "Crystal growth" is the subsequent growth of the nuclei that succeed in achieving the critical cluster size.

Supercooling

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Main article: Supercooling
Rapid formation of ice crystals in supercool water (home freezer experiment)

Crystallization of pure liquids usually begins at a lower temperature than the melting point, due to high activation energy of homogeneous nucleation. The creation of a nucleus implies the formation of an interface at the boundaries of the new phase. Some energy is expended to form this interface, based on the surface energy of each phase. If a hypothetical nucleus is too small, the energy that would be released by forming its volume is not enough to create its surface, and nucleation does not proceed. Freezing does not start until the temperature is low enough to provide enough energy to form stable nuclei. In presence of irregularities on the surface of the containing vessel, solid or gaseous impurities, pre-formed solid crystals, or other nucleators, heterogeneous nucleation may occur, where some energy is released by the partial destruction of the previous interface, raising the supercooling point to be near or equal to the melting point. The melting point of water at 1 atmosphere of pressure is very close to 0 °C (32 °F; 273 K), and in the presence of nucleating substances the freezing point of water is close to the melting point, but in the absence of nucleators water can supercool to −40 °C (−40 °F; 233 K) before freezing.[4][5] Under high pressure (2,000 atmospheres) water will supercool to as low as −70 °C (−94 °F; 203 K) before freezing.[6]

Exothermicity

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Main article: Enthalpy of fusion

Freezing is almost always an exothermic process, meaning that as liquid changes into solid, heat and pressure are released. This is often seen as counter-intuitive, since the temperature of the material does not rise during freezing, except if the liquid were supercooled. But this can be understood since heat must be continually removed from the freezing liquid or the freezing process will stop. The energy released upon freezing is a latent heat, and is known as the enthalpy of fusion and is exactly the same as the energy required to melt the same amount of the solid.

Low-temperature helium is the only known exception to the general rule.[7] Helium-3 has a negative enthalpy of fusion at temperatures below 0.3 K. Helium-4 also has a very slightly negative enthalpy of fusion below 0.8 K. This means that, at appropriate constant pressures, heat must be added to these substances in order to freeze them.[8]

Vitrification

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Main article: Glass transition

Certain materials, such as glass and glycerol, may harden without crystallizing; these are called amorphous solids. Amorphous materials, as well as some polymers, do not have a freezing point, as there is no abrupt phase change at any specific temperature. Instead, there is a gradual change in their viscoelastic properties over a range of temperatures. Such materials are characterized by a glass transition that occurs at a glass transition temperature, which may be roughly defined as the "knee" point of the material's density vs. temperature graph. Because vitrification is a non-equilibrium process, it does not qualify as freezing, which requires an equilibrium between the crystalline and liquid state.

Freezing of living organisms

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Main article: Cryobiology

Many living organisms are able to tolerate prolonged periods of time at temperatures below the freezing point of water. Most living organisms accumulate cryoprotectants such as anti-nucleating proteins, polyols, and glucose to protect themselves against frost damage by sharp ice crystals. Most plants, in particular, can safely reach temperatures of −4 °C to −12 °C. Certain bacteria, notably Pseudomonas syringae, produce specialized proteins that serve as potent ice nucleators, which they use to force ice formation on the surface of various fruits and plants at about −2 °C.[9] The freezing causes injuries in the epithelia and makes the nutrients in the underlying plant tissues available to the bacteria.[10]

Bacteria

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Three species of bacteria, Carnobacterium pleistocenium, as well as Chryseobacterium greenlandensis and Herminiimonas glaciei, have reportedly been revived after surviving for thousands of years frozen in ice.[citation needed]

Plants

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Many plants undergo a process called hardening, which allows them to survive temperatures below 0 °C for weeks to months.

Animals

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The nematode Haemonchus contortus can survive 44 weeks frozen at liquid nitrogen temperatures. Other nematodes that survive at temperatures below 0 °C include Trichostrongylus colubriformis and Panagrolaimus davidi. Many species of reptiles and amphibians survive freezing.

Human gametes and 2-, 4- and 8-cell embryos can survive freezing and are viable for up to 10 years, a process known as cryopreservation.

Experimental attempts to freeze human beings for later revival are known as cryonics.

Food preservation

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Main article: Frozen food

Freezing is a common method of food preservation that slows both food decay and the growth of micro-organisms. Besides the effect of lower temperatures on reaction rates, freezing makes water less available for bacteria growth. Freezing is a widely used method of food preservation. Freezing generally preserves flavours, smell and nutritional content. Freezing became commercially viable.

Table

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Phase transitions of matter (
)
To
From
 Solid Liquid Gas Plasma
Solid
Melting Sublimation
Liquid Freezing
Vaporization
Gas Deposition Condensation
Ionization
Plasma Recombination

See also

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References

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[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Freezing

View on Grokipedia
from Grokipedia
Freezing is the conversion of a into a phase, typically achieved by lowering the below the substance's freezing point or by applying sufficient . This occurs at the freezing point, defined as the specific at which the and phases are in under given conditions, such as standard atmospheric pressure. For pure , this equilibrium is 0 °C (32 °F) at 1 atm, marking a fundamental reference point in thermometry and . During freezing, the process is endothermic from the perspective of the surroundings but exothermic for the system, as is released when molecules transition from disordered motion to a more ordered crystalline lattice structure. Most substances contract upon freezing due to closer molecular packing in the state, though uniquely expands by about 9% because its crystals form an open that traps air pockets, leading to lower in the form. This anomaly contributes to phenomena like floating on and the expansion of pipes in cold weather. Freezing can also involve supercooling, where a persists below its freezing point without solidifying until is triggered, such as by agitation or seeding with crystals; this metastable state is common in pure liquids and affects processes like formation in the atmosphere. In practical applications, freezing is essential for , as it slows microbial growth and enzymatic reactions by immobilizing molecules, though rapid freezing minimizes cellular damage from large crystals. Industrially, controlled freezing is used in for preserving biological samples, in via freeze-thaw cycles, and in for creating alloys with specific microstructures. These aspects highlight freezing's role in both natural processes and human-engineered systems, governed by principles of and phase equilibria.

Physical Principles of Freezing

Definition and Phase Transition

Freezing is the in which a transforms into a solid upon the removal of , typically occurring when the reaches or falls below the substance's freezing point. This process is reversible, with the reverse transition known as , where the solid reverts to a upon heating. In pure substances, freezing represents a phase transition characterized by discontinuities in properties such as and . In the of a pure substance, the freezing-melting boundary is depicted as a line separating the and regions, indicating the equilibrium conditions where both phases coexist at a specific and . This boundary slopes positively for most substances, meaning the freezing point increases with , though exceptions like exhibit a negative due to the lower density of compared to liquid . Crossing this line from the liquid side initiates freezing, while the reverse path leads to . Freezing occurs under conditions of at the freezing temperature, generally at constant , where the rate of solidification equals any potential . For , this equilibrium is achieved at 0°C (32°F) under standard of 1 , allowing pure liquid to coexist with . The process requires the liquid to be cooled sufficiently to reach this point, after which heat extraction drives the phase change. The initiation of freezing relies on , the formation of stable solid clusters within the liquid that serve as seeds for further . Homogeneous nucleation occurs spontaneously in the bulk liquid without impurities, demanding significant to overcome the energy barrier for cluster formation, typically below -38°C for . In contrast, heterogeneous nucleation is facilitated by impurities, surfaces, or foreign particles that lower the energy barrier, allowing freezing at warmer temperatures closer to the equilibrium point. This distinction is crucial, as most practical freezing events involve heterogeneous mechanisms due to the rarity of perfectly pure systems.

Thermodynamics and Energy Changes

The first law of , which states that is conserved in any process, applies directly to freezing as a where the change of the system is balanced by to the surroundings without net work done in a . During freezing, the liquid's molecular decreases as it forms a more ordered solid structure, resulting in the release of while the remains constant at the freezing point until the transition is complete. This conservation ensures that the energy liberated equals the associated with the phase change. The of fusion, LfL_f, quantifies the energy change per unit during the or of a substance at its freezing point, with the total QQ given by Q=mLfQ = m L_f, where mm is the . For , this value is 333.5 J/g at 0°C, meaning 333.5 joules of are released per gram of as it freezes into . This energy release arises from the stronger intermolecular forces in the lattice compared to the state, requiring input of equivalent energy to reverse the process during . Freezing is an exothermic process, as the system releases heat to the surroundings to achieve the lower-energy solid state, maintaining the temperature at the freezing point throughout the transition due to the absorption of this latent heat by the environment. This heat liberation is why freezing can warm nearby materials or air until solidification is fully achieved. At the freezing point, the Gibbs free energy change ΔG\Delta G for the liquid-to-solid transition is zero, indicating thermodynamic equilibrium, where ΔG=ΔH−TΔS=0\Delta G = \Delta H - T \Delta S = 0, so the enthalpy change ΔH\Delta H equals TΔST \Delta S, with ΔH\Delta H negative for the exothermic freezing and ΔS\Delta S negative due to decreased disorder./15%3A_Thermodynamics_of_Chemical_Equilibria/15.04%3A_Free_Energy_and_the_Gibbs_Function) In some systems, thermal hysteresis occurs between and temperatures because kinetic barriers, such as the energy required for , prevent the from happening at the exact equilibrium point, leading to freezing at a lower than melting. This deviation highlights the distinction between and kinetic limitations in real processes.

Crystallization Mechanisms

Crystallization during freezing begins with , the formation of initial solid clusters or embryos from the supercooled liquid, followed by growth, where these clusters expand into macroscopic crystals through the addition of molecules to the lattice. According to , nucleation involves overcoming an energy barrier due to the positive surface free energy of the embryo contrasted against the negative volume free energy gain from phase transformation; embryos smaller than a critical size dissolve, while larger ones grow spontaneously. This process is described by the nucleation rate J=Aexp⁡(−ΔG∗/kT)J = A \exp(-\Delta G^*/kT), where ΔG∗\Delta G^* is the free energy barrier, kk is Boltzmann's constant, TT is , and AA is a kinetic prefactor. Homogeneous nucleation occurs in pure liquids without impurities or surfaces, requiring significant undercooling—typically around 40°C below the for —to achieve the high needed for spontaneous cluster formation. It is rare in natural systems due to the substantial energy barrier, with rates becoming appreciable only at temperatures near -40°C for pure . In contrast, heterogeneous nucleation, the dominant mechanism in most practical freezing scenarios, is facilitated by impurities, container walls, or foreign particles that lower the energy barrier by providing a template for embryo attachment. The reduced barrier is quantified by a factor f(m)=(2+m)(1−m)24f(m) = \frac{(2 + m)(1 - m)^2}{4}, where mm is a parameter depending on the substrate; effective nucleants like mimic the lattice to minimize interfacial tension. During growth, molecules from the liquid incorporate into the lattice, propagating the ordered structure; for , this typically forms ice Ih, a with protons arranged in a wurtzite-like configuration, where oxygen atoms occupy lattice points and hydrogen bonds form a tetrahedral network. The rate of is influenced by the , which drives heat dissipation and solute rejection ahead of the advancing interface, and by the cooling rate, which determines the degree of undercooling and thus the frequency—faster cooling generally increases density but can lead to finer sizes due to limited growth time. release during growth sustains a local temperature rise at the interface, moderating the overall process.

Key Phenomena in Freezing

Supercooling

, also known as undercooling, refers to the phenomenon where a is cooled below its equilibrium freezing point without solidifying, thereby existing in a metastable that is prone to rapid upon perturbation. This state arises because the lacks sufficient kinetic or thermodynamic driving force to initiate solidification spontaneously at temperatures just below the freezing point. The occurrence of supercooling is primarily facilitated by the absence of heterogeneous nucleation sites, such as impurities or container walls that could otherwise promote formation, combined with the high purity of the and controlled, slow cooling rates that minimize the probability of homogeneous . In pure systems, these factors create a significant energy barrier to , allowing the to persist in its undercooled form, as explored further in crystallization mechanisms. The degree of varies by substance; for , it typically ranges from 5 to 15°C under conditions, while certain metals like can achieve undercoolings up to 230 K due to their atomic bonding characteristics. Once supercooled, the liquid can be induced to freeze abruptly by external triggers, including mechanical agitation that disrupts molecular arrangements, the addition of seed crystals to provide a nucleation template, or the introduction of impurities acting as heterogeneous nuclei. A practical example is observed in atmospheric science, where supercooled water droplets in clouds—often at temperatures around -15°C—remain liquid until they collide with subfreezing surfaces or ice particles, leading to the rapid formation of freezing rain or glaze ice.

Exothermicity

Freezing is an wherein the transition from to phase releases of fusion, which compensates for the cooling applied and maintains the at the equilibrium freezing point until the phase change is complete. This heat release arises from the stronger intermolecular forces in the ordered structure compared to the , resulting in a net energy expulsion to the surroundings. In practical observations, such as during experiments, the of a partially frozen sample exhibits a characteristic plateau at the freezing point, where cooling halts until solidification is fully achieved, directly attributable to the counteracting the drop. For instance, in a controlled cooling setup, the system's remains stable despite ongoing extraction, illustrating the balancing effect of the exothermic . The exothermicity of freezing has significant implications for processes in controlled environments, where the released heat can induce localized reheating, thereby slowing overall freezing rates and influencing the uniformity of solidification. This effect is particularly notable in applications requiring precise temperature management, as the must be efficiently dissipated to prevent delays in achieving complete freezing. A quantitative example is the freezing of at 0°C, where 1 kg releases approximately 334 kJ of , sufficient to prevent further temperature decline until the entire mass solidifies. The concept of and its role in exothermic freezing was first systematically explored in the by , whose thermodynamic studies on phase changes laid the foundation for understanding these energy dynamics.

Vitrification and Amorphous Solids

refers to the process by which a , upon rapid cooling, transforms into a non-crystalline, glass-like without undergoing , thereby preserving the disordered molecular structure of the state. This occurs when cooling rates are sufficiently high to suppress the and growth of , trapping the material in a kinetically arrested, highly viscous state that behaves as a solid. In the context of , vitrification produces amorphous ice, a form distinct from ordinary crystalline , and requires avoiding the thermodynamic pathway toward ordered lattice formation. The temperature, denoted as TgT_g, marks the point during cooling where the of the supercooled liquid reaches approximately 101210^{12} Pa·s, effectively rendering it an infinite-viscosity solid without a discrete phase change. For pure , TgT_g is approximately 136 K (-137°C) under hyper conditions, though this value can vary slightly depending on the method of preparation, such as vapor deposition or rapid of micrometer-sized droplets. Above TgT_g, the material exhibits liquid-like relaxation dynamics, while below it, structural changes occur gradually over extended timescales due to the frozen-in disorder. Unlike crystalline freezing, which involves a phase transition with the abrupt release of and the formation of a periodic atomic lattice, vitrification is a second-order kinetic process characterized by no evolution and continuous, non-abrupt changes in properties such as specific heat and . This absence of change distinguishes amorphous solids from crystals, as the former retain isotropic, short-range order similar to the parent liquid, leading to mechanical properties like and transparency in thin films. By circumventing , vitrification enables the study of deeply supercooled states that would otherwise crystallize. In theoretical applications, serves as a model for understanding amorphous phases in planetary ices, such as those on the satellites of , where and low temperatures stabilize non-crystalline water ice forms over geological timescales. These amorphous ices provide insights into the structural diversity of water under extreme conditions, informing models of icy body evolution in the outer solar system. Achieving vitrification poses significant challenges, primarily due to the need for ultra-fast cooling rates—typically on the order of 10610^6 K/s or higher for small samples of pure water—to outpace crystallization kinetics. For instance, experiments using laser-induced flash freezing have measured a critical cooling rate of about 6.4 × 10^6 K/s for micrometer-scale water samples, beyond which amorphous ice forms reliably. Slower rates allow sufficient time for molecular rearrangement into crystals, limiting vitrification to laboratory or specialized high-pressure environments.

Factors Influencing Freezing

Freezing Point Determination

The freezing point of a substance is determined experimentally by observing the temperature at which the liquid phase transitions to solid during controlled cooling, marking the onset of phase equilibrium between liquid and solid. Historically, cooling curve analysis has been a fundamental method for freezing point determination, involving the monitoring of temperature as a function of time while cooling a sample. In this technique, the temperature decreases steadily until the freezing point is reached, at which point a plateau appears on the plot due to the release of latent heat of fusion, maintaining a constant temperature until solidification is complete. This method, dating back to early 20th-century physical chemistry experiments, allows identification of the freezing point as the temperature of the plateau. Modern techniques, such as (DSC), provide higher precision by measuring the heat flow associated with the . In DSC, a sample and reference are heated or cooled at a constant rate, and the instrument detects the onset of the exothermic freezing process through a peak in the heat flow curve corresponding to the release, enabling accurate determination of the . This approach is widely used in and pharmaceutical analysis for its sensitivity to small thermal events. Standard procedures for freezing point measurements emphasize calibration using pure substances, with water serving as a primary reference due to its well-defined at 0.01°C. Calibration ensures accuracy to within 0.01°C by comparing the instrument's response to the known freezing behavior of under , as outlined in international temperature scales like the International Temperature Scale of 1990 (ITS-90). Such standards are maintained by organizations like NIST to support reproducible results across laboratories. Factors affecting the accuracy of freezing point determinations include variations in pressure, which alter the equilibrium according to the Clausius-Clapeyron equation: dTdP=T(Vl−Vs)ΔH\frac{dT}{dP} = \frac{T (V_\mathrm{l} - V_\mathrm{s})}{\Delta H} where TT is the , VlV_\mathrm{l} and VsV_\mathrm{s} are the molar volumes of the and , and ΔH\Delta H is the for ; for , the smaller volume of compared to results in a negative value, leading to a slight depression of the freezing point with increasing pressure. Common instrumentation for these measurements includes thermocouples, which provide reliable temperature sensing through voltage differences generated at junctions of dissimilar metals, offering precision suitable for cooling curve plots. Cryoscopes, specialized devices often equipped with automated cooling and stirring mechanisms, are employed for high-accuracy determinations, particularly in controlled environments to minimize effects.

Colligative Properties and Depression

is a of solutions in which the presence of a non-volatile solute lowers the freezing point of the compared to the pure , with the magnitude of the depression depending on the number of solute particles rather than their identity. This effect arises because the solute disrupts the 's ability to form a pure phase at the original freezing temperature, requiring a lower temperature to achieve equilibrium between the and phases. The freezing point depression, denoted as ΔTf\Delta T_f, is quantitatively described by the equation: ΔTf=Kf⋅m⋅i\Delta T_f = K_f \cdot m \cdot i where KfK_f is the cryoscopic constant specific to the solvent (for water, Kf=1.86∘C⋅kg/molK_f = 1.86^\circ \text{C} \cdot \text{kg/mol}), mm is the molality of the solute (moles of solute per kilogram of solvent), and ii is the van't Hoff factor representing the number of particles produced per solute molecule (e.g., i=2i = 2 for NaCl assuming complete dissociation). This relationship holds under ideal conditions and is derived from the principles of thermodynamics applied to phase equilibria. The underlying mechanism stems from Raoult's law, which states that the vapor pressure of the solvent in the solution is reduced proportionally to the mole fraction of the solute (P=Xsolvent⋅P∘P = X_{\text{solvent}} \cdot P^\circ), where P∘P^\circ is the vapor pressure of the pure solvent. At the freezing point, the vapor pressure of the solid phase must equal that of the liquid phase; the lowered vapor pressure of the solution shifts this equilibrium to a lower temperature. Practical examples illustrate this effect: , with a typical of about 3.5% (roughly 0.6 m NaCl equivalent), freezes at approximately -2°C rather than 0°C due to the dissolved salts. Similarly, automotive , typically a 50/50 mixture of and water (about 8.6 m), depresses the freezing point to -37°C, preventing damage in cold climates. This colligative property is most accurate for dilute solutions, where solute-solute interactions are negligible and the van't Hoff factor ii accurately reflects dissociation. In concentrated solutions or with strong electrolytes, deviations occur due to ion pairing, incomplete dissociation, or non-ideal behavior, leading to smaller-than-expected depressions.

Freezing in Biological Systems

Bacteria and Microorganisms

exhibit significant sensitivity to freezing, primarily due to the mechanical damage inflicted by formation on cell membranes and the accompanying stress. As forms extracellularly, it concentrates solutes in the unfrozen fraction, creating an osmotic gradient that draws water out of cells, leading to shrinkage, membrane rupture, and leakage of intracellular contents. This process can also denature proteins and disrupt cellular integrity, rendering many non-spore-forming non-viable upon thawing. To counter these challenges, certain bacteria employ survival mechanisms such as endospore formation, particularly in genera like Bacillus. Endospores are dormant structures with thick protective coats and minimal water content, enabling them to withstand extreme cold; for instance, Bacillus subtilis spores have demonstrated survival at temperatures as low as 10 K in simulated extraterrestrial conditions. Another key adaptation involves the production of extracellular polysaccharides (EPS), which bind to ice crystal surfaces, inhibiting their growth and recrystallization while maintaining a liquid microenvironment around cells to mitigate dehydration. Psychrophilic bacteria, adapted to polar environments, exemplify thriving in subzero conditions without freezing; species in and maintain metabolic activity below 0°C in supercooled , avoiding intracellular through membrane modifications and compatible solute accumulation. Laboratory studies on freezing tolerance reveal high survival in supercooled states for select strains; for example, rhamnosus achieves approximately 90% viability after short-term exposure to -196°C, highlighting the role of rapid cooling in preserving cellular structure. In contexts, freezing at -20°C serves as a preservation method that inactivates many bacterial pathogens by exacerbating damage and halting metabolic repair, significantly reducing populations of contaminants like and over storage time, though spores and psychrotolerant strains may persist.

Plants and Frost Resistance

Plants experience frost damage primarily through two mechanisms: extracellular freezing and intracellular freezing. Extracellular freezing occurs when ice forms outside the cells in the apoplast, drawing water from the protoplast and causing cellular dehydration, which can lead to membrane damage and impaired function if prolonged. Intracellular freezing, in contrast, involves ice crystal formation within the cell, resulting in mechanical rupture of cell structures and immediate cell death due to the expansion of ice. This type of damage is particularly lethal and often occurs during rapid temperature drops or in non-acclimated tissues. To mitigate these risks, plants employ adaptive strategies such as deep in cells, where cellular water remains liquid below 0°C, sometimes reaching -40°C in species like the katsura tree () or red osier dogwood (). This avoids ice nucleation within cells, preserving viability until extracellular ice forms. Additionally, certain plants produce antifreeze proteins (AFPs) that bind to ice crystals, inhibiting their growth and recrystallization in extracellular spaces, as observed in overwintering species like winter rye (Secale cereale). Deciduous plants enhance frost resistance by shedding leaves in autumn, which minimizes and water loss from the vascular system during winter, thereby reducing the risk of and in stems. Evergreens, retaining foliage year-round, rely on structural modifications such as thicker s and reduced cell wall porosity to promote and limit ice propagation into sensitive tissues. Seasonal acclimation, known as cold hardening, enables plants to gradually increase freezing tolerance over weeks of exposure to low but non-freezing temperatures. This process involves (ABA) hormone signaling, which upregulates the accumulation of soluble sugars like and , acting as cryoprotectants that stabilize membranes and depress the freezing point through colligative effects. These sugars also maintain osmotic balance during dehydration from extracellular freezing. Frost events pose significant agricultural challenges, leading to substantial losses; for instance, the 2007 spring freeze caused severe damage to fruits, vegetables, and field crops across the eastern U.S., including , with economic impacts exceeding hundreds of millions of dollars due to untimely budding and flowering. Such incidents highlight the vulnerability of non-acclimated crops to sudden temperature drops, underscoring the need for frost-resistant varieties in farming.

Animals and Cryoprotection

Animals exhibit diverse strategies to survive subzero temperatures, primarily through freeze tolerance, where organisms endure partial or extensive formation in their bodies, or freeze avoidance, where they prevent freezing altogether via or mechanisms. In freeze-tolerant species, such as the wood frog (Rana sylvatica), up to 65% of body water can freeze extracellularly while vital organs remain unfrozen, facilitated by rapid mobilization of glucose from liver stores acting as a cryoprotectant to minimize cellular damage. This contrasts with freeze-avoiding animals that maintain a supercooled state without nucleation. Cryoprotectants play a central role in both strategies by depressing the freezing point colligatively and stabilizing biomolecules against and low-temperature stresses. In insects, polyhydric alcohols like and methylamines such as serve as key cryoprotectants, accumulating in and tissues to promote extracellular ice formation while limiting intracellular freezing, thereby reducing size and preventing mechanical injury to cells. These compounds also stabilize membranes and enzymes by forming hydrogen bonds that counteract the disruptive effects of and , enabling species like the (Eurosta solidaginis) to tolerate temperatures as low as -40°C. In hibernating mammals, such as the (Urocitellus parryii), controlled allows body temperature to drop to approximately -3°C without freezing, supported by seasonal that enhances tolerance to ischemia and minimizes nucleation through low metabolic rates and adjustments. Freezing imposes severe physiological challenges, including ischemia from ice-blocked blood vessels that halts circulation and oxygen delivery, leading to anoxic conditions in tissues. Upon thawing, reperfusion can trigger oxidative stress through the production of reactive oxygen species (ROS), potentially damaging DNA, proteins, and lipids if not mitigated by upregulated antioxidants like glutathione and catalase, as observed in freeze-tolerant vertebrates. Evolutionary adaptations, such as antifreeze glycoproteins in Arctic codfishes (Boreogadus saida), bind to nascent ice crystals in blood plasma to inhibit their growth and prevent lethal freezing at seawater temperatures around -1.9°C, demonstrating how specialized proteins evolved de novo from ancestral genes to confer freeze avoidance in polar marine environments.

Applications and Preservation Techniques

Food Preservation Methods

Freezing serves as a key method for by halting microbial growth and enzymatic reactions through the formation of crystals that immobilize and reduce molecular mobility in matrices. The primary underlying effective food freezing is the rate at which the process occurs: rapid freezing minimizes the size of ice crystals formed, thereby reducing cellular damage and preserving texture, flavor, and structural compared to slow freezing, which produces larger crystals that can rupture cell walls and lead to drip loss upon thawing. Several commercial freezing methods are employed to achieve these rapid rates, tailored to food type and scale. Air-blast freezing involves exposing to high-velocity air at temperatures around -30°C, allowing efficient for large batches like and while being cost-effective for industrial use. Immersion freezing submerges in a cryogenic medium such as or , enabling ultra-rapid cooling that forms fine ice crystals and is particularly suited for irregularly shaped items like to maintain quality. Cryogenic freezing, often using at -196°C, provides the fastest rates by direct contact, minimizing processing time and oxidation, though it is more energy-intensive and typically reserved for high-value products like berries or prepared meals. Regarding nutritional impacts, freezing causes minimal degradation of vitamins and minerals compared to thermal methods like , as it avoids heat-induced losses and retains water-soluble nutrients such as in fruits and . However, some activity persists in the partially frozen state between -10°C and -18°C, potentially leading to gradual breakdown of quality attributes unless mitigated by pre-treatments like blanching. Freezing also inactivates most by disrupting their cellular structures, though spores may survive, emphasizing the need for proper handling to prevent post-thaw contamination. For optimal preservation, the U.S. (FDA) and U.S. Department of (USDA) recommend storing frozen foods at -18°C (0°F) or below to maintain safety and quality over extended periods. At this temperature, varies by food type; for example, frozen meats such as steaks and roasts retain acceptable quality for 9-12 months, while ground meats last 3-4 months before potential or flavor changes occur. These guidelines ensure that frozen foods remain safe indefinitely if kept continuously frozen, though quality diminishes with time due to sublimation and oxidation. The modern practice of quick freezing traces its roots to the 1920s innovations of Clarence Birdseye, who developed a rapid freezing process inspired by Inuit preservation techniques, which was patented in 1930 (U.S. Patent 1,773,079), revolutionizing the industry by enabling commercial-scale production of high-quality frozen foods like fish and vegetables. This breakthrough addressed prior limitations of slow freezing methods, paving the way for the global frozen food market and emphasizing the importance of speed in preserving sensory attributes.

Cryopreservation in Medicine and Biology

Cryopreservation in and preserves viable cells, tissues, and organs at subzero temperatures, typically using at -196°C, to halt metabolic activity and enable long-term storage for therapeutic applications such as preservation and transplantation. The two dominant techniques are slow freezing and . Slow freezing involves gradual cooling (e.g., -1°C to -2°C per minute) with penetrating cryoprotectants like (DMSO) at concentrations of 1-2 M to minimize intracellular ice formation by promoting extracellular crystallization and osmotic dehydration; this method has been standard for and early-stage embryos since the mid-20th century. In contrast, employs ultra-rapid cooling rates (>10,000°C per minute) with high concentrations of cryoprotectants (e.g., and ) to achieve a glass-like, amorphous state without ice crystals, making it the preferred approach for s and blastocysts due to superior post-thaw viability. Meta-analyses indicate yields oocyte survival rates of 90-95%, compared to 70-80% for slow freezing, with similar advantages for embryos.01261-7/fulltext) Key applications include and banking for assisted reproductive technologies (), as well as emerging organ preservation. cryopreservation, pioneered in the 1950s, enabled the first successful human pregnancies via with frozen-thawed in 1953, leading to the establishment of sperm banks that now store millions of samples annually for preservation in cases of or delayed parenthood. and cryopreservation support in vitro fertilization (IVF) cycles, allowing elective freezing for social reasons or medical necessity, with over 100,000 babies born worldwide from cryopreserved embryos by 2020. For organs, combined with has achieved functional recovery in rat kidneys stored for up to 100 days, addressing the critical shortage of transplantable organs by extending preservation beyond the current 24-48 hour limit for static cold storage. banking, initiated clinically after the first successful hematopoietic stem cell transplant in 1988 for a with , provides a source of unmatched donors for treating over 80 diseases, including leukemias. Despite these successes, challenges persist, particularly from ice-induced damage during freezing and thawing. In oocytes, intracellular ice formation can disrupt the —the protective layer—leading to structural cracks, impaired fertilization, and reduced IVF implantation rates; this is exacerbated in mature oocytes due to their large volume and low surface-to-volume ratio. Slow freezing amplifies these risks through osmotic stress and , though mitigates them, achieving survival rates of approximately 90% and live birth rates comparable to fresh transfers (30-40% per cycle). Organ-scale faces additional hurdles like fractures from uneven rewarming, limiting scalability for clinical use.00593-9/fulltext) Post-2020 advances have addressed these limitations through innovative rewarming and protocol optimization. Nanowarming, utilizing (e.g., at 0.5-1% v/v) excited by alternating , enables uniform thawing of large volumes (up to 50 mL) in seconds, preventing thermal gradients and cracking in vitrified tissues like rabbit kidneys and porcine arteries, with post-thaw viabilities exceeding 80%. As of 2025, further progress includes nanowarming techniques for liter-scale cryoprotectant volumes, enabling potential of human-sized organs like kidneys and hearts without cracking. (AI) has optimized by analyzing imaging data to select high-quality oocytes and s for freezing, predicting post-thaw outcomes with accuracies around 70-80% via models, and automating protocol adjustments for personalized cooling rates in ART labs. Ethically, raises concerns over embryo disposition and equitable access, while regulatory frameworks ensure safety; the U.S. (FDA) has overseen products since the first 1988 transplant, formalizing regulations in 2005 under 21 CFR 1271 for donor screening, processing, and to minimize risks.

Physical Data and Examples

Freezing Points of Common Substances

The freezing points of substances vary widely depending on their , molecular structure, and external conditions such as pressure and purity. These values are typically measured at standard (1 atm) for pure substances, providing a benchmark for phase transitions from to . Impurities or solutes can alter these points, often lowering them through colligative effects. Standard values are compiled from authoritative sources like the National Institute of Standards and Technology (NIST), which maintains high-precision thermodynamic data in collaboration with the International Union of Pure and Applied Chemistry (IUPAC). The following table presents freezing (or melting) points for selected common substances, illustrating the range from everyday liquids to metals and solutions. All values are for pure substances unless noted, at 1 atm , and reflect s or equilibrium conditions where applicable.
SubstanceFreezing Point (°C)Notes
(H₂O)0.00Pure, defines the scale; at 0.01°C under 611.657 Pa.
Mercury (Hg)-38.83; used as a fixed point in temperature scales.
(C₂H₅OH)-114.1Pure anhydrous; highly volatile liquid at .
Sodium chloride solution (23.3 wt% NaCl in H₂O)-21.1Eutectic point; demonstrates in aqueous solutions.
Iron (Fe)1538Pure metal; high value typical of .
These data highlight key trends in freezing points. For metals like iron, values are elevated due to strong , often exceeding 1000°C and requiring specialized high-temperature measurement techniques. In contrast, organic liquids such as exhibit low freezing points, reflecting weaker intermolecular forces. Gases, when liquefied under elevated pressure (e.g., oxygen at -218.8°C after above its critical pressure), show even lower freezing points, enabling cryogenic applications but complicating measurements at ambient conditions. Modern IUPAC-recommended values incorporate spectroscopic and calorimetric methods for precision exceeding 0.01°C in many cases.

Phase Diagrams and Visual Representations

Phase diagrams provide graphical representations of the equilibrium conditions under which different phases of a substance—solid, liquid, and gas—coexist as functions of and . For a pure substance, the diagram delineates distinct regions: the phase dominates at low temperatures and high pressures, the phase occupies intermediate temperatures and moderate pressures, and the gas phase prevails at high temperatures and low pressures. The boundaries between these regions are equilibrium curves, including the freezing curve, which traces the -liquid boundary where the substance transitions between frozen and molten states at varying pressures. The marks the unique intersection of the solid-liquid, liquid-gas, and solid-gas equilibrium curves, where all three phases coexist in stable equilibrium at a specific and pressure. Below the pressure, the substance sublimes directly from solid to gas without an intermediate liquid phase, while above it, the familiar sequence of , , , and occurs. These diagrams, rooted in thermodynamic principles, illustrate how behaves as a first-order along the solid-liquid boundary. In diagrams, which depict systems of two components such as and salt, the freezing behavior of mixtures is visualized through composition-temperature relationships at constant pressure. These diagrams feature curves that converge at the eutectic point, the lowest temperature at which the mixture can remain entirely ; beyond this point, the system freezes abruptly into a mixture of two solid phases. For the salt- , the eutectic occurs where (concentrated saltwater) coexists with and hydrohalite (NaCl·2H₂O) crystals, enabling freezing at temperatures below that of pure . Such diagrams highlight colligative effects, where solute addition lowers the freezing point of the solvent. The slope of the freezing curve in pressure-temperature diagrams reveals how applied influences the freezing point: for most substances, increased raises the freezing temperature because the phase is denser than the , favoring solidification under compression. Water exhibits an anomalous negative slope, where higher lowers the freezing point, as its form () is less dense than the due to the open hexagonal structure formed by hydrogen bonding. This inversion arises from the volume increase upon freezing, making the phase more stable under . Visual representations of water's anomalous expansion often depict versus graphs, showing a peak at approximately 4°C followed by a decline as water cools to 0°C and freezes into , which floats due to its lower (about 0.917 g/cm³ compared to 1 g/cm³ for ). Schematic sketches illustrate molecules in form clustering tightly but expanding into a lattice in , with voids accommodating the hydrogen-bonded framework. These diagrams underscore the practical implications, such as insulating bodies of . Educational tools for understanding these concepts include interactive animations that simulate phase transitions, such as the PhET States of Matter , which allows users to adjust and to observe freezing in real-time . Historically, J. Willard Gibbs laid the foundational framework for phase diagrams in the 1870s through his graphical methods and , describing equilibrium in multi-phase systems without visual illustrations but enabling later diagrammatic representations. Gibbs's 1876-1878 papers on heterogeneous substances provided the thermodynamic basis for interpreting freezing curves and eutectics.

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