Electronegativity of halogens

  • Thread starter triplej
  • Start date
  • #1
7
0

Main Question or Discussion Point

Can anyone explain concisely why Trifluoroacetic Acid is more acidic than Trichloroacetic Acid, but HCl is more acidic than HF?
 

Answers and Replies

  • #2
183
5
Well, let's start with HF and HCl. We know that HCl is stronger because Cl is more electronegative, making the H+ ion more likely to leave.

In the case of halogenated acids, you have to look at the partial charges. The halogens, because they have a strong partial negative charge, put a partial positive charge on the carbon they're attached to. That, in turn, puts a partial negative charge charge on the carbon in the carboxylic acid functional group. That partial negative makes the oxygen in the OH group slightly positive.

Since Cl has a stronger partial negative than F, the oxygen on the trichloroacetic acid will have a stronger positive charge than the oxygen in the trifluoroacetic acid, making it a weaker base, because of the same reasoning used with HCl and HF.

Does that make sense? It's all about keeping track of the partial charges and knowing how electronegativity of the atom attached to the acidic H affects the strength of the acid.
 
  • #3
397
21
Well, let's start with HF and HCl. We know that HCl is stronger because Cl is more electronegative, making the H+ ion more likely to leave.

In the case of halogenated acids, you have to look at the partial charges. The halogens, because they have a strong partial negative charge, put a partial positive charge on the carbon they're attached to. That, in turn, puts a partial negative charge charge on the carbon in the carboxylic acid functional group. That partial negative makes the oxygen in the OH group slightly positive.

Since Cl has a stronger partial negative than F, the oxygen on the trichloroacetic acid will have a stronger positive charge than the oxygen in the trifluoroacetic acid, making it a weaker base, because of the same reasoning used with HCl and HF.

Does that make sense? It's all about keeping track of the partial charges and knowing how electronegativity of the atom attached to the acidic H affects the strength of the acid.
Flourine is the most electronegative atom on the periodic table. If it was all about electronegativity than HF would be the strongest haloacid and HI the weakest (which is not the case at all).

To the OP:
In general when you want to decide on why acids/bases are stronger/weaker you should deprotonate/protonate the thing and look for any type of stablizing/destabilizing effects, with the two biggest effects to watch out for being the inductive effect and resonance stabilization. Bear in mind that the only thing these two effects are doing is taking a charge which is localized in space and spreading it out thus decreasing 'charge density.'

So lets take HF and HCl and deprotonate them. You get F- and Cl-. Obviously since HCl is a strong acid the Cl is "happier" hanging out with a negative charge. Why? Well one explanation that I know of, and which fits with most observations, is that Cl is a larger atom than F thus when there is a negative charge localized on the atom itself, the charge density is smaller (meaning the charge is more spread in space-same concept of induction and resonance stabilization). So yes F is more electronegative but once deprotonated there is the same amount of charge (-1) localized around a smaller atom. This explains acid strength of halo-acids (HX).

HF < HCl < HBr < HI (pKa: 3.2 > -7 > -8 > -99 respectively)

Now when we are talking halogen substituted COOH's its a different story and all about electronegativity. Simply put F is more electronegative than Cl and will exert a greater "pull" on the carboxylate anion thus stabilizing the anion through induction. Notice though that the negative charge will not be on the halogen and its all about the sigma withdrawing properties of halogens/electronegative atoms, in general.
 
  • #5
7
0
Thank you both, I get it now. Should have remembered the stability vs. electronegativity thing.
 
  • #6
Yes, just look at the Bond Energy of H-F versus H-Cl or H-Br. HF has a pretty strong bond, look it up in a table, its tightly packed so the electronegativity is high (also because its nearly a noble gas), but F makes strong bonds in general. Because the bond strength is high, the acidity is low.. The HF bond is actually lower than it should be because fluorine is abnormal in the trends in its group. However, it still has higher bond strength than HCl and HBr.

Bond Energies: HF = 135 kcal/mole, HCl = 103 kcal/mole, HBr ~ 80, HI ~ 60 kcal/mole

If you pay attention to detail and get real values for these, and put in a bit of work on your own, you can actually calculate the pH directly from these numbers using DelG=-RTlnK, the definition of equilibrium constant and Definition of pH. However, the equations are only good for strong acids, requiring some simple approximations that do not hold for HF, they are off by about 0.3 pH units or so. Anyways, its just thermodynamics is easiest way to calculate.
Thermodynamics determines almost everything in chemistry, in general.

Keep in mind, that if you compare HF as an acid to say CH4, that you NEED to include the heat of formation of the reactants, which varies. Acids do not have the same heat of formation of the reactants as the alkanes.

If you want to understand all these trends in general, you need to take a look at the full tables of heats of formation listed in say Atkins Physical Chemistry, the appendix, and they will tell you the DelG for any reaction. Memorize the trends and you can tell what is an acid and what is not.
 
  • #7
183
5
A bit late on that, eh?
 
  • #8
81
0
Flourine is the most electronegative atom on the periodic table. If it was all about electronegativity than HF would be the strongest haloacid and HI the weakest (which is not the case at all).
Acid strength = enthalpy of dehydration
+ enthalpy of dissociation
+ ionization energy of H+
+ electron affinity of X-
+ enthalpy of hydration of H+ and X-

ΔG = -RTlnκ = ΔH - TΔS

The reason why HF has low total ΔH value is:
1) Enthalpies of dissociation show that H-F(1.7 A°)bond is much stronger that of H-I(1.0 A°). Hence its enthalpy is almost twice greater.
2) The enthalpy of dehydration is much higher because of strong hydrogen bonds in HF.
3) The very low value of electron affinity of F- also contributes. And though the its hydration energy of F- is very high, it is not enough to offset the above terms.

I hope this helps.
 
  • #9
81
0
why Trifluoroacetic Acid is more acidic than Trichloroacetic Acid
The concept that should be applied in this case is of inductive effect (+I and -I)

The acidic strength of a compound can be characterized by its ability or ease to donate H+ ions or proton.
In case of carboxylic group R-COOH; when it donates a proton, what left is R-COO-. Now the the group which will stabilize the negative charge on R-COO- will be favorable and the proton can be released easily if such a group is attached. This means that the acidic strength will increase with the ease to donate H+. Such a group must be -I group which will "put" a partial positive charge on O- which will stabilize its negative charge. And the fact is F has a greater inductive power than Cl and hence Trifluoroacetic Acid will be more acidic than Trichloroacetic Acid.

I hope this helps too.
 

Related Threads on Electronegativity of halogens

  • Last Post
Replies
5
Views
11K
  • Last Post
Replies
10
Views
2K
  • Last Post
Replies
9
Views
26K
  • Last Post
Replies
2
Views
11K
  • Last Post
Replies
2
Views
2K
Replies
1
Views
5K
  • Last Post
Replies
3
Views
2K
  • Last Post
Replies
6
Views
3K
Replies
19
Views
25K
Replies
2
Views
2K
Top