In order to predict the solubilities of the given compounds, it is useful to define the primary intermolecular forces each experiences when introduced to water. I: Hydrogen bonding dominates interaction between methanol and water the two are miscible. II: Hydrogen bonding is present, but solubility is reduced by the presence of a multi-carbon chain, which adds significant nonpolar character to the structure.
III: Perchloric acid is a strong acid stronger than nitric acid and sulfuric acid , meaning it completely dissociates in water, forming very strong ion-dipole interactions with water. The molecule with the lowest vapor pressure is the molecule with the strongest intermolecular forces. All of these molecules except pentane have the capability to hydrogen bond. However, can donate two hydrogen bonds one at each alcohol , and can accept four hydrogen bonds one at each oxygen.
All oxygen, fluorine, and nitrogen atoms are hydrogen bond acceptors, whether or not they are attached to hydrogen. However, a hydrogen bond donor is a hydrogen that must be bound to any of these three atoms. Of the following intermolecular forces, which force would typically provide a pure compound with the highest possible boiling point? At first glance, we would be eager to jump to ionic bonding as the correct answer, as ionic bonding provides for very high boiling points.
The correct answer, however, is a rare type of intermolecular force called network covalent bonding. Network covalent bonding is typically seen in diamond and quartz, and is a stronger intermolecular force than ionic bonding. Hydrogen bonding is the next strongest intermolecular force and also increases the boiling points of pure substances.
In general, increased intermolecular interraction and higher magnitude of intermolecular forces lead to an increase in a molecule's boiling point. Inversely, decreased intermolecular interraction and lower magnitude of intermolecular forces lead to a decrease in a molecule's boiling point.
In this case, the only intermolecular force exhibited by any of these molecules are London dispersion forces. The magnitude of London dispersion forces decreases with a decrease in molecule size carbon chain length and molecular surface area. Therefore, the shortest, most branched molecule in this problem will have the lowest boiling point. The correct answer is isobutane, a four membered, branched hydrocarbon. When discussing boiling points of hydrocarbons, it is important to remember that branching decreases a molecule's boiling point.
We can first eliminate hexane and pentane as our answers, as neither are branched. From here, we can come upon 2,3-dimethylbutane as our answer because it is more branched than 2-methylpentane. Also important when ranking hydrocarbons in terms of boiling point is the number of carbons - more carbons means a higher boiling point.
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Louis, MO Subject optional. Email address: Your name:. R hydroxypentanoic acid. Explanation : Boiling point is highly dependent on the intermolecular forces of a compound. Report an Error. Possible Answers: The triple bond of an alkyne consists of three pi-bonds. Terminal alkynes are less acidic than internal alkynes. Internal alkynes are more stable than terminal alkynes. Terminal alkynes are stronger compounds than internal alkynes. Correct answer: Internal alkynes are more stable than terminal alkynes.
Explanation : The answer is "Internal alkynes are more stable than terminal alkynes" as it is the only true statement in regards to alkynes. Since all observable samples of compounds and mixtures contain a very large number of molecules ca.
Indeed, many of the physical characteristics of compounds that are used to identify them e. All atoms and molecules have a weak attraction for one another, known as van der Waals attraction. This attractive force has its origin in the electrostatic attraction of the electrons of one molecule or atom for the nuclei of another.
If there were no van der Waals forces, all matter would exist in a gaseous state, and life as we know it would not be possible. It should be noted that there are also smaller repulsive forces between molecules that increase rapidly at very small intermolecular distances. For general purposes it is useful to consider temperature to be a measure of the kinetic energy of all the atoms and molecules in a given system. As temperature is increased, there is a corresponding increase in the vigor of translational and rotation motions of all molecules, as well as the vibrations of atoms and groups of atoms within molecules.
Experience shows that many compounds exist normally as liquids and solids; and that even low-density gases, such as hydrogen and helium, can be liquified at sufficiently low temperature and high pressure. A clear conclusion to be drawn from this fact is that intermolecular attractive forces vary considerably, and that the boiling point of a compound is a measure of the strength of these forces. Thus, in order to break the intermolecular attractions that hold the molecules of a compound in the condensed liquid state, it is necessary to increase their kinetic energy by raising the sample temperature to the characteristic boiling point of the compound.
The following table illustrates some of the factors that influence the strength of intermolecular attractions. The formula of each entry is followed by its formula weight in parentheses and the boiling point in degrees Celsius. First there is molecular size. Large molecules have more electrons and nuclei that create van der Waals attractive forces, so their compounds usually have higher boiling points than similar compounds made up of smaller molecules.
It is very important to apply this rule only to like compounds. The examples given in the first two rows are similar in that the molecules or atoms are spherical in shape and do not have permanent dipoles.
Molecular shape is also important, as the second group of compounds illustrate. The upper row consists of roughly spherical molecules, whereas the isomers in the lower row have cylindrical or linear shaped molecules. The attractive forces between the latter group are generally greater. Finally, permanent molecular dipoles generated by polar covalent bonds result in even greater attractive forces between molecules, provided they have the mobility to line up in appropriate orientations.
The last entries in the table compare non-polar hydrocarbons with equal-sized compounds having polar bonds to oxygen and nitrogen. Halogens also form polar bonds to carbon, but they also increase the molecular mass, making it difficult to distinguish among these factors. The melting points of crystalline solids cannot be categorized in as simple a fashion as boiling points.
The distance between molecules in a crystal lattice is small and regular, with intermolecular forces serving to constrain the motion of the molecules more severely than in the liquid state.
Molecular size is important, but shape is also critical, since individual molecules need to fit together cooperatively for the attractive lattice forces to be large. Spherically shaped molecules generally have relatively high melting points, which in some cases approach the boiling point. This reflects the fact that spheres can pack together more closely than other shapes. This structure or shape sensitivity is one of the reasons that melting points are widely used to identify specific compounds.
The data in the following table serves to illustrate these points. Notice that the boiling points of the unbranched alkanes pentane through decane increase rather smoothly with molecular weight, but the melting points of the even-carbon chains increase more than those of the odd-carbon chains. Even-membered chains pack together in a uniform fashion more compactly than do odd-membered chains. Hydrogen Bonding. The most powerful intermolecular force influencing neutral uncharged molecules is the hydrogen bond.
This is shown graphically in the following chart. The exceptionally strong dipole-dipole attractions that cause this behavior are called the hydrogen bond.
Hydrogen forms polar covalent bonds to more electronegative atoms such as oxygen, and because a hydrogen atom is quite small, the positive end of the bond dipole the hydrogen can approach neighboring nucleophilic or basic sites more closely than can other polar bonds. Coulombic forces are inversely proportional to the sixth power of the distance between dipoles, making these interactions relatively strong, although they are still weak ca.
The unique properties of water are largely due to the strong hydrogen bonding that occurs between its molecules.
In the following diagram the hydrogen bonds are depicted as magenta dashed lines. The molecule providing a polar hydrogen for a hydrogen bond is called a donor. The molecule that provides the electron rich site to which the hydrogen is attracted is called an acceptor. Water and alcohols may serve as both donors and acceptors, whereas ethers, aldehydes, ketones and esters can function only as acceptors. Similarly, primary and secondary amines are both donors and acceptors, but tertiary amines function only as acceptors.
Once you are able to recognize compounds that can exhibit intermolecular hydrogen bonding, the relatively high boiling points they exhibit become understandable. The data in the following table serve to illustrate this point. Compound Formula Mol. Also, O—H O hydrogen bonds are clearly stronger than N—H N hydrogen bonds, as we see by comparing propanol with the amines.
As expected, the presence of two hydrogen bonding functions in a compound raises the boiling point even further. Acetic acid the ninth entry is an interesting case. A dimeric species, shown on the right, held together by two hydrogen bonds is a major component of the liquid state.
Thus, the dimeric hydrogen bonded structure appears to be a good representation of acetic acid in the condensed state. A related principle is worth noting at this point. Although the hydrogen bond is relatively weak ca. The hydrogen bonds between cellulose fibers confer great strength to wood and related materials. Properties of Crystalline Solids. Some decompose before melting, a few sublime, but a majority undergo repeated melting and crystallization without any change in molecular structure.
When a pure crystalline compound is heated, or a liquid cooled, the change in sample temperature with time is roughly uniform. However, if the solid melts, or the liquid freezes, a discontinuity occurs and the temperature of the sample remains constant until the phase change is complete.
For a given compound, this temperature represents its melting point or freezing point , and is a reproducible constant as long as the external pressure does not change. The length of the horizontal portion depends on the size of the sample, since a quantity of heat proportional to the heat of fusion must be added or removed before the phase change is complete.
Now it is well known that the freezing point of a solvent is lowered by a dissolved solute, e. This provides a useful means for establishing the identity or non-identity of two or more compounds, since the melting points of numerous solid organic compounds are documented and commonly used as a test of purity. The phase diagram on the right shows the melting point behavior of mixtures ranging from pure A on the left to pure B on the right.
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