The Bent Structure and Polarity of Water
Group IV hydrides on the periodic table include carbon (C), silicon (Si), germanium (Ge) and tin (Sn), with each having four valence electrons. Each valence electron can form a covalent bond with hydrogen; the resultant molecule is symmetric and includes methane (CH4, SiH4, GeH4, SnH4), with bond angles of 109.5o, which only interact by weak van der Waals interactions arising from transient dipole moments and are nonpolar molecules. Conversely, chalcogen atoms, or group VI hydrides, include oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po), lack symmetry due to lone electron pairs in their valence shells and result in asymmetric molecules when covalently bound to two hydrogen atoms (i.e., H2O, H2S, H2Se, and H2Te) forming polar molecules (hydrogen chalcogenides).
Hydrogen chalcogenides, H2X, are triatomic molecules with a chalcogen, bent molecular configuration, and the presence of lone electron pairs that causes a deviation in the tetrahedral bond angle from 109.5 in methane (CH4) to 104.5 in the bent configuration for water and to 92.1o for H2S, 91.0o for H2Se and 90.0o for H2Te. Due to the uneven charge distribution, water can align to and shield both negative and positive charges on dissolved molecules, making them ideal solvents able to solubilize a wide range of molecules. Unlike methane, water molecules attract each other due to the permanent dipole moments arising from the valance electron pairs on oxygen, causing an uneven charge distribution between the electronegative oxygen and positively charged hydrogen atoms. When two water molecules hydrogen bond, one partially donates a hydrogen atom to the other molecule’s lone pair of electrons on its oxygen. Interactions between the positively charged hydrogen atom on one molecule of water and a negatively charged oxygen atom on another neutralize the dipole moment and result in a network of water molecules that form cooperative and anti-cooperative hydrogen bonds (Pauling 1945). A water molecule that donates an electron pair to form a hydrogen bond with another water molecule causes an electron-withdrawing effect, thereby increasing the negative charge of the oxygen atom
. The subsequent increase in electronegativity amplifies the strength of the dipole of the adjacent water molecule, which is donating its hydrogen, causing a cooperative interaction between the next water molecule, which hydrogen bonds (i.e., a stronger hydrogen bond) with the next water molecule (Bartha et al. 2003). The other water molecule that accepts a hydrogen atom attains a greater positive charge, causing its lone pair of electrons to be less electronegative, causing the next water molecule to interact less strongly. This type of hydrogen bond is termed anti-cooperative hydrogen bonding (Bartha et al. 2003). Hence, in bulk water, there is a distribution in hydrogen bond distance, strength and angle due to the complex interactions between molecules.
Dipole Moment and Autoionization
Temperature modifies the dipole moment resulting in changes to the length and strength of the hydrogen bond (Dougherty 1998), positively correlating with cohesion, dielectric constant, specific heat, and viscosity while inversely correlating with adhesion forces, density, diffusion coefficient, and thermal conductivity. The strength of hydrogen bonding is maximum at ~4oC concurring with the highest possible density for water. Although we often oversimplify by considering water a single entity when in fact, for bulk water, there is a distribution in hydrogen bond distances, strengths and angles due to the complex interactions between molecules while water coexists in numerous chemical states. For example, water contains trace isotopes of oxygen (i.e., 17O or 18O), leading to a more electronegative molecule with stronger hydrogen bonding; while others contain deuterated water (2H), which forms weaker hydrogen bonding because hydrogen protons are more positive, reducing the magnitude of the dipole moment. Additionally, due to autoionization, which endothermically generates ions OH– and H+ from water due to an excited O-H stretching overtone vibration existing for femtoseconds before recombining. The autoionization constant Kw represents an equilibrium constant between the formation of water and hydroxyl ions (OH–) and hydronium ions (H3O+ which is often simplified to H+).
2 H20(l) –>OH– (aq)+ H3O+ (aq)
Kw = [OH–][H3O+]
Any changes to a system in dynamic equilibrium follow Le Châtelier’s principle, which states that if an equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change to reestablish a new equilibrium. Changing water temperature alters dynamic equilibrium, resulting in altered autoionization constants, Kw, for water making the ability to dissociate and pH temperature dependent, not constant. The pH of pure water is only 7 at 288K (or 14.5oC), as the autoionization constant is a function of temperature. Although pH changes as a function of temperature, because there is an equal proportion of OH– to H3O+ and since neither is in excess, the solution remains neutral and does not become acidic at low temperatures nor basic at elevated temperatures.
| Temperature (K) | Kw | pH |
| 273 | 0.11×10-14 | 7.47 |
| 283 | 0.68X10-14 | 7.27 |
| 288 | 1.00X10-14 | 7.00 |
| 313 | 2.92×10-14 | 6.77 |
| 373 | 51.3×10-14 | 6.14 |
Phase Transition Temperatures and the Density of Water and Ice
The physical state of water influences food quality, microbial growth, physical disruptions of cellular components, rates of chemical reactions and sensory properties. Phase transitions, including water to ice, ice to water, water to vapor, and others, are in thermodynamic equilibrium in simple systems; however, going from a liquid to solid, as would be the case for ice cream, thermodynamic equilibrium is not achieved and is impeded by a kinetic parameter such as high viscosity. As ice cream crystallizes, water forms into ice, and the remaining unfrozen water is more concentrated in sucrose, causing a freezing point depression and increase in viscosity, impeding crystallization from reaching a thermodynamic free energy minimum endpoint. Since frozen foods are not in equilibrium, ice crystals continue to grow. Water exists in three states, solid, liquid and vapor, while both ice and liquid phases contain sub-phases with varied molecular packing into lattice structures. The liquid-vapor transition temperature of the water is pressure dependent, and water boils only boils at 100oC under atmospheric pressure (101kPa); while it would boil at ~71oC atop Mt. Everest. Elevated pressure increases the liquid-vapor transition temperature, while decreased pressure, or under vacuum, allows water to boil below 100oC. The pressure affects the vaporization temperature with important consequences on food processing when sterilizing canned liquid foods; for example, the can is placed into a pressurized retort and uses steam to increase temperature and pressure allowing the contents to heat above 100oC while remaining liquid. Cans sterilized in an unpressurized vessel cannot exceed 100oC because water vaporization causes expansion of the can, which may rupture. Conversely, a vacuum is applied to lower the boiling point when making concentrate for heat-sensitive foods prone to browning reactions, such as orange juice. Rising film evaporators reduce the pressure to ~2.9 kPa, allowing water to be removed at or near ambient temperature and preserving the color and taste. The ice-liquid boundary is less sensitive to pressure than the vapor-liquid or solid-vapor boundaries because the density of liquid and ice are more similar than vapor density; thus, no processing techniques use pressure to modify the ice-liquid boundary.
The three-phase boundaries converge at the triple point, and at this pressure/temperature combination, all three states, solid, liquid, and vapor, coexist in equilibrium. Hence, a water molecule in the vapor phase can condense into a liquid or a solid; a water molecule from ice either melts into a liquid or sublimates into a vapor. Allowing foods to sublimate below the triple point allows frozen food to be freeze-dried preserving structure and nutrients. Above the critical point imparts unique physical properties; below, the superheated vapor is commonly used in processing plants to transfer thermal energy.
Supercritical fluids behave as gas that can diffuse through porous materials and as liquid because they can dissolve small molecules. For example, supercritical CO2 removes caffeine from coffee beans and tea. Transitions between these three phases are essential for numerous aspects of food science, such as freezing crystallizing melting chocolate, evaporating water from foods to make concentrates, isolates, and powders, and sublimating ice while retaining the structure of freeze-dried berries found in breakfast cereals and ‘astronaut’ freeze-dried ice cream.
The bent configuration and dipole moment of water gives rise to its unique physical properties (including phase transition temperatures, adhesion, cohesion, latent and specific heat capacities); symmetric molecules containing group IV hydrides, CH4. SiH4, GeH4, and SnH4, are nonpolar, and the weak van der Waals forces govern adhesion resulting in significantly lower melting and boiling points compared to group VI hydrides, H2O, H2S, H2Se, and H2Te, which lack symmetry due to the presence of two lone electron pairs in their valence shells are polar.
Typically, an increase in molecular mass for elements going down a column of the periodic table corresponds to the increased strength of van der Waals interactions and, in turn, an elevation in the boiling and melting temperatures. Hence London dispersion forces explain why boiling point decreases from H2Te, which boils at -2oC, followed by H2Se at -42oC, and H2S at -60oC. If only governed by van der Waals, H2O would boil below -60oC, but because of the ability to form hydrogen bonds does not boil until 100oC.
The physical properties of water are temperature dependent; as the dipole moment of water changes, so do the strength and length of the hydrogen bond (Dougherty 1998). The maximum hydrogen bonding strength occurs at ~4oC resulting in the highest possible density for water. The density of liquid water at 100oC is 958.4 kg/m3 increasing to 998.2 kg/m3 at 20oC and 999.9 kg/m3 at 4oC before decreasing to 998.0 kg/m3 at 0oC in the liquid state.
The phase transition from liquid water to solid ice alters the molecular arrangement, and because the tetrahedral structure adopted by ice leads to small voids in volume, there is an expansion and subsequent decrease in the solid density compared to liquid density. Since ice is less dense than water, ice floats, which is extremely important because if the density increased, the ice would sink to the bottom of oceans and accumulate. Instead, natural convection that arises upon cooling causes surface water to cool to 4oC; the increased density causes natural convection and the cooled water sinks before freezing. However, the volume expansion of ice causes significant challenges in food science, specifically when freezing whole foods, including meat, fruits and vegetables. When ice forms, the expansion in volume can lead to ice crystals rupturing and decompartmentalizing cells leading to undesirable characteristics. Cellular decompartmentalization is exacerbated when crystallization occurs near the melting temperature as few nuclei form resulting in more ice crystal growth on fewer, larger crystals. Decompartmentalization is obvious when freezing whole bananas, which turn brown due to cell rupture allowing polyphenols and polyphenol oxidase to come into contact causing enzymatic browning.
Heat Capacity and Latent Heat of Fusion and Vaporization
The heat capacity (cp) (amount of heat required to produce a unit change in its temperature), thermal conductivity (k), or the ability to conduct heat, and thermal diffusivity (α) (rate of transfer of heat of a material from the hot end to the cold end) differ based on temperature and the physical state of water (e.g., solid vs. liquid). As a product freezes, the water turns to ice at the surface and propagates inwards. As the ice layer propagates from the surface of the material, it begins to freeze quicker; this is due to an increase in the thermal diffusivity (kwater(o0C) = 0.558 W/moC; kice(0oC) = 0.558 W/moC) and thermal conductivity (αwater(o0C) = 0.131×10-6 m2/s; αice(0oC) = 11.82 x10-6 m2/s) of ice compared to water. The formation of the ice layer accelerates freezing due to an increase in thermal diffusivity and thermal conductivity; however, the opposite occurs when thawed, where water at the surface acts as an insulator resulting in significantly longer thawing than freezing times. The specific heat capacity (cp) for the major constituents of the body significantly differ, with water being the highest (4.2 kJ/kg/K) and significantly greater than fat (1.7 kJ/kg/K ), protein (1.5 kJ/kg/K), carbohydrate (1.4 kJ/kg/K), and ash (0.8 kJ/kg/K). Considering the adult human body contains 60% water, the high specific heat capacity regulates body temperature at 37o, requiring significant energy input or loss to cause a deviation in this temperature. Additionally, the body further regulates temperature through evaporative cooling during perspiration, taking advantage of the high latent heat of vaporization. The high specific heat of water is also essential in maintaining climate as large bodies of water contain significant thermal energy, which resists temperature changes when air temperature drops or increases; as solar radiation increases the temperature near the equator, natural convection establishes ocean currents regulating global climate. The equation below is significant in determining heat flow during sterilization processes such as canning and pasteurization to ensure safe food.
Adhesion and Cohesion of Water
Water cohesion is the attraction between water molecules, while adhesion is the attraction of water to other molecules. Water has the highest cohesion of all non-metallic liquids arising from the strong dipole moment and the ability of water to hydrogen bond with other water molecules. The high cohesive attraction also gives rise to high surface tensions and viscosities. When other water molecules surround water, they arrange their dipole moments to create the “lowest energy state,” balancing the net dipole moments of water molecules in the lattice, which minimizes the surface area forming spherical drops.
Water molecules at the air-water surface of the droplet or a non-compatible interface contain unbalanced forces giving rise to a high surface tension which is the tendency to resist rupture by directing the net force towards the interior of the drop. In a glass capillary, when the adhesive force (attraction of water to glass) exceeds the cohesive forces (attraction of water to water), capillary action generates an upward force drawing liquid up at the liquid edges, forming a meniscus. The capillary action is essential to move water from the roots to the leaves of plants and through the capillaries of the human vascular system; adhesion is also central to solubilizing and dispersing nutrients in blood.