Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions. When vapourized, the ions are still not freed of one another. For example, in the vapour phase sodium chloride exists as diatomic "molecules".
Melting and boiling points
Ionic compound
Ionic compound | Melting and boiling points
Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions. When vapourized, the ions are still not freed of one another. Electrostatic forces between particles are strongest when the charges are high, and the distance between the nuclei of the ions is small. In such cases, the compounds generally have very high melting and boiling points and a low vapour pressure. Electrostatic forces between particles are strongest when the charges are high, and the distance between the nuclei of the ions is small. Electrostatic forces between particles are strongest when the charges are high, and the distance between the nuclei of the ions is small, which results in very high melting and boiling points
describe in general terms the structure of ionic solids such as nacl
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Ions typically pack into extremely regular crystalline structures, in an arrangement that minimizes the lattice energy.
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what can you tell me about their boiling point?
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Boiling points exhibit similar trends to melting points in terms of the size of ions and strength of other interactions. When vapourized, the ions are still not freed of one another. For example, in the vapour phase sodium chloride exists as diatomic "molecules".
Ionic compound | Melting and boiling points
Electrostatic forces between particles are strongest when the charges are high, and the distance between the nuclei of the ions is small. In such cases, the compounds generally have very high melting and boiling points and a low vapour pressure. Trends in melting points can be even better explained when the structure and ionic size ratio is taken into account. Above their melting point ionic solids melt and become molten salts (although some ionic compounds such as aluminium chloride and iron(III) chloride show molecule-like structures in the liquid phase). Inorganic compounds with simple ions typically have small ions, and thus have high melting points, so are solids at room temperature.
Ionic compound | Melting and boiling points
Some substances with larger ions, however, have a melting point below or near room temperature (often defined as up to 100 °C), and are termed ionic liquids. Ions in ionic liquids often have uneven charge distributions, or bulky substituents like hydrocarbon chains, which also play a role in determining the strength of the interactions and propensity to melt. Even when the local structure and bonding of an ionic solid is disrupted sufficiently to melt it, there are still strong long-range electrostatic forces of attraction holding the liquid together and preventing ions boiling to form a gas phase. This means that even room temperature ionic liquids have low vapour pressures, and require substantially higher temperatures to boil.
Ionic compound | Melting and boiling points
If the compound's normal boiling point is higher, then that compound can exist as a liquid or solid at that given temperature at atmospheric external pressure, and will so exist in equilibrium with its vapor (if volatile) if its vapors are contained. If a compound's vapors are not contained, then some volatile compounds can eventually evaporate away in spite of their higher boiling points. In general, compounds with ionic bonds have high normal boiling points, if they do not decompose before reaching such high temperatures. Many metals have high boiling points, but not all. Very generally—with other factors being equal—in compounds with covalently bonded molecules, as the size of the molecule (or molecular mass) increases, the normal boiling point increases.
Boiling point | Boiling point as a reference property of a pure compound
As previously discussed, ionic interactions require considerably more energy to break than hydrogen bonds, which in turn are require more energy than dipole–dipole interactions. The trends observed in their boiling points (figure 8) shows exactly the correlation expected, where sodium n-butoxide requires significantly more heat energy (higher temperature) to boil than n-butanol, which boils at a much higher temperature than diethyl ether. The heat energy required for a compound to change from liquid to gas is associated with the energy required to break the intermolecular forces each molecule experiences in its liquid state.
Non-covalent interaction | Boiling points
The characteristic melting point of metals and ionic solids is ~ 1000 °C and greater, while molecular solids typically melt closer to 300 °C (see table), thus many corresponding substances are either liquid (ice) or gaseous (oxygen) at room temperature. This is due to the elements involved, the molecules they form, and the weak intermolecular interactions of the molecules. Allotropes of phosphorus are useful to further demonstrate this structure-property relationship. White phosphorus, a molecular solid, has a relatively low density of 1.82 g/cm and melting point of 44.1 °C; it is a soft material which can be cut with a knife.
Molecular solid | Melting and boiling points
Salts characteristically have high melting points. For example, sodium chloride melts at 801 °C. Some salts with low lattice energies are liquid at or near room temperature. These include molten salts, which are usually mixtures of salts, and ionic liquids, which usually contain organic cations. These liquids exhibit unusual properties as solvents.
Salt (chemistry) | Melting point
When the molecular size becomes that of a macromolecule, polymer, or otherwise very large, the compound often decomposes at high temperature before the boiling point is reached. Another factor that affects the normal boiling point of a compound is the polarity of its molecules. As the polarity of a compound's molecules increases, its normal boiling point increases, other factors being equal. Closely related is the ability of a molecule to form hydrogen bonds (in the liquid state), which makes it harder for molecules to leave the liquid state and thus increases the normal boiling point of the compound. Simple carboxylic acids dimerize by forming hydrogen bonds between molecules.
Boiling point | Boiling point as a reference property of a pure compound
Ionic compounds containing basic ions hydroxide (OH) or oxide (O) are classified as bases. Ionic compounds without these ions are also known as salts and can be formed by acid–base reactions. Ionic compounds can also be produced from their constituent ions by evaporation of their solvent, precipitation, freezing, a solid-state reaction, or the electron transfer reaction of reactive metals with reactive non-metals, such as halogen gases. Ionic compounds typically have high melting and boiling points, and are hard and brittle. As solids they are almost always electrically insulating, but when melted or dissolved they become highly conductive, because the ions are mobilized.
Ionic compound | Introduction
Some compounds decompose at higher temperatures before reaching their normal boiling point, or sometimes even their melting point. For a stable compound, the boiling point ranges from its triple point to its critical point, depending on the external pressure. Beyond its triple point, a compound's normal boiling point, if any, is higher than its melting point. Beyond the critical point, a compound's liquid and vapor phases merge into one phase, which may be called a superheated gas. At any given temperature, if a compound's normal boiling point is lower, then that compound will generally exist as a gas at atmospheric external pressure.
Boiling point | Boiling point as a reference property of a pure compound
Any salt that melts without decomposing or vaporizing usually yields an ionic liquid. Sodium chloride (NaCl), for example, melts at into a liquid that consists largely of sodium cations () and chloride anions (). Conversely, when an ionic liquid is cooled, it often forms an ionic solid—which may be either crystalline or glassy. The ionic bond is usually stronger than the Van der Waals forces between the molecules of ordinary liquids. For that reason, common salts tend to melt at higher temperatures than other solid molecules. Some salts are liquid at or below room temperature. Examples include compounds based on the 1-ethyl-3-methylimidazolium (EMIM) cation and include: EMIM:Cl, EMIM dicyanamide, ()()·, that melts at ; and 1-butyl-3,5-dimethylpyridinium bromide which becomes a glass below .
Ionic liquid | Introduction
A standard ionic solid consists of atoms held together by ionic bonds, that is by the electrostatic attraction of opposite charges (the result of transferring electrons from atoms with lower electronegativity to atoms with higher electronegativity). Among the ionic solids are compounds formed by alkali and alkaline earth metals in combination with halogens; a classic example is table salt, sodium chloride. Ionic solids are typically of intermediate strength and extremely brittle. Melting points are typically moderately high, but some combinations of molecular cations and anions yield an ionic liquid with a freezing point below room temperature. Vapour pressures in all instances are extraordinarily low; this is a consequence of the large energy required to move a bare charge (or charge pair) from an ionic medium into free space.
Bonding in solids | Ionic solids
Bonds with partially ionic and partially covalent character are called polar covalent bonds. Ionic compounds conduct electricity when molten or in solution, typically not when solid. Ionic compounds generally have a high melting point, depending on the charge of the ions they consist of. The higher the charges the stronger the cohesive forces and the higher the melting point. They also tend to be soluble in water; the stronger the cohesive forces, the lower the solubility.
Ionic bonding | Introduction
The first atom is located at each lattice point, and the second atom is located halfway between lattice points along the fcc unit cell edge. Solid sodium chloride has a melting point of 801 °C. Thermal conductivity of sodium chloride as a function of temperature has a maximum of 2.03 W/(cm K) at and decreases to 0.069 at . It also decreases with doping.
Sodium chloride | Solid sodium chloride
As can be seen from the above plot of the logarithm of the vapor pressure vs. the temperature for any given pure chemical compound, its normal boiling point can serve as an indication of that compound's overall volatility. A given pure compound has only one normal boiling point, if any, and a compound's normal boiling point and melting point can serve as characteristic physical properties for that compound, listed in reference books. The higher a compound's normal boiling point, the less volatile that compound is overall, and conversely, the lower a compound's normal boiling point, the more volatile that compound is overall.
Boiling point | Boiling point as a reference property of a pure compound
In 1926, Yakov Frenkel suggested that in an ionic crystal like AgI, in thermodynamic equilibrium, a small fraction of the cations, α, are displaced from their regular lattice sites into interstitial positions. He related α with the Gibbs energy for the formation of one mol of Frenkel pairs, ΔG, as α = exp(-ΔG/2RT), where T is temperature and R is the gas constant; for a typical value of ΔG = 100 kJ/mol, α ~ 1 at 100 °C and ~6 at 400 °C. This idea naturally explained the presence of an appreciable fraction of mobile ions in otherwise defect-free ionic crystals, and thus the ionic conductivity in them.
Solid state ionics | Point defects in ionic crystals
This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons. As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points. (The increased nuclear charge is not a relevant factor due to the shielding effect.)
Alkali metal | Melting and boiling points
Most ionic compounds are very brittle. Once they reach the limit of their strength, they cannot deform malleably, because the strict alignment of positive and negative ions must be maintained. Instead the material undergoes fracture via cleavage. As the temperature is elevated (usually close to the melting point) a ductile–brittle transition occurs, and plastic flow becomes possible by the motion of dislocations.
Ionic compound | Brittleness
The dinitrobenzenes are all crystalline solids. The boiling points of the three isomers are relatively close; however, the melting points significantly differ. 1,4-Dinitrobenzene, which has the highest symmetry, has the highest melting point.
Dinitrobenzene | Properties
As quoted at http://www.webelements.com/ from these sources:
Boiling points of the elements (data page) | WebEl
The molar mass of salt is 58.443 g/mol, its melting point is 801 °C (1,474 °F) and its boiling point 1,465 °C (2,669 °F). Its density is 2.17 grams per cubic centimetre and it is readily soluble in water. When dissolved in water it separates into Na and Cl ions, and the solubility is 359 grams per litre. From cold solutions, salt crystallises as the dihydrate NaCl·2HO. Solutions of sodium chloride have very different properties from those of pure water; the freezing point is −21.12 °C (−6.02 °F) for 23.31 wt% of salt, and the boiling point of saturated salt solution is around 108.7 °C (227.7 °F).
Salt | Chemistry
The concept of ionic radii is based on the assumption of a spherical ion shape. However, from a group-theoretical point of view the assumption is only justified for ions that reside on high-symmetry crystal lattice sites like Na and Cl in halite or Zn and S in sphalerite. A clear distinction can be made, when the point symmetry group of the respective lattice site is considered, which are the cubic groups "O" and "T" in NaCl and ZnS. For ions on lower-symmetry sites significant deviations of their electron density from a spherical shape may occur. This holds in particular for ions on lattice sites of polar symmetry, which are the crystallographic point groups "C", "C", "C" or "C", "n" = 2, 3, 4 or 6.
Ionic radius | Non-spherical ions
In some cases the anions take on a simple cubic packing, and the resulting common structures observed are: Some ionic liquids, particularly with mixtures of anions or cations, can be cooled rapidly enough that there is not enough time for crystal nucleation to occur, so an ionic glass is formed (with no long-range order).
Ionic compound | Structure
The melting point of a substance is the point where it changes state from solid to liquid while the boiling point of a substance (in liquid state) is the point where the vapour pressure of the liquid equals the environmental pressure surrounding the liquid and all the liquid changes state to gas. As a metal is heated to its melting point, the metallic bonds keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal's boiling point. Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali metals decreases down the group.
Alkali metal | Melting and boiling points
Ionic compounds can be produced from their constituent ions by evaporation, precipitation, or freezing. Reactive metals such as the alkali metals can react directly with the highly electronegative halogen gases to form an ionic product. They can also be synthesized as the product of a high temperature reaction between solids. If the ionic compound is soluble in a solvent, it can be obtained as a solid compound by evaporating the solvent from this electrolyte solution. As the solvent is evaporated, the ions do not go into the vapour, but stay in the remaining solution, and when they become sufficiently concentrated, nucleation occurs, and they crystallize into an ionic compound.
Ionic compound | Formation
For greater pressures, published data (as of 2007) still varies by tens of gigapascals and over a thousand kelvin.
Iron | Melting and boiling points
The compressibility of an ionic compound is strongly determined by its structure, and in particular the coordination number. For example, halides with the caesium chloride structure (coordination number 8) are less compressible than those with the sodium chloride structure (coordination number 6), and less again than those with a coordination number of 4.
Ionic compound | Compressibility
If the solvation energy exceeds the lattice energy, the negative net enthalpy change of solution provides a thermodynamic drive to remove ions from their positions in the crystal and dissolve in the liquid. In addition, the entropy change of solution is usually positive for most solid solutes like ionic compounds, which means that their solubility increases when the temperature increases. There are some unusual ionic compounds such as cerium(III) sulfate, where this entropy change is negative, due to extra order induced in the water upon solution, and the solubility decreases with temperature.
Ionic compound | Solubility
A minor factor affecting boiling points is the shape of a molecule. Making the shape of a molecule more compact tends to lower the normal boiling point slightly compared to an equivalent molecule with more surface area. Most volatile compounds (anywhere near ambient temperatures) go through an intermediate liquid phase while warming up from a solid phase to eventually transform to a vapor phase. By comparison to boiling, a sublimation is a physical transformation in which a solid turns directly into vapor, which happens in a few select cases such as with carbon dioxide at atmospheric pressure. For such compounds, a sublimation point is a temperature at which a solid turning directly into vapor has a vapor pressure equal to the external pressure.
Boiling point | Boiling point as a reference property of a pure compound
Polarity is key in determining relative boiling point as it causes increased intermolecular forces, thereby raising the boiling point. In the same manner, symmetry is key in determining relative melting point as it allows for better packing in the solid state, even if it does not alter the polarity of the molecule. One example of this is the relationship between oleic acid and elaidic acid; oleic acid, the "cis" isomer, has a melting point of 13.4 °C, making it a liquid at room temperature, while the "trans" isomer, elaidic acid, has the much higher melting point of 43 °C, due to the straighter "trans" isomer being able to pack more tightly, and is solid at room temperature.
Cis–trans isomerism | Comparison of physical properties
Non-covalent interactions have a significant effect on the boiling point of a liquid. Boiling point is defined as the temperature at which the vapor pressure of a liquid is equal to the pressure surrounding the liquid. More simply, it is the temperature at which a liquid becomes a gas. As one might expect, the stronger the non-covalent interactions present for a substance, the higher its boiling point. For example, consider three compounds of similar chemical composition: sodium n-butoxide (CHONa), diethyl ether (CHO), and n-butanol (CHOH). The predominant non-covalent interactions associated with each species in solution are listed in the above figure.
Non-covalent interaction | Boiling points
The boiling point is 2072 °C (3762 °F), higher than that of thallium, but lower than gallium, conversely to the general trend of melting points, but similarly to the trends down the other post-transition metal groups because of the weakness of the metallic bonding with few electrons delocalized. The density of indium, 7.31 g/cm, is also greater than gallium, but lower than thallium. Below the critical temperature, 3.41 K, indium becomes a superconductor. Indium crystallizes in the body-centered tetragonal crystal system in the space group "I"4/"mmm" (lattice parameters: "a" = 325 pm, "c" = 495 pm): this is a slightly distorted face-centered cubic structure, where each indium atom has four neighbours at 324 pm distance and eight neighbours slightly further (336 pm).
Indium | Physical
Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as magnesium oxide. Pauling's rules provide guidelines for predicting and rationalizing the crystal structures of ionic crystals
Ionic bonding | Structures
"Ionic" metal halides (predominantly of the alkali and alkali earth metals) tend to have very high melting and boiling points. They freely dissolve in water, and some are deliquescent. They are generally poorly soluble in organic solvents. Some low-oxidation state transition metals have halides which dissolve well in water, such as ferrous chloride, nickelous chloride, and cupric chloride. Metal cations with a high oxidation state tend to undergo hydrolysis instead, e.g. ferric chloride, aluminium chloride, and titanium tetrachloride. Discrete metal halides have lower melting and boiling points. For example, titanium tetrachloride melts at −25 °C and boils at 135 °C, making it a liquid at room temperature.
Metal halides | Structure and reactivity
Rhenium hexafluoride is a liquid at room temperature. At 18.5 °C, it freezes into a yellow solid. The boiling point is 33.7 °C. The solid structure measured at −140 °C is orthorhombic space group "Pnma". Lattice parameters are "a" = 9.417 Å, "b" = 8.570 Å, and "c" = 4.965 Å. There are four formula units (in this case, discrete molecules) per unit cell, giving a density of 4.94 g·cm. The ReF molecule itself (the form important for the liquid or gas phase) has octahedral molecular geometry, which has point group ("O"). The Re–F bond length is 1.823 Å.
Rhenium hexafluoride | Description
Unless otherwise stated, melting points, boiling points, densities, crystalline structures, ionisation energies, electron affinities, and electronegativity values are from the "CRC Handbook of Physics and Chemistry;" standard electrode potentials are from the 1989 compilation by Steven Bratsch.
Nonmetal | Data sources
Room-temperature ionic liquids (RTILs) consist of bulky and asymmetric organic cations such as 1-alkyl-3-methylimidazolium, 1-alkylpyridinium, "N"-methyl-"N"-alkylpyrrolidinium and ammonium ions. Phosphonium cations are less common, but offer some advantageous properties. A wide range of anions are employed, ranging from simple halides, which generally suffer high melting points, to inorganic anions such as tetrafluoroborate and hexafluorophosphate, and to large organic anions like bistriflimide, triflate or tosylate. There are also many potential uses of ionic liquids with simple non-halogenated organic anions such as formate, alkylsulfate, alkylphosphate or glycolate. The melting point of 1-butyl-3-methylimidazolium tetrafluoroborate is about and it is a colorless liquid with high viscosity at room temperature.
Ionic liquid | Room-temperature varieties
Note: ρ is density, "n" is refractive index at 589 nm, and η is viscosity, all at 20 °C; "T" is the equilibrium temperature between two phases: ice/liquid solution for "T" < 0–0.1 °C and NaCl/liquid solution for "T" above 0.1 °C.
Water (data page) | Water with dissolved NaCl
The melting point of iron is experimentally well defined for pressures less than 50 GPa. For greater pressures, published data (as of 2007) put the γ-ε-liquid triple point at pressures that differ by tens of gigapascals and 1000 K in the melting point. Generally speaking, molecular dynamics computer simulations of iron melting and shock wave experiments suggest higher melting points and a much steeper slope of the melting curve than static experiments carried out in diamond anvil cells. The melting and boiling points of iron, along with its enthalpy of atomization, are lower than those of the earlier group 3d elements from scandium to chromium, showing the lessened contribution of the 3d electrons to metallic bonding as they are attracted more and more into the inert core by the nucleus; however, they are higher than the values for the previous element manganese because that element has a half-filled 3d subshell and consequently its d-electrons are not easily delocalized.
Allotropes of iron | Melting and boiling points
In such a lattice, it is usually not possible to distinguish discrete molecular units, so that the compounds formed are not molecular in nature. However, the ions themselves can be complex and form molecular ions like the acetate anion or the ammonium cation. For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na), and the chlorine atoms each gain an electron to form anions (Cl). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl). However, to maintain charge neutrality, strict ratios between anions and cations are observed so that ionic compounds, in general, obey the rules of stoichiometry despite not being molecular compounds.
Ionic bonding | Formation
The boiling point of neopentane is only 9.5 °C, significantly lower than those of isopentane (27.7 °C) and normal pentane (36.0 °C). Therefore, neopentane is a gas at room temperature and atmospheric pressure, while the other two isomers are (barely) liquids. The melting point of neopentane (−16.6 °C), on the other hand, is 140 degrees higher than that of isopentane (−159.9 °C) and 110 degrees higher than that of "n"-pentane (−129.8 °C). This anomaly has been attributed to the better solid-state packing assumed to be possible with the tetrahedral neopentane molecule; but this explanation has been challenged on account of it having a lower density than the other two isomers.
Neopentane | Boiling and melting points
Technetium hexafluoride is a golden-yellow solid at room temperature. Its melting point is 37.4 °C and its boiling point is 55.3 °C. Technetium hexafluoride undergoes a solid phase transition at −4.54 °C. Above this temperature (measured at 10 °C), the solid structure is cubic. Lattice parameters are "a" = 6.16 Å. There are two formula units (in this case, discrete molecules) per unit cell, giving a density of 3.02 g·cm. Below this temperature (measured at −19 °C), the solid structure is orthorhombic space group "Pnma". Lattice parameters are "a" = 9.55 Å, "b" = 8.74 Å, and "c" = 5.02 Å.
Technetium hexafluoride | Description
The melting points of the alkanes follow a similar trend to boiling points for the same reason as outlined above. That is, (all other things being equal) the larger the molecule the higher the melting point. There is one significant difference between boiling points and melting points. Solids have more rigid and fixed structure than liquids. This rigid structure requires energy to break down. Thus the better put together solid structures will require more energy to break apart. For alkanes, this can be seen from the graph above (i.e., the blue line). The odd-numbered alkanes have a lower trend in melting points than even numbered alkanes.
Alkane | Melting points
The upper limits of thermal stability of ionic liquids reported in the literature are usually based upon fast (about 10 K/min) TGA scans, and they do not imply long-term (several hours) thermal stability of ionic liquids, which is limited to less than 500 K for most ionic liquids. The solubility properties of ILs are diverse. Saturated aliphatic compounds are generally only sparingly soluble in ionic liquids, whereas alkenes show somewhat greater solubility, and aldehydes can be completely miscible. Solubility differences can be exploited in biphasic catalysis, such as hydrogenation and hydrocarbonylation processes, allowing for relatively easy separation of products and/or unreacted substrate(s).
Ionic liquid | Characteristics
Ionic radius, "r", is the radius of a monatomic ion in an ionic crystal structure. Although neither atoms nor ions have sharp boundaries, they are sometimes treated as if they were hard spheres with radii such that the sum of ionic radii of the cation and anion gives the distance between the ions in a crystal lattice. Ionic radii are typically given in units of either picometers (pm) or angstroms (Å), with 1 Å = 100 pm. Typical values range from 31 pm (0.3 Å) to over 200 pm (2 Å). The concept can be extended to solvated ions in liquid solutions taking into consideration the solvation shell.
Ionic radius | Introduction
Molecular solids tend to have lower fracture toughness (sucrose, K = 0.08 MPa m) than metal (iron, K = 50 MPa m), ionic (sodium chloride, K = 0.5 MPa m), and covalent solids (diamond, K = 5 MPa m). Molecular solids have low melting (T) and boiling (T) points compared to metal (iron), ionic (sodium chloride), and covalent solids (diamond). Examples of molecular solids with low melting and boiling temperatures include argon, water, naphthalene, nicotine, and caffeine (see table below). The constituents of molecular solids range in size from condensed monatomic gases to small molecules (i.e. naphthalene and water) to large molecules with tens of atoms (i.e.
Molecular solid | Introduction
Organic compounds typically melt and many boil. In contrast, while inorganic materials generally can be melted, many do not boil, and instead tend to degrade. In earlier times, the melting point (m.p.) and boiling point (b.p.) provided crucial information on the purity and identity of organic compounds. The melting and boiling points correlate with the polarity of the molecules and their molecular weight. Some organic compounds, especially symmetrical ones, sublime. A well-known example of a sublimable organic compound is para-dichlorobenzene, the odiferous constituent of modern mothballs. Organic compounds are usually not very stable at temperatures above 300 °C, although some exceptions exist.
Organic chemistry | Melting and boiling properties
The boiling point of a substance is the temperature at which it can change its state from a liquid to a gas. Boiling point may also refer to:
Boiling point (disambiguation) | Introduction
The word "ion" is the Greek , "ion", "going", the present participle of , "ienai", "to go". This term was introduced by English physicist and chemist Michael Faraday in 1834 for the then-unknown species that "goes" from one electrode to the other through an aqueous medium. In 1913 the crystal structure of sodium chloride was determined by William Henry Bragg and William Lawrence Bragg. This revealed that there were six equidistant nearest-neighbours for each atom, demonstrating that the constituents were not arranged in molecules or finite aggregates, but instead as a network with long-range crystalline order. Many other inorganic compounds were also found to have similar structural features.
Ionic compound | History of discovery
This is a simple way to control the concentration and ionic strength. The concentration of solutes affects many colligative properties, including increasing the osmotic pressure, and causing freezing-point depression and boiling-point elevation. Because the solutes are charged ions they also increase the electrical conductivity of the solution. The increased ionic strength reduces the thickness of the electrical double layer around colloidal particles, and therefore the stability of emulsions and suspensions. The chemical identity of the ions added is also important in many uses. For example, fluoride containing compounds are dissolved to supply fluoride ions for water fluoridation.
Ionic compound | Uses
The lattice energy of an ionic compound depends upon charges of the ions that comprise the solid. More subtly, the relative and absolute sizes of the ions influence "ΔH".
Lattice energy | Theoretical treatments
When a reasonable form is assumed for the additional repulsive energy, the total lattice energy can be modelled using the Born–Landé equation, the Born–Mayer equation, or in the absence of structural information, the Kapustinskii equation. Using an even simpler approximation of the ions as impenetrable hard spheres, the arrangement of anions in these systems are often related to close-packed arrangements of spheres, with the cations occupying tetrahedral or octahedral interstices. Depending on the stoichiometry of the ionic compound, and the coordination (principally determined by the radius ratio) of cations and anions, a variety of structures are commonly observed, and theoretically rationalized by Pauling's rules.
Ionic compound | Structure
The melting (98 °C) and boiling (883 °C) points of sodium are lower than those of lithium but higher than those of the heavier alkali metals potassium, rubidium, and caesium, following periodic trends down the group. These properties change dramatically at elevated pressures: at 1.5 Mbar, the color changes from silvery metallic to black; at 1.9 Mbar the material becomes transparent with a red color; and at 3 Mbar, sodium is a clear and transparent solid. All of these high-pressure allotropes are insulators and electrides. In a flame test, sodium and its compounds glow yellow because the excited 3s electrons of sodium emit a photon when they fall from 3p to 3s; the wavelength of this photon corresponds to the D line at about 589.3 nm.
Sodium | Physical
When ionic compounds dissolve, the individual ions dissociate and are solvated by the solvent and dispersed throughout the resulting solution. Because the ions are released into solution when dissolved, and can conduct charge, soluble ionic compounds are the most common class of strong electrolytes, and their solutions have a high electrical conductivity. The solubility is highest in polar solvents (such as water) or ionic liquids, but tends to be low in nonpolar solvents (such as petrol/gasoline). This is principally because the resulting ion–dipole interactions are significantly stronger than ion-induced dipole interactions, so the heat of solution is higher. When the oppositely charged ions in the solid ionic lattice are surrounded by the opposite pole of a polar molecule, the solid ions are pulled out of the lattice and into the liquid.
Ionic compound | Solubility
Moreover, its enthalpy of fusion is lower than the enthalpies of fusion of both "n"-pentane and isopentane, thus indicating that its high melting point is due to an entropy effect resulting from higher molecular symmetry. Indeed, the entropy of fusion of neopentane is about 4 times lower than that of "n"-pentane and isopentane.
Neopentane | Boiling and melting points
Its boiling point is expected to be about 80±30 °C, so that it is probably neither noble nor a gas; as a liquid it is expected to have a density of about 5 g/cm. It is expected to have a barely positive electron affinity (estimated as 5 kJ/mol) and a moderate ionisation energy of about 860 kJ/mol, which is rather low for a nonmetal and close to those of the metalloids tellurium and astatine. The oganesson fluorides OgF and OgF are expected to show significant ionic character, suggesting that oganesson may have at least incipient metallic properties. The oxides of oganesson, OgO and OgO, are predicted to be amphoteric.
Nonmetal | Group 18
An ionic crystal is a crystal consisting of ions bound together by their electrostatic attraction. Examples of such crystals are the alkali halides, including potassium fluoride, potassium chloride, potassium bromide, potassium iodide, sodium fluoride, and other combinations of sodium, caesium, rubidium, or lithium ions with fluoride, bromide, chloride or iodide ions. NaCl has a 6:6 co-ordination. The properties of NaCl reflect the strong interactions that exist between the ions. It is a good conductor of electricity when molten, but very poor in the solid state. When fused the mobile ions carry charge through the liquid. They are characterized by strong absorption of infrared radiation and have planes along which they cleave easily.
Ionic crystal | Introduction
that approaching the critical temperature), the solubility of ionic solutes tends to decrease due to the change of properties and structure of liquid water; the lower dielectric constant results in a less polar solvent and in a change of hydration energy affecting the ΔG of the dissolution reaction. Gaseous solutes exhibit more complex behavior with temperature. As the temperature is raised, gases usually become less soluble in water (exothermic dissolution reaction related to their hydration) (to minimum, which is below 120 °C for most permanent gases), but more soluble in organic solvents (endothermic dissolution reaction related to their solvatation). The chart shows solubility curves for some typical solid inorganic salts (temperature is in degrees Celsius i.e.
Solubility | Temperature
As a pure material, bronopol has a melting point of about 130 °C. However, due to its polymorphic characteristics, bronopol undergoes a lattice rearrangement at 100 to 105 °C and this can often be wrongly interpreted as the melting point. At temperatures above 140 °C, bronopol decomposes exothermically releasing hydrogen bromide and oxides of nitrogen.
Bronopol | Melting point
They have been found to hydrolyse. Because of their distinctive properties, ionic liquids have been investigated for many applications. Some ionic liquids can be distilled under vacuum conditions at temperatures near 300 °C. In the original work by Martyn Earle, et al., the authors wrongly concluded that the vapor was made up of individual, separated ions, but was later proven that the vapors formed consisted of ion pairs. Some ionic liquids (such as 1-butyl-3-methylimidazolium nitrate) generate flammable gases on thermal decomposition. Thermal stability and melting point depend on the liquid's components. The thermal stability of a task-specific ionic liquid, protonated betaine bis(trifluoromethanesulfonyl)imide is of about and N-butyl-N-methyl pyrrolidinium bis(trifluoromethanesulfonyl)imide was thermally stable up to 640 K.
Ionic liquid | Characteristics
In solid sodium chloride, each ion is surrounded by six ions of the opposite charge as expected on electrostatic grounds. The surrounding ions are located at the vertices of a regular octahedron. In the language of close-packing, the larger chloride ions are arranged in a cubic array whereas the smaller sodium ions fill all the cubic gaps (octahedral voids) between them. This same basic structure is found in many other compounds and is commonly known as the halite or rock-salt crystal structure. It can be represented as a face-centered cubic (fcc) lattice with a two-atom basis or as two interpenetrating face centered cubic lattices.
Sodium chloride | Solid sodium chloride
The exact arrangement of ions in an ionic lattice varies according to the size of the ions in the solid.
Ionic crystal | Introduction
Alkanes experience intermolecular van der Waals forces. Stronger intermolecular van der Waals forces give rise to greater boiling points of alkanes. There are two determinants for the strength of the van der Waals forces: Under standard conditions, from CH to CH alkanes are gaseous; from CH to CH they are liquids; and after CH they are solids. As the boiling point of alkanes is primarily determined by weight, it should not be a surprise that the boiling point has almost a linear relationship with the size (molecular weight) of the molecule. As a rule of thumb, the boiling point rises 20–30 °C for each carbon added to the chain; this rule applies to other homologous series.
Alkane | Boiling point
Indeed, the ionic interaction is the electrostatic term with "l" = 0 and "l" = 0. Using that the irregular harmonics for "L" = 0 is simply and that the monopole moments and their Clebsch–Gordan coupling are we recover—as to be expected—Coulomb's law For shorter distances, where the charge distributions of the monomers overlap, the ions will repel each other because of inter-monomer exchange of the electrons. Ionic compounds have high melting and boiling points due to the large amount of energy required to break the forces between the charged ions. When molten they are also good conductors of heat and electricity, due to the free or delocalized ions.
Quantum-mechanical explanation of intermolecular interactions | Ionic interactions
The boiling point of a liquid at a given external pressure is the temperature (formula_10) at which the vapor pressure of the liquid equals the external pressure. The "normal boiling point" is the boiling point at a pressure equal to 1 atm. The boiling point of a pure solvent is increased by the addition of a non-volatile solute, and the elevation can be measured by ebullioscopy. It is found that Here "i" is the van 't Hoff factor as above, "K" is the ebullioscopic constant of the solvent (equal to 0.512 °C kg/mol for water), and "m" is the molality of the solution.
Colligative properties | Boiling point elevation (ebullioscopy)
Together with the formula above, the boiling-point elevation can in principle be used to measure the degree of dissociation or the molar mass of the solute. This kind of measurement is called ebullioscopy (Greek "boiling-viewing"). However, since superheating is difficult to avoid, precise "ΔT" measurements are difficult to carry out, which was partly overcome by the invention of the Beckmann thermometer. Furthermore, the cryoscopic constant that determines freezing-point depression is larger than the ebullioscopic constant, and since the freezing point is often easier to measure with precision, it is more common to use cryoscopy. Among many urban legends related to the effect of ebullioscopic increase, one of them leads to adding salt when cooking pasta only after water has started boiling.
Boiling-point elevation | Uses
The upper one, representing the mole fraction of the vapor at various temperatures, is called the "dew point curve".
Vapor–liquid equilibrium | Boiling-point diagrams
The melting and boiling points of iron, along with its enthalpy of atomization, are lower than those of the earlier 3d elements from scandium to chromium, showing the lessened contribution of the 3d electrons to metallic bonding as they are attracted more and more into the inert core by the nucleus; however, they are higher than the values for the previous element manganese because that element has a half-filled 3d subshell and consequently its d-electrons are not easily delocalized. This same trend appears for ruthenium but not osmium. The melting point of iron is experimentally well defined for pressures less than 50 GPa.
Iron | Melting and boiling points
In the preceding section, boiling points of pure compounds were covered. Vapor pressures and boiling points of substances can be affected by the presence of dissolved impurities (solutes) or other miscible compounds, the degree of effect depending on the concentration of the impurities or other compounds. The presence of non-volatile impurities such as salts or compounds of a volatility far lower than the main component compound decreases its mole fraction and the solution's volatility, and thus raises the normal boiling point in proportion to the concentration of the solutes. This effect is called boiling point elevation. As a common example, salt water boils at a higher temperature than pure water.
Boiling point | Impurities and mixtures
Boiling point is the temperature at which the vapor pressure of a liquid is equal to the surrounding pressure, causing the liquid to rapidly evaporate, or boil. It is closely related to vapor pressure, but is dependent on pressure. The normal boiling point is the boiling point at atmospheric pressure, but it can also be reported at higher and lower pressures.
Volatility (chemistry) | Boiling point
For an isolated spherical solid particle of diameter formula_1 in its own liquid, the Gibbs–Thomson equation for the structural melting point depression can be written: where: The "4" in the above equation comes from the spherical geometry of the solid-liquid interface.
Gibbs–Thomson equation | For particles
Although ionic compounds contain charged atoms or clusters, these materials do not typically conduct electricity to any significant extent when the substance is solid. In order to conduct, the charged particles must be mobile rather than stationary in a crystal lattice. This is achieved to some degree at high temperatures when the defect concentration increases the ionic mobility and solid state ionic conductivity is observed. When the ionic compounds are dissolved in a liquid or are melted into a liquid, they can conduct electricity because the ions become completely mobile. This conductivity gain upon dissolving or melting is sometimes used as a defining characteristic of ionic compounds.
Ionic compound | Electrical conductivity
The boiling point of an element at a given pressure is a characteristic attribute of the element. This is also true for many simple compounds including water and simple alcohols. Once boiling has started and provided that boiling remains stable and the pressure is constant, the temperature of the boiling liquid remains constant. This attribute led to the adoption of boiling points as the definition of 100°C.
Boiling | Physics
Many ionic compounds exhibit significant solubility in water or other polar solvents. Unlike molecular compounds, salts dissociate in solution into anionic and cationic components. The lattice energy, the cohesive forces between these ions within a solid, determines the solubility. The solubility is dependent on how well each ion interacts with the solvent, so certain patterns become apparent. For example, salts of sodium, potassium and ammonium are usually soluble in water. Notable exceptions include ammonium hexachloroplatinate and potassium cobaltinitrite. Most nitrates and many sulfates are water-soluble. Exceptions include barium sulfate, calcium sulfate (sparingly soluble), and lead(II) sulfate, where the 2+/2− pairing leads to high lattice energies.
Salt (chemistry) | Solubility
Carboxylic acids tend to have higher boiling points than water, because of their greater surface areas and their tendency to form stabilised dimers through hydrogen bonds. For boiling to occur, either the dimer bonds must be broken or the entire dimer arrangement must be vaporised, increasing the enthalpy of vaporization requirements significantly.
Carboxylic acid | Boiling points
The boiling points of the pentane isomers range from about 9 to 36 °C. As is the case for other alkanes, the more thickly branched isomers tend to have lower boiling points. The same tends to be true for the melting points of alkane isomers, and that of isopentane is 30 °C lower than that of "n"-pentane. However, the melting point of neopentane, the most heavily branched of the three, is 100 °C "higher" than that of isopentane. The anomalously high melting point of neopentane has been attributed to the tetrahedral molecules packing more closely in solid form. But this explanation is contradicted by the fact that neopentane has a lower density than the other two isomers.
Pentane | Physical properties
BrF has the highest thermal stability of the interhalogens with four atoms. ICl has the lowest. Chlorine trifluoride has a boiling point of −12 °C. Bromine trifluoride has a boiling point of 127 °C and is a liquid at room temperature. Iodine trichloride melts at 101 °C.
Interhalogen | Properties
In chemistry, an ionic compound is a chemical compound composed of ions held together by electrostatic forces termed ionic bonding. The compound is neutral overall, but consists of positively charged ions called cations and negatively charged ions called anions. These can be simple ions such as the sodium (Na) and chloride (Cl) in sodium chloride, or polyatomic species such as the ammonium () and carbonate () ions in ammonium carbonate. Individual ions within an ionic compound usually have multiple nearest neighbours, so are not considered to be part of molecules, but instead part of a continuous three-dimensional network, usually in a crystalline structure.
Ionic compound | Introduction
In organic chemistry, Carnelley's rule, established in 1882 by Thomas Carnelley, states that "high molecular symmetry is associated with high melting point". Carnelley based his rule on examination of 15,000 chemical compounds. For example, for three structural isomers with molecular formula CH the melting point increases in the series isopentane −160 °C (113 K) n-pentane −129.8 °C (143 K) and neopentane −16.4 °C (256.8 K). Likewise in xylenes and also dichlorobenzenes the melting point increases in the order meta, ortho and then para. Pyridine has a lower symmetry than benzene hence its lower melting point but the melting point again increases with diazine and triazines.
Melting point | Carnelley's rule
Data obtained from "CRC Handbook of Chemistry (Page 2117)" Azeotropic mixture
Ethanol (data page) | Boiling points of aqueous solutions
Sodium peroxide crystallizes with hexagonal symmetry. Upon heating, the hexagonal form undergoes a transition into a phase of unknown symmetry at 512 °C. With further heating above the 657 °C boiling point, the compound decomposes to NaO, releasing O.
Sodium peroxide | Properties
The distance between two ions in an ionic crystal can be determined by X-ray crystallography, which gives the lengths of the sides of the unit cell of a crystal. For example, the length of each edge of the unit cell of sodium chloride is found to be 564.02 pm. Each edge of the unit cell of sodium chloride may be considered to have the atoms arranged as Na∙∙∙Cl∙∙∙Na, so the edge is twice the Na-Cl separation. Therefore, the distance between the Na and Cl ions is half of 564.02 pm, which is 282.01 pm. However, although X-ray crystallography gives the distance between ions, it doesn't indicate where the boundary is between those ions, so it doesn't directly give ionic radii.
Ionic radius | Determination
The chloride of N-methylpyridinium behaves as an ionic liquid in the molten state. Its properties with different mixtures of zinc chloride have been characterised by several authors in the temperature range 150 – 200 °C (423 – 473 K).
Methylpyridinium | Ionic liquid
For a solid crystalline ionic compound the enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born–Haber cycle. It can also be calculated (predicted) using the Born–Landé equation as the sum of the electrostatic potential energy, calculated by summing interactions between cations and anions, and a short-range repulsive potential energy term. The electrostatic potential can be expressed in terms of the interionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The further away from the nucleus the weaker the shield.
Ionic bonding | Strength of the bonding
Four crystalline forms (polymorphs) of ZnCl are known: α, β, γ, and δ, and in each case the Zn ions are tetrahedrally coordinated to four chloride ions. Here "a", "b", and "c" are lattice constants, "Z" is the number of structure units per unit cell, and ρ is the density calculated from the structure parameters. The pure anhydrous orthorhombic form (δ) rapidly changes to one of the other forms on exposure to the atmosphere, and a possible explanation is that the OH ions originating from the absorbed water facilitate the rearrangement. Rapid cooling of molten ZnCl gives a glass. The covalent character of the anhydrous material is indicated by its relatively low melting point of 275 °C.
Zinc chloride | Structure and properties
The boiling point is 45.5 °C, and the melting point is -122.9. The predicted heat of formation for liquid trimethyldiborane is ΔH=-48 kcal/mol, and for the gas -41 kcal/mol. Heat of vapourisation ΔH was measured at 7.0 kcal/mol. A gas chromatograph can be used to determine the amounts of the methyl boranes in a mixture. The order they pass through are diborane, monomethyldiborane, trimethylborane, 1,1-dimethyldiborane, 1,2-dimethyldiborane, trimethyldiborane, and last tetramethyldiborane.
Trimethyldiborane | Properties
The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor. The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water boils at at sea level, but at at altitude. For a given pressure, different liquids will boil at different temperatures.
Boiling point | Introduction
The attraction between the Na and Cl ions in the solid is so strong that only highly polar solvents like water dissolve NaCl well. When dissolved in water, the sodium chloride framework disintegrates as the Na and Cl ions become surrounded by the polar water molecules. These solutions consist of metal aquo complex with the formula [Na(HO)], with the Na–O distance of 250 pm. The chloride ions are also strongly solvated, each being surrounded by an average of 6 molecules of water. Solutions of sodium chloride have very different properties from pure water. The freezing point is for 23.31 wt% of salt, and the boiling point of saturated salt solution is near .
Sodium chloride | Aqueous solutions
The phase diagram on the right shows a "negative" azeotrope of ideal constituents, X and Y. Again the bottom trace illustrates the boiling temperature at various compositions, and again, below the bottom trace the mixture must be entirely liquid phase. The top trace again illustrates the condensation temperature of various compositions, and again, above the top trace the mixture must be entirely vapor phase. The point, A, shown here is a boiling point with a composition chosen very near to the azeotrope. The vapor is collected at the same temperature at point B. That vapor is cooled, condensed, and collected at point C.
Azeotrope | "Maximum-boiling" or "Negative" azeotrope
In even more simple terms a simple ionic compound (with positive and negative ions) such as sodium chloride (common salt) is easily soluble in a highly polar solvent (with some separation of positive (δ+) and negative (δ-) charges in the covalent molecule) such as water, as thus the sea is salty as it accumulates dissolved salts since early geological ages. The solubility is favored by entropy of mixing (Δ"S") and depends on enthalpy of dissolution (Δ"H") and the hydrophobic effect. The free energy of dissolution (Gibbs energy) depends on temperature and is given by the relationship: Δ"G" = Δ"H" – TΔ"S".
Solubility | Polarity
The quoted boiling point in the table below is thus a prediction. Despite this situation, some sources describe it without comment as a known compound.
Hexafluoride | Binary hexafluorides
For ionic compounds with ions occupying lattice sites with crystallographic point groups "C", "C""", "C" or "C" ("n" = 2, 3, 4 or 6) the concept of the lattice energy and the Born–Haber cycle has to be extended. In these cases the polarization energy "E" associated with ions on polar lattice sites has to be included in the Born–Haber cycle and the solid formation reaction has to start from the already polarized species. As an example, one may consider the case of iron-pyrite FeS, where sulfur ions occupy lattice site of point symmetry group "C". The lattice energy defining reaction then reads where pol S stands for the polarized, gaseous sulfur ion.
Lattice energy | Effect of polarisation
The boiling point is thus about 85 K lower than that of methanol. This fact can be explained by the absence of intramolecular H—F bonds, which are also not visible in the infrared gas phase spectrum.
Trifluoromethanol | Synthesis
An alternative to using HF and HSO as alkylation catalysts is the use of ionic liquid (IL). ILs are liquid salts that have melting points below 100 °C. They exhibit strong acid properties, so they can be used as acid catalysis without using conventional liquid acids. Ionic liquids are salts in liquid state, composed mostly of ions that convert C4 paraffins and other olefins into excellent gasoline-range blending products. Many parameters are available for fine-tuning IL properties for specific applications, and the choice of cation and anion affects the IL's physical properties, such as melting point, viscosity, density, water solubility, and reactivity.
Alkylation unit | Ionic liquids
TCL is a white crystalline solid at room temperature, soluble in common organic solvents. Its melting point is 81.5-83 °C, its boiling point is 265 °C. It is corrosive.
Terephthaloyl chloride | Properties
For cooling applications, a mixture with common sodium chloride salt (NaCl) lowers the melting point to 18 °C (64 °F). The heat of fusion of NaCl·NaSO·10HO, is actually "increased" slightly to 286 kJ/kg.
Sodium sulfate | Thermal storage
Ionic liquids are often moderate to poor conductors of electricity, non-ionizing, highly viscous and frequently exhibit low vapor pressure. Their other properties are diverse: many have low combustibility, are thermally stable, with wide liquid regions, and favorable solvating properties for a range of polar and non-polar compounds. Many classes of chemical reactions, such as Diels-Alder reactions and Friedel-Crafts reactions, can be performed using ionic liquids as solvents. ILs can serve as solvents for biocatalysis. The miscibility of ionic liquids with water or organic solvents varies with side chain lengths on the cation and with choice of anion. They can be functionalized to act as acids, bases, or ligands, and are precursors salts in the preparation of stable carbenes.
Ionic liquid | Characteristics
This process occurs widely in nature, and is the means of formation of the evaporite minerals. Another method of recovering the compound from solution involves saturating a solution at high temperature and then reducing the solubility by reducing the temperature until the solution is supersaturated and the solid compound nucleates. Insoluble ionic compounds can be precipitated by mixing two solutions, one with the cation and one with the anion in it. Because all solutions are electrically neutral, the two solutions mixed must also contain counterions of the opposite charges. To ensure that these do not contaminate the precipitated ionic compound, it is important to ensure they do not also precipitate.
Ionic compound | Formation
In a binary boiling-point diagram, temperature () is graphed vs. . At any given temperature where both phases are present, vapor with a certain mole fraction is in equilibrium with liquid with a certain mole fraction. The two mole fractions often differ. These vapor and liquid mole fractions are represented by two points on the same horizontal isotherm (constant ) line. When an entire range of temperatures vs. vapor and liquid mole fractions is graphed, two (usually curved) lines result. The lower one, representing the mole fraction of the boiling liquid at various temperatures, is called the "bubble point curve".
Vapor–liquid equilibrium | Boiling-point diagrams
Values of the ebullioscopic constants "K" for selected solvents:
Boiling-point elevation | Ebullioscopic constants