Answer
Verified
Hint: Carboxylic acid is made up of the carbonyl group $\text{(-C=O)}$ and the hydroxyl group$\text{-OH}$. The carboxylic acid may be aliphatic $\text{R-COOH}$ or aromatic $\text{Ar-COOH}$ depending upon the attachment of the group to the aliphatic or aryl group respectively.
Complete answer:
> The carboxylic acid has a higher boiling point. The higher boiling point is due to the presence of intermolecular hydrogen bonding. The hydrogen bond is the stronger bond
between the electronegative atoms like $\text{O,S,F,etc}$ between the $\text{H}$ atoms.
> In carboxylic acid, the hydrogen bonding is established between the hydrogen of one carboxylic acid and the oxygen of the carbonyl group of the other carboxylic acid.
> Due to hydrogen bonding, the carboxylic acid exists as dimers. The hydrogen bonds are not broken easily and completely even in the vapour phase. Thus they require higher heat energy to break or separate the molecules from each
other.
> On comparing it with the corresponding alcohol, carboxylic acid has a higher boiling than the corresponding aldehyde, ketone, and even alcohols of the comparable molecular masses.
Let's take an example of ethanoic acid and propanol. They have the same molecular mass 60. The boiling point of the ethanoic acid is $\text{391K}$ whereas that of the propanol is $\text{370K}$.
The higher boiling points of carboxylic acid compared to alcohol is due to the following reasons:
(i)
As compared to alcohols, the $\text{O-H}$ bond in carboxylic acids is more strongly polarised due to the presence of adjacent electron-withdrawing carbonyl groups. Therefore, carboxylic acids can form a stronger hydrogen bond.
(ii) The molecules of carboxylic acids are held together by two hydrogen bonds and therefore, form cyclic dimers.
We can say that the carboxylic acid molecules are held together by the strong attractive forces and therefore, they have a boiling point.
Such types of attractive forces do not exist between the alcohols. Alcohols do have the hydrogen bonding in it. But the hydrogen bonding is not as strong as that of the carboxylic acid.
Note: Remember that the boiling point is the heat energy required to separate the molecules. The boiling point is directly proportional to the molecular mass. Hydrogen and oxygen forms a hydrogen bonding. Hydrogen bonding is the reason behind the various properties.
Acidity
The most important property of carboxylic acids, and the one that is responsible for naming them such, is their acidity. An acid is any compound that donates a hydrogen ion, H+ (also called a proton), to another compound, termed a base. Carboxylic acids do this much more readily than most other classes of organic compounds, so they are said to be stronger acids, even though they are much weaker than the most important mineral acids—sulfuric (H2SO4), nitric (HNO3), and hydrochloric (HCl). The reason for the enhanced acidity of this group of compounds can best be demonstrated by a comparison of their acidity with that of alcohols, both of which contain an ―OH group. Alcohols are neutral compounds in aqueous solution. When an alcohol donates its proton, it becomes a negative ion called an alkoxide ion, RO−. When a carboxylic acid donates its proton, it becomes a negatively charged ion, RCOO−, called a carboxylate ion.
A carboxylate ion is much more stable than the corresponding alkoxide ion because of the existence of resonance structures for the carboxylate ion which disperse its negative charge. Only one structure can be drawn for an alkoxide ion, but two structures can be drawn for a carboxylate ion. When two or more structures that differ only in the positions of valence electrons can be drawn for a molecule or ion, it means that its valence electrons are delocalized, or spread over more than two atoms. This phenomenon is called resonance, and the structures are called resonance forms. A double-headed arrow is used to show that the two or more structures are related by resonance. Because there are two resonance forms but only one real ion, it follows that neither of these forms is an accurate representation of the actual ion. The real structure incorporates aspects of both resonance structures but duplicates neither. Resonance always stabilizes a molecule or ion, even if charge is not involved. The stability of an anion determines the strength of its parent acid. A carboxylic acid is, therefore, a much stronger acid than the corresponding alcohol, because, when it loses its proton, a more stable ion results.
Some atoms or groups, when attached to a carbon, are electron-withdrawing, as compared with a hydrogen atom in the same position. For example, consider chloroacetic acid (Cl―CH2COOH) compared with acetic acid (H―CH2COOH). Because chlorine has a higher electronegativity than hydrogen, the electrons in the Cl―C bond are drawn farther from the carbon than the electrons in the corresponding H―C bond. Thus, chlorine is considered to be an electron-withdrawing group. This is one example of the so-called inductive effect, in which a substituent affects a compound’s distribution of electrons. There are a number of such effects, and atoms or groups may be electron-withdrawing or electron-donating as compared with hydrogen. The presence of such groups near the COOH group of a carboxylic acid often has an effect on the acidity. In general, electron-withdrawing groups increase acidity by increasing the stability of the carboxylate ion. In contrast, electron-donating groups decrease acidity by destabilizing the carboxylate ion. For example, the methyl group, ―CH3, is generally regarded as electron-donating, and acetic acid, CH3 COOH, is about 10 times weaker as an acid than formic acid, HCOOH. Similarly, chloroacetic acid, ClCH2 COOH, in which the strongly electron-withdrawing chlorine replaces a hydrogen atom, is about 100 times stronger as an acid than acetic acid, and nitroacetic acid, NO2CH2 COOH, is even stronger. (The NO2 group is a very strong electron-withdrawing group.) An even greater effect is found in trichloroacetic acid, Cl3CCOOH, whose acid strength is about the same as that of hydrochloric acid.
Solubility
The solubility of carboxylic acids in water is similar to that of alcohols, aldehydes, and ketones. Acids with fewer than about five carbons dissolve in water; those with a higher molecular weight are insoluble owing to the larger hydrocarbon portion, which is hydrophobic. The sodium, ammonium, and potassium salts of carboxylic acids, however, are generally quite soluble in water. Thus, almost any carboxylic acid can be made to dissolve in water by converting it to such a salt, which is easily done by adding a strong base—most commonly sodium hydroxide (NaOH) or potassium hydroxide, (KOH). The calcium and sodium salts of propanoic (propionic) acid are used as preservatives, chiefly in cheese, bread, and other baked goods.
Boiling point
Carboxylic acids have much higher boiling points than hydrocarbons, alcohols, ethers, aldehydes, or ketones of similar molecular weight. Even the simplest carboxylic acid, formic acid, boils at 101 °C (214 °F), which is considerably higher than the boiling point of ethanol (ethyl alcohol), C2H5OH, which boils at 78.5 °C (173 °F), although the two have nearly identical molecular weights. The difference is that two molecules of a carboxylic acid form two hydrogen bonds with each other (two alcohol molecules can only form one). Thus, carboxylic acids exist as dimers (pairs of molecules), not only in the liquid state but even to some extent in the gaseous state.
Therefore, boiling a carboxylic acid requires the addition of more heat than boiling the corresponding alcohol, because (1) if the dimer persists in the gaseous state, the molecular weight is in effect doubled; and, (2) if the dimer is broken upon boiling, extra energy is required to break the two hydrogen bonds. Carboxylic acids with higher molecular weights are solids at room temperature (e.g., benzoic and palmitic acids). Virtually all salts of carboxylic acids are solids at room temperature, as can be expected for ionic compounds.
Odour
Unbranched-chain carboxylic acids (fatty acids) that are liquids at room temperature, especially those from propanoic (C3) to decanoic (C10) acid, have very foul, disagreeable odours. An example is butanoic (butyric) acid (C4), which is the main ingredient in stale perspiration and thus the chief cause of “locker-room” odour.