When compounds with different functional groups are given how do we determine which is more acidic or basic
.
First I'll discuss what makes something acidic. I don't know your background in chemistry so feel free to breeze over this part.
The definition of a Brønsted acid is a species that has the tendency to lose a proton. When that proton is removed, you're left with the conjugate base and any features that stabilize that conjugate base result in a stronger acid. For example, CH4 (methane) has CH3- as its conjugate base and its pKa is ~48. HF (hydrofluric acid) has F- as its conjugate base, and a pKa of around 3. This is because fluorine is much more electronegative than carbon (i.e. it likes to pull electron density to itself) and so it can better stabilize the negative charge.
Another feature that can stabilize the conjugate base is the ability to delocalise that negative charge over more than one electronegative atom. This is easier to see on an image, using perchloric acid as an example (pKa about -10):
(from the learning point)
Electron diffraction studies show that the charge is spread out evenly over all oxygen atoms (better demonstrated in the final picture, with the dashed bonds). The bonds are all the same length - if the charge was not delocalised in this way, you would expect single and double bond lengths (which differ by around 20 picometres).
Having electronegative atoms directly attached to, or near, protons also makes them more acidic (as their electron density will be pulled towards that atom). An example of this is in carbonyl compounds:
(Alpha and beta carbon)
The carbon double bonded to the oxygen is the carbonyl compound, and the carbon directly bonded to it is known as the α-carbon. The hydrogen atoms attached to that carbon, then, are known as α-hydrogens (or α-protons). Their close proximity to the electron-withdrawing carbonyl group makes them more acidic, and α-protons have a pKa of ~20 (compare this to the pKa value of 40-50 for alkyl C-H).
There are of course many other factors in what makes something acidic but I believe these are the most important before we look at functional groups.
So onto your question.
Here's a list of functional groups (from Compound Interest)
For the purpose of this question I'll just consider a few of the functional groups that have heteroatoms, i.e. atoms that aren't carbon or hydrogen. The first such functional group on that image is an alcohol. It has one oxygen atom which can stabilize the negative charge (once the proton attached to it is removed), but apart from that there aren't any other stabilising features. We would expect alcohols to be somewhat acidic, and they are with a pKa of ~15.
Ethers aren't very acidic. Although protons close to the oxygen atom will be somewhat acidified, try carrying out a "curly arrow" mechanism to see what happens when you try remove a proton from an ether (such as diethyl ether). Protonated ethers, on the other hand, are very acidic (the oxygen atom will be protonated, and have three bonds leading to a formal positive charge. It wants to get rid of this).
Aldehydes will have those α-protons, so these will be somewhat acidic (but NOT the aldehyde proton! Try a curly arrow mechanism to see why), and the same goes for ketones. You can get more and less acidic aldehydes and ketones, and this will vary depending on what exactly the 'R' groups are.
Carboxylic acids have a pKa of ~5. The negative charge is now delocalised over 2 oxygen atoms, and this makes a big difference to the acidity.
Amides are only weakly acidic. This is partially because nitrogen is less electronegative than oxygen, and partially because the nitrogen atom can delocalise its available lone pair of electrons like so:
![]()
(Introduction to Carbonyls)
This would destabilize the conjugate base as instead of "spreading out" the charge it actually puts more negative charge in.
Amines are not acidic at all - you can deprotonate them if you use a very strong base (e.g. an alkyllithium) but this gives you a very unstable conjugate base (it really wants a proton back).
Hope that helps somewhat!
Let’s start with the concept….
Higher electron density within the molecule means that it is…more likely to share electrons with another molecule stronger basicity
Lower electron density within a molecule means that it is…less likely to share electrons with another molecule
weaker basicity
Basically, higher electron density means stronger base and lower electron base means weaker
base!
So what features can influence basicity?
- Resonance
- Electronegativity
- Size of atomic radius
- Inductive effect
- Formal charges
Let’s go through all of these features and learn how they
influence a molecule’s basicity!!!
1. Resonance
Lets recall from the that resonance delocalizes electrons.
If an atom is involved in resonance has a negative formal charge, it delocalizes that
formal charge.
Ex:
Key point: Resonance usually DELOCALIZES electrons
from the atom, thus REDUCING electron density. This
causes the molecule to have
.
First I'll discuss what makes something acidic. I don't know your background in chemistry so feel free to breeze over this part.
The definition of a Brønsted acid is a species that has the tendency to lose a proton. When that proton is removed, you're left with the conjugate base and any features that stabilize that conjugate base result in a stronger acid. For example, CH4 (methane) has CH3- as its conjugate base and its pKa is ~48. HF (hydrofluric acid) has F- as its conjugate base, and a pKa of around 3. This is because fluorine is much more electronegative than carbon (i.e. it likes to pull electron density to itself) and so it can better stabilize the negative charge.
Another feature that can stabilize the conjugate base is the ability to delocalise that negative charge over more than one electronegative atom. This is easier to see on an image, using perchloric acid as an example (pKa about -10):
(from the learning point)
Electron diffraction studies show that the charge is spread out evenly over all oxygen atoms (better demonstrated in the final picture, with the dashed bonds). The bonds are all the same length - if the charge was not delocalised in this way, you would expect single and double bond lengths (which differ by around 20 picometres).
Having electronegative atoms directly attached to, or near, protons also makes them more acidic (as their electron density will be pulled towards that atom). An example of this is in carbonyl compounds:
(Alpha and beta carbon)
The carbon double bonded to the oxygen is the carbonyl compound, and the carbon directly bonded to it is known as the α-carbon. The hydrogen atoms attached to that carbon, then, are known as α-hydrogens (or α-protons). Their close proximity to the electron-withdrawing carbonyl group makes them more acidic, and α-protons have a pKa of ~20 (compare this to the pKa value of 40-50 for alkyl C-H).
There are of course many other factors in what makes something acidic but I believe these are the most important before we look at functional groups.
So onto your question.
Here's a list of functional groups (from Compound Interest)
For the purpose of this question I'll just consider a few of the functional groups that have heteroatoms, i.e. atoms that aren't carbon or hydrogen. The first such functional group on that image is an alcohol. It has one oxygen atom which can stabilize the negative charge (once the proton attached to it is removed), but apart from that there aren't any other stabilising features. We would expect alcohols to be somewhat acidic, and they are with a pKa of ~15.
Ethers aren't very acidic. Although protons close to the oxygen atom will be somewhat acidified, try carrying out a "curly arrow" mechanism to see what happens when you try remove a proton from an ether (such as diethyl ether). Protonated ethers, on the other hand, are very acidic (the oxygen atom will be protonated, and have three bonds leading to a formal positive charge. It wants to get rid of this).
Aldehydes will have those α-protons, so these will be somewhat acidic (but NOT the aldehyde proton! Try a curly arrow mechanism to see why), and the same goes for ketones. You can get more and less acidic aldehydes and ketones, and this will vary depending on what exactly the 'R' groups are.
Carboxylic acids have a pKa of ~5. The negative charge is now delocalised over 2 oxygen atoms, and this makes a big difference to the acidity.
Amides are only weakly acidic. This is partially because nitrogen is less electronegative than oxygen, and partially because the nitrogen atom can delocalise its available lone pair of electrons like so:
(Introduction to Carbonyls)
This would destabilize the conjugate base as instead of "spreading out" the charge it actually puts more negative charge in.
Amines are not acidic at all - you can deprotonate them if you use a very strong base (e.g. an alkyllithium) but this gives you a very unstable conjugate base (it really wants a proton back).
Hope that helps somewhat!
Let’s start with the concept….
Higher electron density within the molecule means that it is…more likely to share electrons with another molecule stronger basicity
Lower electron density within a molecule means that it is…less likely to share electrons with another molecule
weaker basicity
Basically, higher electron density means stronger base and lower electron base means weaker
base!
So what features can influence basicity?
- Resonance
- Electronegativity
- Size of atomic radius
- Inductive effect
- Formal charges
Let’s go through all of these features and learn how they
influence a molecule’s basicity!!!
1. Resonance
Lets recall from the that resonance delocalizes electrons.
If an atom is involved in resonance has a negative formal charge, it delocalizes that
formal charge.
Ex:
Key point: Resonance usually DELOCALIZES electrons
from the atom, thus REDUCING electron density. This
causes the molecule to have
