2.2B: Water’s States: Gas, Liquid, and Solid - Biology

2.2B: Water’s States: Gas, Liquid, and Solid - Biology

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  • Explain the biological significance of ice’s ability to float on water

Water’s States: Gas, Liquid, and Solid

The formation of hydrogen bonds is an important quality of liquid water that is crucial to life as we know it. As water molecules make hydrogen bonds with each other, water takes on some unique chemical characteristics compared to other liquids, and since living things have a high water content, understanding these chemical features is key to understanding life. In liquid water, hydrogen bonds are constantly formed and broken as the water molecules slide past each other. The breaking of these bonds is caused by the motion (kinetic energy) of the water molecules due to the heat contained in the system. When the heat is raised as water is boiled, the higher kinetic energy of the water molecules causes the hydrogen bonds to break completely and allows water molecules to escape into the air as gas (steam or water vapor). On the other hand, when the temperature of water is reduced and water freezes, the water molecules form a crystalline structure maintained by hydrogen bonding (there is not enough energy to break the hydrogen bonds). This makes ice less dense than liquid water, a phenomenon not seen in the solidification of other liquids.

Phases of matter: See what happens to intermolecular bonds during phase changes in this interactive.

Water’s lower density in its solid form is due to the way hydrogen bonds are oriented as it freezes: the water molecules are pushed farther apart compared to liquid water. With most other liquids, solidification when the temperature drops includes the lowering of kinetic energy between molecules, allowing them to pack even more tightly than in liquid form and giving the solid a greater density than the liquid.

The low density of ice, an anomaly, causes it to float at the surface of liquid water, such as an iceberg or the ice cubes in a glass of water. In lakes and ponds, ice forms on the surface of the water creating an insulating barrier that protects the animals and plant life in the pond from freezing. Without this layer of insulating ice, plants and animals living in the pond would freeze in the solid block of ice and could not survive. The detrimental effect of freezing on living organisms is caused by the expansion of ice relative to liquid water. The ice crystals that form upon freezing rupture the delicate membranes essential for the function of living cells, irreversibly damaging them. Cells can only survive freezing if the water in them is temporarily replaced by another liquid like glycerol.

Key Points

  • As water is boiled, kinetic energy causes the hydrogen bonds to break completely and allows water molecules to escape into the air as gas (steam or water vapor).
  • When water freezes, water molecules form a crystalline structure maintained by hydrogen bonding.
  • Solid water, or ice, is less dense than liquid water.
  • Ice is less dense than water because the orientation of hydrogen bonds causes molecules to push farther apart, which lowers the density.
  • For other liquids, solidification when the temperature drops includes the lowering of kinetic energy, which allows molecules to pack more tightly and makes the solid denser than its liquid form.
  • Because ice is less dense than water, it is able to float at the surface of water.

Key Terms

  • density: A measure of the amount of matter contained by a given volume.

What Are the States of Matter?

Matter occurs in four states: solids, liquids, gases, and plasma. Often the state of matter of a substance may be changed by adding or removing heat energy from it. For example, the addition of heat can melt ice into liquid water and turn water into steam.

Key Takeaways: States of Matter

  • Matter has mass and takes up space.
  • The four main states of matter are solids, liquids, gases, and plasma.
  • Under exceptional conditions, other states of matter also exist.
  • A solid has a definite shape and volume. A liquid has a definite volume, but takes the shape of its container. A gas lacks either a defined shape or volume. Plasma is similar to a gas in that its particles are very far apart, but a gas is electrically neutral and plasma has a charge.


Examples of phase transitions include:

  • The transitions between the solid, liquid, and gaseous phases of a single component, due to the effects of temperature and/or pressure:
  • A eutectic transformation, in which a two-component single-phase liquid is cooled and transforms into two solid phases. The same process, but beginning with a solid instead of a liquid is called a eutectoid transformation.
  • A metastable to equilibrium phase transformation. A metastable polymorph which forms rapidly due to lower surface energy will transform to an equilibrium phase given sufficient thermal input to overcome an energetic barrier.
  • A peritectic transformation, in which a two-component single-phase solid is heated and transforms into a solid phase and a liquid phase.
  • A spinodal decomposition, in which a single phase is cooled and separates into two different compositions of that same phase.
  • Transition to a mesophase between solid and liquid, such as one of the "liquid crystal" phases.
  • The transition between the ferromagnetic and paramagnetic phases of magnetic materials at the Curie point.
  • The transition between differently ordered, commensurate or incommensurate, magnetic structures, such as in cerium antimonide.
  • The martensitic transformation which occurs as one of the many phase transformations in carbon steel and stands as a model for displacive phase transformations.
  • Changes in the crystallographic structure such as between ferrite and austenite of iron.
  • Order-disorder transitions such as in alpha-titanium aluminides.
  • The dependence of the adsorption geometry on coverage and temperature, such as for hydrogen on iron (110).
  • The emergence of superconductivity in certain metals and ceramics when cooled below a critical temperature.
  • The transition between different molecular structures (polymorphs, allotropes or polyamorphs), especially of solids, such as between an amorphous structure and a crystal structure, between two different crystal structures, or between two amorphous structures.
  • Quantum condensation of bosonic fluids (Bose–Einstein condensation). The superfluid transition in liquid helium is an example of this.
  • The breaking of symmetries in the laws of physics during the early history of the universe as its temperature cooled. occurs during a phase transition, the ratio of light to heavy isotopes in the involved molecules changes. When water vapor condenses (an equilibrium fractionation), the heavier water isotopes ( 18 O and 2 H) become enriched in the liquid phase while the lighter isotopes ( 16 O and 1 H) tend toward the vapor phase. [1]

Phase transitions occur when the thermodynamic free energy of a system is non-analytic for some choice of thermodynamic variables (cf. phases). This condition generally stems from the interactions of a large number of particles in a system, and does not appear in systems that are too small. It is important to note that phase transitions can occur and are defined for non-thermodynamic systems, where temperature is not a parameter. Examples include: quantum phase transitions, dynamic phase transitions, and topological (structural) phase transitions. In these types of systems other parameters take the place of temperature. For instance, connection probability replaces temperature for percolating networks.

At the phase transition point (for instance, boiling point) the two phases of a substance, liquid and vapor, have identical free energies and therefore are equally likely to exist. Below the boiling point, the liquid is the more stable state of the two, whereas above the gaseous form is preferred.

It is sometimes possible to change the state of a system diabatically (as opposed to adiabatically) in such a way that it can be brought past a phase transition point without undergoing a phase transition. The resulting state is metastable, i.e., less stable than the phase to which the transition would have occurred, but not unstable either. This occurs in superheating, supercooling, and supersaturation, for example.

Ehrenfest classification Edit

Paul Ehrenfest classified phase transitions based on the behavior of the thermodynamic free energy as a function of other thermodynamic variables. [2] Under this scheme, phase transitions were labeled by the lowest derivative of the free energy that is discontinuous at the transition. First-order phase transitions exhibit a discontinuity in the first derivative of the free energy with respect to some thermodynamic variable. [3] The various solid/liquid/gas transitions are classified as first-order transitions because they involve a discontinuous change in density, which is the (inverse of the) first derivative of the free energy with respect to pressure. Second-order phase transitions are continuous in the first derivative (the order parameter, which is the first derivative of the free energy with respect to the external field, is continuous across the transition) but exhibit discontinuity in a second derivative of the free energy. [3] These include the ferromagnetic phase transition in materials such as iron, where the magnetization, which is the first derivative of the free energy with respect to the applied magnetic field strength, increases continuously from zero as the temperature is lowered below the Curie temperature. The magnetic susceptibility, the second derivative of the free energy with the field, changes discontinuously. Under the Ehrenfest classification scheme, there could in principle be third, fourth, and higher-order phase transitions.

The Ehrenfest classification implicitly allows for continuous phase transformations, where the bonding character of a material changes, but there is no discontinuity in any free energy derivative. An example of this occurs at the supercritical liquid–gas boundaries.

Modern classifications Edit

In the modern classification scheme, phase transitions are divided into two broad categories, named similarly to the Ehrenfest classes: [2]

First-order phase transitions are those that involve a latent heat. During such a transition, a system either absorbs or releases a fixed (and typically large) amount of energy per volume. During this process, the temperature of the system will stay constant as heat is added: the system is in a "mixed-phase regime" in which some parts of the system have completed the transition and others have not. [4] [5] Familiar examples are the melting of ice or the boiling of water (the water does not instantly turn into vapor, but forms a turbulent mixture of liquid water and vapor bubbles). Imry and Wortis showed that quenched disorder can broaden a first-order transition. That is, the transformation is completed over a finite range of temperatures, but phenomena like supercooling and superheating survive and hysteresis is observed on thermal cycling. [6] [7] [8]

Second-order phase transitions are also called "continuous phase transitions". They are characterized by a divergent susceptibility, an infinite correlation length, and a power law decay of correlations near criticality. Examples of second-order phase transitions are the ferromagnetic transition, superconducting transition (for a Type-I superconductor the phase transition is second-order at zero external field and for a Type-II superconductor the phase transition is second-order for both normal-state–mixed-state and mixed-state–superconducting-state transitions) and the superfluid transition. In contrast to viscosity, thermal expansion and heat capacity of amorphous materials show a relatively sudden change at the glass transition temperature [9] which enables accurate detection using differential scanning calorimetry measurements. Lev Landau gave a phenomenological theory of second-order phase transitions.

Apart from isolated, simple phase transitions, there exist transition lines as well as multicritical points, when varying external parameters like the magnetic field or composition.

Several transitions are known as infinite-order phase transitions. They are continuous but break no symmetries. The most famous example is the Kosterlitz–Thouless transition in the two-dimensional XY model. Many quantum phase transitions, e.g., in two-dimensional electron gases, belong to this class.

The liquid–glass transition is observed in many polymers and other liquids that can be supercooled far below the melting point of the crystalline phase. This is atypical in several respects. It is not a transition between thermodynamic ground states: it is widely believed that the true ground state is always crystalline. Glass is a quenched disorder state, and its entropy, density, and so on, depend on the thermal history. Therefore, the glass transition is primarily a dynamic phenomenon: on cooling a liquid, internal degrees of freedom successively fall out of equilibrium. Some theoretical methods predict an underlying phase transition in the hypothetical limit of infinitely long relaxation times. [10] [11] No direct experimental evidence supports the existence of these transitions.

The gelation transition of colloidal particles has been shown to be a second-order phase transition under nonequilibrium conditions. [12]

Phase coexistence Edit

A disorder-broadened first-order transition occurs over a finite range of temperatures where the fraction of the low-temperature equilibrium phase grows from zero to one (100%) as the temperature is lowered. This continuous variation of the coexisting fractions with temperature raised interesting possibilities. On cooling, some liquids vitrify into a glass rather than transform to the equilibrium crystal phase. This happens if the cooling rate is faster than a critical cooling rate, and is attributed to the molecular motions becoming so slow that the molecules cannot rearrange into the crystal positions. [13] This slowing down happens below a glass-formation temperature Tg, which may depend on the applied pressure. [9] [14] If the first-order freezing transition occurs over a range of temperatures, and Tg falls within this range, then there is an interesting possibility that the transition is arrested when it is partial and incomplete. Extending these ideas to first-order magnetic transitions being arrested at low temperatures, resulted in the observation of incomplete magnetic transitions, with two magnetic phases coexisting, down to the lowest temperature. First reported in the case of a ferromagnetic to anti-ferromagnetic transition, [15] such persistent phase coexistence has now been reported across a variety of first-order magnetic transitions. These include colossal-magnetoresistance manganite materials, [16] [17] magnetocaloric materials, [18] magnetic shape memory materials, [19] and other materials. [20] The interesting feature of these observations of Tg falling within the temperature range over which the transition occurs is that the first-order magnetic transition is influenced by magnetic field, just like the structural transition is influenced by pressure. The relative ease with which magnetic fields can be controlled, in contrast to pressure, raises the possibility that one can study the interplay between Tg and Tc in an exhaustive way. Phase coexistence across first-order magnetic transitions will then enable the resolution of outstanding issues in understanding glasses.

Critical points Edit

In any system containing liquid and gaseous phases, there exists a special combination of pressure and temperature, known as the critical point, at which the transition between liquid and gas becomes a second-order transition. Near the critical point, the fluid is sufficiently hot and compressed that the distinction between the liquid and gaseous phases is almost non-existent. This is associated with the phenomenon of critical opalescence, a milky appearance of the liquid due to density fluctuations at all possible wavelengths (including those of visible light).

Symmetry Edit

Phase transitions often involve a symmetry breaking process. For instance, the cooling of a fluid into a crystalline solid breaks continuous translation symmetry: each point in the fluid has the same properties, but each point in a crystal does not have the same properties (unless the points are chosen from the lattice points of the crystal lattice). Typically, the high-temperature phase contains more symmetries than the low-temperature phase due to spontaneous symmetry breaking, with the exception of certain accidental symmetries (e.g. the formation of heavy virtual particles, which only occurs at low temperatures). [21]

Order parameters Edit

An order parameter is a measure of the degree of order across the boundaries in a phase transition system it normally ranges between zero in one phase (usually above the critical point) and nonzero in the other. [22] At the critical point, the order parameter susceptibility will usually diverge.

An example of an order parameter is the net magnetization in a ferromagnetic system undergoing a phase transition. For liquid/gas transitions, the order parameter is the difference of the densities.

From a theoretical perspective, order parameters arise from symmetry breaking. When this happens, one needs to introduce one or more extra variables to describe the state of the system. For example, in the ferromagnetic phase, one must provide the net magnetization, whose direction was spontaneously chosen when the system cooled below the Curie point. However, note that order parameters can also be defined for non-symmetry-breaking transitions.

Some phase transitions, such as superconducting and ferromagnetic, can have order parameters for more than one degree of freedom. In such phases, the order parameter may take the form of a complex number, a vector, or even a tensor, the magnitude of which goes to zero at the phase transition. [23]

There also exist dual descriptions of phase transitions in terms of disorder parameters. These indicate the presence of line-like excitations such as vortex- or defect lines.

Relevance in cosmology Edit

Symmetry-breaking phase transitions play an important role in cosmology. As the universe expanded and cooled, the vacuum underwent a series of symmetry-breaking phase transitions. For example, the electroweak transition broke the SU(2)×U(1) symmetry of the electroweak field into the U(1) symmetry of the present-day electromagnetic field. This transition is important to explain the asymmetry between the amount of matter and antimatter in the present-day universe, according to electroweak baryogenesis theory.

Progressive phase transitions in an expanding universe are implicated in the development of order in the universe, as is illustrated by the work of Eric Chaisson [24] and David Layzer. [25]

Critical exponents and universality classes Edit

Continuous phase transitions are easier to study than first-order transitions due to the absence of latent heat, and they have been discovered to have many interesting properties. The phenomena associated with continuous phase transitions are called critical phenomena, due to their association with critical points.

It turns out that continuous phase transitions can be characterized by parameters known as critical exponents. The most important one is perhaps the exponent describing the divergence of the thermal correlation length by approaching the transition. For instance, let us examine the behavior of the heat capacity near such a transition. We vary the temperature T of the system while keeping all the other thermodynamic variables fixed and find that the transition occurs at some critical temperature Tc. When T is near Tc, the heat capacity C typically has a power law behavior:

The heat capacity of amorphous materials has such a behaviour near the glass transition temperature where the universal critical exponent α = 0.59 [26] A similar behavior, but with the exponent ν instead of α, applies for the correlation length.

The exponent ν is positive. This is different with α. Its actual value depends on the type of phase transition we are considering.

It is widely believed that the critical exponents are the same above and below the critical temperature. It has now been shown that this is not necessarily true: When a continuous symmetry is explicitly broken down to a discrete symmetry by irrelevant (in the renormalization group sense) anisotropies, then some exponents (such as γ , the exponent of the susceptibility) are not identical. [27]

For −1 < α < 0, the heat capacity has a "kink" at the transition temperature. This is the behavior of liquid helium at the lambda transition from a normal state to the superfluid state, for which experiments have found α = −0.013 ± 0.003. At least one experiment was performed in the zero-gravity conditions of an orbiting satellite to minimize pressure differences in the sample. [28] This experimental value of α agrees with theoretical predictions based on variational perturbation theory. [29]

For 0 < α < 1, the heat capacity diverges at the transition temperature (though, since α < 1, the enthalpy stays finite). An example of such behavior is the 3D ferromagnetic phase transition. In the three-dimensional Ising model for uniaxial magnets, detailed theoretical studies have yielded the exponent α ≈ +0.110.

Some model systems do not obey a power-law behavior. For example, mean field theory predicts a finite discontinuity of the heat capacity at the transition temperature, and the two-dimensional Ising model has a logarithmic divergence. However, these systems are limiting cases and an exception to the rule. Real phase transitions exhibit power-law behavior.

Several other critical exponents, β, γ, δ, ν, and η, are defined, examining the power law behavior of a measurable physical quantity near the phase transition. Exponents are related by scaling relations, such as

It can be shown that there are only two independent exponents, e.g. ν and η.

It is a remarkable fact that phase transitions arising in different systems often possess the same set of critical exponents. This phenomenon is known as universality. For example, the critical exponents at the liquid–gas critical point have been found to be independent of the chemical composition of the fluid.

More impressively, but understandably from above, they are an exact match for the critical exponents of the ferromagnetic phase transition in uniaxial magnets. Such systems are said to be in the same universality class. Universality is a prediction of the renormalization group theory of phase transitions, which states that the thermodynamic properties of a system near a phase transition depend only on a small number of features, such as dimensionality and symmetry, and are insensitive to the underlying microscopic properties of the system. Again, the divergence of the correlation length is the essential point.

Critical slowing down and other phenomena Edit

There are also other critical phenomena e.g., besides static functions there is also critical dynamics. As a consequence, at a phase transition one may observe critical slowing down or speeding up. The large static universality classes of a continuous phase transition split into smaller dynamic universality classes. In addition to the critical exponents, there are also universal relations for certain static or dynamic functions of the magnetic fields and temperature differences from the critical value.

Percolation theory Edit

Another phenomenon which shows phase transitions and critical exponents is percolation. The simplest example is perhaps percolation in a two dimensional square lattice. Sites are randomly occupied with probability p. For small values of p the occupied sites form only small clusters. At a certain threshold pc a giant cluster is formed, and we have a second-order phase transition. [30] The behavior of P near pc is P

(ppc) β , where β is a critical exponent. Using percolation theory one can define all critical exponents that appear in phase transitions. [31] [30] External fields can be also defined for second order percolation systems [32] as well as for first order percolation [33] systems. Percolation has been found useful to study urban traffic and for identifying repetitive bottlenecks. [34] [35]

Phase transitions in biological systems Edit

Phase transitions play many important roles in biological systems. Examples include the lipid bilayer formation, the coil-globule transition in the process of protein folding and DNA melting, liquid crystal-like transitions in the process of DNA condensation, and cooperative ligand binding to DNA and proteins with the character of phase transition. [36]

In biological membranes, gel to liquid crystalline phase transitions play a critical role in physiological functioning of biomembranes. In gel phase, due to low fluidity of membrane lipid fatty-acyl chains, membrane proteins have restricted movement and thus are restrained in exercise of their physiological role. Plants depend critically on photosynthesis by chloroplast thylakoid membranes which are exposed cold environmental temperatures. Thylakoid membranes retain innate fluidity even at relatively low temperatures because of high degree of fatty-acyl disorder allowed by their high content of linolenic acid, 18-carbon chain with 3-double bonds. [37] Gel-to-liquid crystalline phase transition temperature of biological membranes can be determined by many techniques including calorimetry, fluorescence, spin label electron paramagnetic resonance and NMR by recording measurements of the concerned parameter by at series of sample temperatures. A simple method for its determination from 13-C NMR line intensities has also been proposed. [38]

It has been proposed that some biological systems might lie near critical points. Examples include neural networks in the salamander retina, [39] bird flocks [40] gene expression networks in Drosophila, [41] and protein folding. [42] However, it is not clear whether or not alternative reasons could explain some of the phenomena supporting arguments for criticality. [43] It has also been suggested that biological organisms share two key properties of phase transitions: the change of macroscopic behavior and the coherence of a system at a critical point. [44]

The characteristic feature of second order phase transitions is the appearance of fractals in some scale-free properties. It has long been known that protein globules are shaped by interactions with water. There are 20 amino acids that form side groups on protein peptide chains range from hydrophilic to hydrophobic, causing the former to lie near the globular surface, while the latter lie closer to the globular center. Twenty fractals were discovered in solvent associated surface areas of > 5000 protein segments. [45] The existence of these fractals proves that proteins function near critical points of second-order phase transitions.

In groups of organisms in stress (when approaching critical transitions), correlations tend to increase, while at the same time, fluctuations also increase. This effect is supported by many experiments and observations of groups of people, mice, trees, and grassy plants. [46]

A variety of methods are applied for studying the various effects. Selected examples are:

Scientists Find That Water Might Exist in a Whole New State

One of the most basic things we are taught in school science classes is that water can exist in three different states, either as solid ice, liquid water, or vapour gas. But an international team of scientists have recently found signs that liquid water might actually come in two different states.

Writing in an experimental paper, published in the International Journal of Nanotechnology, the researchers were surprised to find a number of physical properties of water change their behaviour between 50℃ and 60℃. This sign of a potential change to a second liquid state could spark a heated discussion in the scientific community. And, if confirmed, it could have implications for a range of fields, including nanotechnology and biology.

States of matter, also called “phases”, are a key concept in the study of systems made from atoms and molecules. Roughly speaking, a system formed from many molecules can be arranged in a certain number of configurations depending on its total energy. At higher temperatures (and therefore higher energies), the molecules have more possible configurations and so are more disorganised and can move about relatively freely (the gas phase). At lower temperatures, the molecules have a more limited number of configurations and so form a more ordered phase (a liquid). If the temperature goes down further, they arrange themselves in a very specific configuration, producing a solid.

This picture is common for relatively simple molecules such as carbon dioxide or methane, which have three clear, different states (liquid, solid and gas). But for more complex molecules, there is a larger number of possible configurations and this gives rise to more phases. A beautiful illustration of this is the rich behaviour of liquid crystals, which are formed by complex organic molecules and can flow like liquids, but still have a solid-like crystalline structure

Because the phase of a substance is determined by how its molecules are configured, many physical properties of that substance will change abruptly as it goes from one state to another. In the recent paper, the researchers measured several telltale physical properties of water at temperatures between 0℃ and 100℃ under normal atmospheric conditions (meaning the water was a liquid). Surprisingly, they found a kink in properties such as the water’s surface tension and its refractive index (a measure of how light travels through it) at around 50℃.

How can this be? The structure of a water molecule, H2O, is very interesting and can be pictured like a sort of arrow tip, with the two hydrogen atoms flanking the oxygen atom at the top. The electrons in the molecule tend to be distributed in a rather asymmetric way, making the oxygen side negatively charged relative to the hydrogen side. This simple structural feature leads to a kind of interaction between water molecules known as hydrogen bonding, in which the opposite charges attract each other.

This gives water properties that, in many cases, break the trends observed for other simple liquids. For example, unlike most other substances, a fixed mass of water takes up more room as a solid (ice) than as a (liquid) because of the way it molecules form a specific regular structure. Another example is the surface tension of liquid water, which is roughly twice that of other non-polar, simpler, liquids.

Water is simple enough, but not too simple. This means that one possibility for explaining the apparent extra phase of water is that it behaves a little bit like a liquid crystal. The hydrogen bonds between molecules keep some order at low temperatures, but eventually could take a second, less-ordered liquid phase at higher temperatures. This could explain the kinks observed by the researchers in their data.

If confirmed, the authors’ findings could have many applications. For example, if changes in the environment (such as temperature) cause changes in a substance’s physical properties, then this can potentially be used for sensing applications. Perhaps more fundamentally, biological systems are mostly made of water. How biological molecules (such as proteins) interact with each other likely depends on the specific manner in which water molecules arrange to form a liquid phase. Understanding how water molecules arrange themselves on average at different temperatures could shed light on the workings of how they interact in biological systems.

The discovery is an exciting opportunity for theorists and experimentalists, and a beautiful example of how even the most familiar substance still has secrets hiding within.

This article was originally published on The Conversation. Read the original article.


Right now, you are probably sitting on a chair, using a mouse or a keyboard that is resting on a desk – all these things are solids. Something is usually described as a solid if it can hold its own shape and is hard to compress (squash). The particles in most solids are closely packed together. Even though the particles are locked into place and cannot move or slide past each other, they still vibrate a tiny bit.

Ice is water in its solid form or state. Ice keeps its shape when frozen, even if it is removed from its container. However, ice is different from most solids: its molecules are less densely packed than in liquid water. This is why ice floats.

Closer Look

Water molecules in water vapor have few hydrogen bonds and more space between them, making vapor light and less dense than water or ice. While the H₂O molecules are closer together in liquid water than in solid ice, there are fewer hydrogen bonds in liquid water than in the rigid lattice of ice. Therefore, water is fluid whereas ice is solid. This video demonstrates the three states of water molecules.

Hydrogen bonds are again the key. The number of bonds between molecules determines whether water will be a solid, liquid, or gas. In the solid state, water molecules have the maximum number of hydrogen bonds (4 per molecule), giving water the rigid characteristic of ice. In its liquid state, water has fewer hydrogen bonds, which accounts for its less-structured, fluid character.


In geology, soil liquefaction refers to the process by which water-saturated, unconsolidated sediments are transformed into a substance that acts like a liquid, often in an earthquake. [6] Soil liquefaction was blamed for building collapses in the city of Palu, Indonesia in October 2018. [7]

In a related phenomenon, liquefaction of bulk materials in cargo ships may cause a dangerous shift in the load. [8] [9]

In physics and chemistry, the phase transitions from solid and gas to liquid (melting and condensation, respectively) may be referred to as liquefaction. The melting point (sometimes called liquefaction point) is the temperature and pressure at which a solid becomes a liquid. In commercial and industrial situations, the process of condensing a gas to liquid is sometimes referred to as liquefaction of gases.

Coal Edit

Coal liquefaction is the production of liquid fuels from coal using a variety of industrial processes.

Dissolution Edit

Liquefaction is also used in commercial and industrial settings to refer to mechanical dissolution of a solid by mixing, grinding or blending with a liquid.

Food preparation Edit

In kitchen or laboratory settings, solids may be chopped into smaller parts sometimes in combination with a liquid, for example in food preparation or laboratory use. This may be done with a blender, or liquidiser in British English.

In biology, liquefaction often involves organic tissue turning into a more liquid-like state. For example, liquefactive necrosis in pathology, [10] or liquefaction as a parameter in semen analysis. [11]

Water Cycle Steps

Water cycle steps are becoming less predictable as global warming changes water levels and distribution across the globe. This subcategory of the biogeochemical cycle should also not be discussed as a sequenced number of events as different modes of water uptake, transportation and return occur simultaneously and at different rates according to variances in global or local ecosystems. A mountainous region will experience significantly more sublimation and runoff, for example, when compared to flat, open plains. In fact, when discussing water cycle steps it is easier to look at the movement of water separately: going up and coming down.

Water Goes Up

Water cycle steps in the atmosphere are easy to see wherever a cloud is visible. A cloud is the result of water condensation that is added to the atmosphere by way of water evaporation, water sublimation, and water transpiration. Water can move through the troposphere by way of another water cycle step – water transportation. Water can return to the Earth’s crust through water precipitation and deposition.

The atmospheric water cycle takes place in the lowest layer of our atmosphere or the troposphere. The troposphere extends from the Earth’s surface and reaches heights of 4 miles at the two poles and up to 12 miles at the equator. The layer above – the stratosphere – contains very little water vapor.

Water vapor in the atmosphere is extremely important as these droplets are able to absorb solar energy as well as the heat that radiates from the Earth (thermal radiation). It is water vapor that regulates local climates and air temperatures. Variances in temperature, in turn, cause currents of air known as convection currents that help to create the wind patterns so often typical to a certain region, such as monsoon storms or desert zephyrs.

Transpiration is the conversion of water by plants into water vapor. In ideal conditions, plants only use around 5% of the water they take up through their root systems. One only has to see pictures of the mist above a rainforest to understand this contribution to water vapor levels in the troposphere. Under the sun’s rays, water escapes through leaf pores as a gas. The combination of evaporation and transpiration is called evapotranspiration. While transpiration is probably responsible for 10% of the troposphere’s water content, combined evapotranspiration provides about 99%.

Transportation does not provide water vapor to the troposphere but describes the movement of water via the wind or the jet streams – strong wind currents at the top of the troposphere or at the tropopause, a level of air between the troposphere and stratosphere. We can see the effects of transportation by watching clouds move across the sky. In addition, winds remove water vapor from the air above sources of water. This lowers the saturation levels (or humidity) of the air and allows even more water vapor to enter the atmosphere.

Water Goes Down

Water cycle steps on the Earth’s crust are highly dependent on the type of ecosystem. These steps are water condensation, precipitation, and deposition.

Water does not fall to earth in the form of water vapor. As water vapor rises, it loses heat energy through continuous motion. In addition, gaseous forms of water experience less pressure as they rise. Where there is less pressure, the air is unable to hold as much water as when pressures are high. Furthermore, other substances in the air such as pollen, pollutants, and dust provide a surface on which water vapor can settle and condense. Condensation is the opposite of evaporation and we have all seen the effect of condensation on windows and bathroom mirrors. As warm water vapor hits a cooler surface, energy levels dramatically drop. The water molecules no longer move at rapid rates and settle as water droplets. This also occurs in the atmosphere in the presence of condensation nuclei – small particles onto which water vapor can settle.

Clouds are the result of condensed water vapor. Eventually, they become saturated and are no longer able to hold liquid water droplets. This leads to precipitation.

Rain is the most common example of water cycle precipitation. Other forms are hailstones, sleet, and snow.

Deposition is the opposite of sublimation. In cases of deposition, water vapor is instantly converted from gas state to solid state (ice) without the intermediate liquid phase. In contrast to sublimation, the process of deposition releases energy. Deposition can be seen in snowfall and in the formation of frost.


Intermediary water cycle steps provide a bridge between water landing on the Earth’s surface and water vapor rising into the troposphere.

Infiltration is the absorption of water by the soil and rock of the upper level of the Earth’s crust and is very much dependent on environmental factors such as soil or rock depth, vegetation levels, saturation levels, and porosity. Percolation describes the flow of this infiltrated water through the soil or rock under the force of gravity. Eventually, percolated water will reach an impenetrable layer of non-porous rock. The water settles here in aquifers. You can make your own scale model of an aquifer by digging a deep pit in the sand when next on the beach. The pools or reservoirs of water that form above non-porous rock are called aquifers, but the water they contain is known as groundwater. Groundwater is another named phase of the water cycle and does not describe a step but the result of precipitation, infiltration and percolation.

Plant uptake is another way in which the water provided to the earth’s crust via precipitation and infiltration can be absorbed. Plant root systems take up water, using it as a nutrient source and discharging water vapor through leaf pores in the earlier described transpiration phase.

Where the ground is saturated and unable to deal with high levels of precipitation, another part of the water cycle takes place. This is water runoff. Water runoff is becoming a global problem due to the effects of global warming. Gravity is an extremely important factor when water droplets fall from the clouds. As everyone should know, water moves downhill. Where precipitation is high and the land it falls on is either limited in porosity or already saturated with water, water begins to flow downwards. Runoff may also be the result of snow melts.

Runoff is the combination of surface runoff, interflow, and baseflow. Surface runoff comes in the forms of saturation excess overland flow where the ground is already wet and unable to absorb more water, and overland flow or the runoff from our roofs, sidewalks and roads. As we increase non-porous infrastructures, we simultaneously reduce the globe’s ability to absorb precipitation. Storm runoff also occurs during heavy rainfall.

Interflow but involves water that has already percolated into lower soil levels. With the next heavy rain, this already saturated soil or rock is not given the time to reach the aquifer and water rises upwards to the soil subsurface and pushes upwards to produce increased surface runoff.

Baseflow or fair-weather flow describes how moving bodies of water such as streams and rivers take on infiltrated water over a longer period of time, between precipitation (hence ‘fair weather flow’). This is a delayed response but also contributes to runoff as an already present body of water that can increase dramatically in size in the days that follow precipitation events.

Student activities – keeping it fun while exploring the science

Student activities extend and complement the science articles. Explore states of matter with these activities:

  • Water molecules in drama – students learn about the physical characteristics of water in its different states.
  • Looking at water – solid, liquid or gas – a hands-on experience of water in its different states.
  • Solid to liquid to gas – examines the role of heat energy as water changes states.

Explore surface tension with these activities:

  • Observing water’s thin skin – three simple activities demonstrate water’s thin elastic layer.
  • Investigating bubbles – students work as scientists as they experiment with bubbles.

There are two teacher resources. Alternative conceptions about water’s states of matter matches common alternative conceptions with accurate science concepts. It includes suggested teaching points to help make conceptual changes occur. There is also a unit plan that pulls all of the resources together. The planner is in Word, so teachers can adapt the planner to suit their needs.

Different Parts of a Flame

There are several parts of a flame each is made up of different chemicals.

  • Near the base of a flame, oxygen, and fuel vapor mix as unburned gas. The composition of this part of the flame depends on the fuel that is being used.
  • Above this is the region where the molecules react with each other in the combustion reaction. Again, the reactants and products depend on the nature of the fuel.
  • Above this region, combustion is complete, and the products of the chemical reaction may be found. Typically these are water vapor and carbon dioxide. If combustion is incomplete, a fire may also give off tiny solid particles of soot or ash. Additional gases may be released from incomplete combustion, especially of "dirty" fuel, such as carbon monoxide or sulfur dioxide.

While it's difficult to see it, flames expand outward like other gases. In part, this is hard to observe because we only see the portion of the flame that is hot enough to emit light. A flame isn't round (except in space) because the hot gases are less dense than the surrounding air, so they rise up.

The color of the flame is an indication of its temperature and the chemical composition of the fuel. A flame emits incandescent light, which means that light with the highest energy (the hottest part of the flame) is blue, and that with the least energy (the coolest part of the flame) is redder. The chemistry of the fuel plays its part as well, and this is the basis for the flame test to identify chemical composition. For example, a blue flame may appear green if a boron-containing salt is present.