Contents
II. Factors Affecting the Acidity of Organic Compounds
II.2. Electronegativity and Bond Strength (Size)
II.2.1. Across the Periodic Table, Acidity is Influenced by Electronegativity
II.2.2. Down the Periodic Table, Acidity is Influenced by Bond Strength
II.3.1. Electron-withdrawing Inductive Effect
II.3.2. Electron-donating Inductive Effect
II.4. Resonance (Delocalisation)
II.4.1. Resonance in Carboxylic Acids
II.4.2. Acidity Assisted by Resonance to Benzene Ring
II.4.3. Acidity Assisted by Resonance to Neighbouring Unsaturated Groups
I. Introduction
Now that we have seen the pKa values of several functional group and know what they represent in general, we shall have a look at more specific compounds and trends to see what factors affect the acidity of organic compounds.
As you have seen in the pKa tables, those values are approximate values for a group of compounds with the same functional group. For example, alcohols have pKa approximately 16. However, not all of them have pKa of 16, in fact only methanol does. Ethanol’s pKa is about 16.5, and t-butanol’s pKa is about 19.
So what makes them different? Let’s start our discussion!
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II. Factors Affecting the Acidity of Organic Compounds
You may have noticed that there are two types of acids in general:
- Neutral acids
Are acids with no overall charge, such as HCl, HBr, CH3COOH, PhCOOH (benzoic acid), etc.
These acids release anionic conjugate bases: negatively charged conjugate bases, such as Cl–, Br–, CH3COO–, PhCOO– (benzoate anion), etc.
- Cationic acids
Are positively charged acids. Usually the acids in this class are protonated nitrogens or oxygens, such as NH4+ (ammonium ion), Et3NH+ (triethylammonium ion), C5H5NH+ (pyridinium ion), H3O+ (hydronium ion), and protonated carbonyls (R2C=OH+) and alcohols (R–OH2+).
These acids release neutral conjugate bases: conjugate bases with no overall charge, such as NH3 (ammonia), Et3N (triethylamine), C5H5N (pyridine), H2O (water), and neutral carbonyls (R2C=O) and alcohols (R–OH).
We have also learnt what differs strong acids to weak acids. A strong acid is an acid that releases weak (stable) conjugate base, whilst a weak acid is an acid that releases strong (unstable) conjugate base.
All of these can be translated into two perspectives, depending on the acids:
- Neutral acids
Neutral acids release anionic conjugate bases, which may or may not be stable.
If the anionic conjugate bases are stable, then the acids are strong.
If the anionic conjugate bases are unstable, then the acids are weak.
Therefore, to predict the acidity of neutral acids, we have to examine the stability of their conjugate bases.
- Cationic acids
Cationic acids release neutral conjugate bases. Neutral conjugate bases are generally stable. Cationic acids however, may or may not be stable.
If the cationic acids are unstable, they give away their proton easily, hence they are strong acids.
If the cationic acids are stable, they don’t give away their proton easily, hence they are weak acids.
Therefore, to predict the acidity of cationic acids, we have to examine their own stability.
That being said, factors that affect the acidity of organic compounds can simply be thought of as factors that stabilises the charged species (either the anionic conjugate base, or cationic acids).
If the charged species is stable, then we have a strong acid; if not, then we have a weak acid.
Therefore, the question that we need to address is:
What are the factors that can make the charged species more stable?
In this post, we will look at the affecting factors one by one, and prove them by comparing their pKa values. Remember:
The smaller the pKa value of an acid is, the stronger the acid is, the weaker its conjugate base is.
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II.1. Hybridisation
This is the most common starting point when talking about the factors affecting the acidity of organic compounds. In organic chemistry, you only get three types of hybridisation: sp3 (for single bonds), sp2 (for double bonds), and sp (for triple bonds).
Let’s compare the acidity between ethane (an alkane), ethylene (an alkene), and acetylene (an alkyne).
You can see that an alkene is about 7 orders of magnitude (107 times) more acidic than an alkane, and an alkyne is about 15 orders of magnitude (1015 times) more acidic than alkane!
Why is this? Obviously, hybridsation. Otherwise I won’t put that example in this section!
Alkyne (sp hybridisation) have more s character than alkene (sp2) and alkane (sp3). So what is this s character? Think about it like this:
- In sp3, you have 1 s orbital and 3 p orbitals combining (4 orbitals overall): s, p, p, p
This means, the s orbital makes up about 25% (1 out of 4 orbitals) of the overall
- In sp2, you have 1 s orbital and 2 p orbitals combining (3 orbitals overall): s, p, p
This means, the s orbital makes up about 33% (1 out of 3 orbitals) of the overall
- In sp, you have 1 s orbital and 1 p orbitals combining (2 orbitals overall): s, p
This means, the s orbital only makes up about 50% (1 out of 2 orbitals) of the overall
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The Logic Box Why does having more s character means more acidic? Recall from your general chemistry or physical chemistry lectures that the s orbital is spherical around the centre of the atom (nucleus). On the other hand, each of the p orbitals (px, py, and pz) has two lobes with a node near the nucleus. A node is a part of the orbital that has no electron density.
This means that:
Therefore, a negative charge can be stabilised better by the s orbital rather than by the p orbital. This is why more s character means more negative charge stabilisation. |
The three examples above are neutral acids that release anionic conjugate base. The negative charge of the conjugate base is stabilised better as the s character increases.
This makes sp2 hybridisation better at stabilising negative charge than sp3 hybridisation (alkene is more acidic than alkane). This also makes sp hybridisation (alkyne) the most acidic amongst the three.
Of course, this is also evident in the nitrogen analogues:
The three examples in the nitrogen analogues are (from left to right) amine, imine, and protonated nitrile. Amine and imine are neutral acids but protonated nitrile is a cationic acid. Therefore, the theory above can easily be applied to amine and imine, but what about protonated nitrile?
Remember that the nitrile is sp hybridised, and that more s character means more negative charge stabilisation. The protonated nitrile has a positive charge and therefore gets destabilised by the s character. Because the cationic acid is destabilised, it releases its proton easily and becomes a strong acid, as shown by its pKa value.
This whole theory is also evident in the oxygen analogues:
Isopropanol is a neutral acid while the protonated acetone is a cationic acid. The conjugate base of isopropanol does not get much negative charge stabilisation (25% s character due to sp3 hybridisation) and therefore isopropanol is a weak acid. The protonated acetone has sp2 hybridisation, therefore its positive charge is destabilised by the s character and it becomes a strong acid.
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II.2. Electronegativity and Bond Strength (Size)
Let’s first recall from your general chemistry (or high school chemistry?) lecture:
- Electronegativity increases from left to right
- Electronegativity decreases from top to bottom
- Atomic radius (size) decreases from left to right causing bond strength to increase
- Atomic radius (size) increases from top to bottom causing bond strength to decrease
Consult your general chemistry textbooks if you forget about this, I will not go into too much detail here.
Now, have a look at the trends in acidity within the periodic table:
Some of you may ask yourself this: ‘acidity increases as electronegativity increases when we go across (left to right) the periodic table, but acidity also increases as electronegativity decreases when we go down (top to bottom) the periodic table. Why??’
Others may also ask this: ‘acidity increases as size increases when we go down (top to bottom) the periodic table, but acidity also increases as size decreases when we go across (left to right) the periodic table. Why??’
You may also ask the other way around for each case, but regardless how you ask your question, this seemingly contradictory trend can actually be explained easily:
- Across the periodic table, acidity is influenced by electronegativity
- Down the periodic table, acidity is influenced by bond strength
Let’s have a look at each of the two in more detail, and let’s start with the easy one first.
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II.2.1. Across the Periodic Table, Acidity is Influenced by Electronegativity
‘As we go across the periodic table, acidity increases as electronegativity increases’
This makes sense, because the increase in electronegativity means the species is able to stabilise negative charge better. Compare the following acids from the first row:
The acids are neutral acids, therefore they release anionic conjugate base. The conjugate base gets more and more stabilised by the increased electronegativity. In other words, negatively charged species gets more and more stabilised as its electronegativity increases. Therefore, the acid gets more and more stable as we go across the periodic table.
The bond strength does not significantly affect the acidity when you go across the periodic table, because the elements in the same row have similar energy level and overlap of orbitals between H and the element itself. This will be clear when we talk about it in the next section.
One last point to note: you’ve always been told that HF is a weak acid. Yes it is true when we compare it against HCl, HBr, or HI. But here, when we compare it to the other acids within the same row, HF is the strongest. Just look at its small pKa value compared to its neighbours.
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II.2.2. Down the Periodic Table, Acidity is Influenced by Bond Strength
‘As we go down the periodic table, acidity increases as bond strength decreases’
As you go down the periodic table, the atomic radius (size) becomes larger. This means that the orbital overlap between the element and H gets more and more ineffective. In turn, this causes the acid to release its proton more easily.
Compare the following acids from the halogen group:
The increase in size and the decrease in bond strength affect the acidity more than the increase in electronegativity as we go down the periodic table. Therefore HI becomes the strongest acid amongst the other acids in its group.
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II.3. Inductive Effect
In the section above, we saw how electronegativity affects the acidity of the H+ directly bonded to the electronegative atom. Here, we shall see that an electronegative functional group can also affect the acidity of a H+ that is several bonds away from itself. We shall also see a non-electronegative functional group that affect the acidity of the aforementioned H+. Both of these groups affect the acidity of the H+ through the σ bonds. This effect is known as the inductive effect.
There are two types of inductive effect based on the groups that exhibit the effect:
- Electron-withdrawing inductive effect
- Electron-donating inductive effect
Substituted carboxylic acids are well known examples for this effect. A carboxylic acid is a neutral acid; in solution, it dissociates to give proton and carboxylate anion as its conjugate base. If the carboxylate anion can be stabilised, the carboxylic acid gets stronger. We shall use carboxylic acids as examples in this section.
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II.3.1. Electron-withdrawing Inductive Effect
Electronegative functional groups affect the acidity of H+ by inductively pulling the electron density towards it through the σ bonds, therefore weaking the O–H bond in COOH and making the H+ easier to be released. They also stabilise negatively charged species, therefore the conjugate base is stabilised, making the carboxylic acid a stronger acid.
This type of group is known as the electron-withdrawing group (EWG).
The ‘electron-pulling through the σ bonds’ phenomenon is known as the electron-withdrawing inductive effect, also known as the ‘–I effect‘.
Have a look at the acetic acid derivatives below.
The electronegativity of the chlorine atom pulls electron density away from the acidic H towards itself, making chloroacetic acid about 4 orders of magnitude (104 times) more acidic than normal acetic acid! This shows the inductive effect that chlorine has.
More EWGs means more acidic
As the number of chlorine atom increases, the inductive effect becomes stronger. Hence, trichloroacetic acid is a stronger acid than dichloroacetic acid, which in turn is a stronger acid than (mono)chloroacetic acid. All of them are more acidic compared to unsubstituted acetic acid.
The addition of one chlorine atom to the acetic acid increases the acidity by 1.88 orders of magnitude (about 76x). Addition of another chlorine atom increases it further by about 1.57 orders of magnitude (about 37x), and addition of the last one increases it again by 0.64 orders of magnitude (about 4x).
More electronegative EWGs means more acidic
The electronegativity of the EWG itself also influences the inductive effect. Compare the acidity of trichloroacetic acid (TCA) and trifluoroacetic acid (TFA) below:
Fluorine is more electronegative than chlorine, therefore its inductive effect is greater than that of chlorine, causing the pKa of TFA to be smaller than TCA. This means that TFA is more acidic than TCA by 0.9 order of magnitude (about 8x).
Closer EWG means more acidic
One last thing, distance matters. The closer the EWG is to the COOH, the stronger the inductive effect is, and the more acidic the acid is.
You can see from the picture above that:
- 4-Chlorobutanoic acid is about 2x more acidic than unsubstituted butanoic acid.
- 3-Chlorobutanoic acid is about 3x more acidic than 4-chlorobutanoic acid.
- 2-chlorobutanoic acid is about 15x more acidic than 3-chlorobutanoic acid.
The chlorine atom in 2-chlorobutanoic acid is one atom closer to the carboxylic group than the chlorine atom in 3-chlorobutanoic acid. Look how dramatically the acidity increases!
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II.3.2. Electron-donating Inductive Effect
Alkyl functional group affects the acidity of H+ by inductively pushing the electron density away from it through the σ bonds, therefore strengthening the O–H bond in COOH and making the H+ more difficult to be released. They also stabilise positively charged species, therefore the conjugate base is destabilised, making the carboxylic acid a weaker acid.
This type of group is known as the electron-donating group (EDG).
The ‘electron-pushing through the σ bonds’ phenomenon is known as the electron-donating inductive effect, sometimes written as ‘+I effect‘.
Have a look at the acetic acid derivatives below.
The methyl groups push electron density away from themselves towards the acidic H. It destabilises the anionic conjugate bases, making pivalic acid (2,2-dimethylpropanoic acid) the least acidic and formic acid the most acidic in the series of carboxylic acid above.
Similar effect is also observed in alcohols:
Larger number of methyl groups (EDGs) increases the electron density near the acidic H and therefore making the alcohol less acidic. This causes the t-butanol to have to largest pKa value in the series above.
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II.4. Resonance (Delocalisation)
We have learnt from basic structural organic chemistry that resonance can greatly stabilise charged species. Because the acidity of organic compounds is influenced by the stability of charged species, resonance plays an important role in determining the acidity of various compounds.
Similar to the inductive effect, there are functional groups that are electron-withdrawing and electron-donating through resonance. Through resonance, EWGs increases the acidity of a compound and EDGs decreases the acidity.
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II.4.1. Resonance in Carboxylic Acids
The dissociation of carboxylic acid is stabilised by the resonance within the COO group.
This resonance stabilises the carboxylate anion (conjugate base of carboxylic acid), making carboxylic acids acidic. We can compare the pKas of ethanol and acetic acid to see how the resonance increases the acidity of carboxylic acids:
The resonance stabilises the anionic conjugate base so much so that it increases the acidity of acetic acid by about 11 orders of magnitude!
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II.4.2. Acidity Assisted by Resonance to Benzene Ring
When a benzene ring is attached to the carboxylic group, resonance of the carboxylate group with the ring increases its acidity. Compare acetic acid to benzoic acid:
Resonance with the benzene ring does not only increase the acidity of carboxylic acids but also other functional groups such as alcohol, thiol, and amine.
First, we compare a non-conjugated alcohol (methanol) with a conjugated alcohol (phenol):
The negative charge in the phenoxide anion (anionic conjugate base of phenol) is stabilised by delocalisation to the aromatic ring, whilst there is no conjugate base stabilisation in methanol. This makes phenol more acidic than methanol.
This also applies to other similar compounds such as thiols (methanethiol and thiophenol):
… and between methanamine and aniline too:
You can observe something interesting when comparing neutral and protonated aniline (anilinium cation):
In neutral aniline, the free electron pair of the nitrogen delocalises to the benzene ring, giving stability. Protonation of the aniline nitrogen takes away the conjugation, making the anilinium (which is a cationic acid) unstable. Instability of a cationic acid makes a strong acid. This, in conjunction with the stabilisation of aniline as a conjugate base, makes anilinium a strong acid, as proved by its smaller pKa value of 4.6, compared to aniline’s 28.
Of course, you observe the same thing in neutral and protonated phenol:
Protonation of the phenol oxygen takes away the conjugation that phenol had. This makes the protonated phenol (a cationic acid) unstable, hence making it a stronger acid than neutral phenol by 17 orders of magnitude!
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II.4.3. Acidity Assisted by Resonance to Neighbouring Unsaturated Groups
Hydrogens bonded to a carbon adjacent to one or more unsaturated groups are acidic. This carbon is known as the α-carbon, whilst its hydrogen is known as the α-hydrogen. Deprotonation of α-hydrogens results in the formation of carbanion, called the α-carbanion. Overall, this type of acid is known as carbon acid.
The reason α-hydrogens are acidic is because their conjugate bases are stabilised through conjugation with the unsaturated groups. For example, in butanone:
You’ll probably realise that the carboxylic acid family (acids, esters, and amides) give larger pKa, meaning they are less acidic than aldehydes/ketones. The reason for this is because the OR and NR2 groups in the carboxylic acid family is more involved in resonance with the carbonyl group. Because of this, the α-carbanion has to ‘compete’ and ‘share’ the resonance with the carbonyl group. This makes the α-hydrogen of the carboxylic acid family more acidic than the aldehydes/ketones.
In aldehydes/ketones, the the α-carbanion has carbonyl group only for itself, making the resonance stabilision better. This makes their α-hydrogens more acidic.
The α-hydrogen is even more acidic if it’s flanked by two adjacent unsaturated groups. The figure below shows the comparison of pKas of (from left to right): 1,3-diester (malonate esters), dicyanomethane (malononitrile), β-keto ester, 1,3-diketone, and dinitromethane.
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II.5. Aromaticity
We know that there are aromatic and anti aromatic compounds. Aromaticity stabilises compounds and therefore making acids stronger. If a compound releases an aromatic conjugate base, the aromaticity stabilises the conjugate base, which make the compound a stronger acid. If the released conjugate base is anti aromatic, the anti aromaticity destabilises the conjugate base, making the compound a weaker acid.
The conjugate base of cyclopropene is anti aromatic and therefore very unstable. On the other hand, the conjugate base of cycloheptatriene is not planar, making it neither aromatic nor anti aromatic, and therefore unstable. This makes the pKa value of both compounds very large. Having an anti aromatic conjugate base, cyclopropene is therefore the most unstable of the two. Therefore, cycloheptatriene has smaller pKa than cyclopropene.
On the other hand, the conjugate base of cyclopentadiene is aromatic, making it very stable and dramatically more acidic than the other two. It is about 23 orders of magnitude more acidic than cycloheptatriene and 45 orders of magnitude more acidic than cyclopropene!
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II.6. Solvation
A species is considered an acid if it is able to separate its proton from its conjugate base. In order to do that, solvent plays a great role. In the absence of solvent, it is hard to separate a positive charge (proton) from a negative charge (conjugate base). This separation is greatly assisted by solvent; the solvent molecules orient themselves around the solute, separating the oppositely charged species. This process is known as solvation.
Basically, it is difficult to achieve separation of proton at room temperature without the help of solvent, so the role of solvent is very crucial. In fact, all of the factors explained in Sections II.1.-II.5. are also influenced by solvation. All of the compounds used as examples in the previous sections may become more or less acidic when dissolved in different solvent.
Solvent that can better stabilise the conjugate base of a neutral acid makes it a stronger acid. The ability of the solvent to stabilise the conjugate base depends on many things. One of which is the molecular structure of the acid/conjugate base itself.
We’ve used these examples in Section II.3.2. previously, but they are good examples. Compare the pKa of several decreasingly-bulky alcohols and carboxylic acids:
The bulkiness of the methyl groups hampers the solvation of the conjugate bases by solvent molecules. This destabilises the conjugate bases, making the compounds increasingly less acidic as bulkiness increases. Therefore, aside from the effect of EDG in those examples, solvent also affects their acidity.
The solvent itself is also an important factor. The pKa values I have shown you in the pKa table are mostly measured relative to water as a solvent. The pKa of some compounds are also measured in (dimethylsulfoxide) DMSO, and as you may have guessed, their values are different to those measured in water. Here are a few examples:
| Acid Name |
Acid Formula |
pKa (water) | pKa (DMSO) |
| Trifluorosulfonic acid |
CF3SO3H | –14 | 0.3 |
| Hydrobromic acid |
HBr | –9 | 0.9 |
| Hydrochloric acid |
HCl | –7 | 1.8 |
| Hydrobromic acid |
HF | 3.2 | 15 |
| Acetic acid | CH3COOH | 4.74 | 12.3 |
| Phenol | C6H5OH | 10 | 18 |
| Methanol | CH3OH | 15.5 | 28 |
| Water | H2O | 15.7 | 32 |
| Ammonia | NH3 | 38 | 41 |
You can immediately see that the pKa values in DMSO are larger than the pKa values in water. This means that in DMSO, the acids are less acidic. Some are monumentally larger (e.g. trifluorosulfonic acid is almost 14 orders of magnitude less acidic in DMSO than in water), and some are not (e.g. ammonia is only 3 orders of magnitude less acidic in DMSO than in water).
As an organic solvent, DMSO is not as good as water in stabilising the anionic conjugate base. Therefore, the acids are less acidic.
We will talk more about acidity, basicity, and solvent, including the levelling effect in Acid-Base VIII, Acid-Base IX, and Acid-Base X. Keep in mind that what we have talked about here is about the solvation of certain species in certain solvents.
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III. Conclusion
We have seen various factors that influence the acidity of organic compounds. We have also proved the theories by looking at the pKas of the compounds we used as examples.
- Hybridisation
The sp hybridisation has 50% s character, which stabilises the negative charge of a species. We get stronger acids as we go from sp3 to sp2 to sp hybridisations.
- Electronegativity
As we go across the periodic table, the electronegativity effect outweighs the bond strength effect. We get stronger acids as we go from left to right due to the increase of electronegativity.
- Bond strength
As we go down the periodic table, the effect of bond strength become more significant compared to electronegativity. We get stronger acids due to the decrease in bond strength.
- Inductive effect
Electronegative atoms are electron-withdrawing groups. They pull electron density towards themselves, making the conjugate base of an acid more stable, and therefore making the acid stronger.
Electron-donating group pushes the electron density away from themselves, makin the conjugate base of an acid less stable, and therefore making the acid weaker.
- Resonance
Resonance stabilises charged species. The more they’re stabilised by resonance, the more acidic they become.
- Aromaticity
Aromatic conjugate bases are stable and make for stronger acids. Anti aromatic conjugate bases are unstable and make for weaker acids.
- Solvation
Conjugated bases that can be well solvated by the solvent molecules give stronger acids. Bulkier molecuels are more difficult to be solvated and therefore less acidic.
In solvents that are less good in stabilising anionic conjugate base, acids are less acidic.
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Page last updated: 06ii17
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