[Needaask]

The acidity increases across a period and down a group. I understand that it increases down the group because the covalent bond between the H and the other element decreases in strength due to being bigger causing the charge density to becomes smaller. So being harder to break that bond, its harder to remove the proton so it makes it less acidic.

However, why would the acidity increase across a period? Across a period the bond strength increases so shouldn't acidity decrease rather than increase? If the bond strength increases it'll be harder to cut off that proton so shouldn't it become less acidic across the period too?

Why is this so? Thanks so much for the help

[9-92-6-19]

Well, the acidity increases because, generally, electronegativity increases across a period. This results in a stronger acid, because an increase in electronegativity creates a more ionic compound therefore making the complete disassociation of HX into H⁺ and X^- more feasible; however, this effect is counteracted as one goes down a group where the total bond strength becomes much weaker.

[Needaask]

Well, the acidity increases because, generally, electronegativity increases across a period. This results in a stronger acid, because an increase in electronegativity creates a more ionic compound therefore making the complete disassociation of HX into H⁺ and X^- more feasible; however, this effect is counteracted as one goes down a group where the total bond strength becomes much weaker.

hmm, could you explain that factor? The ionic character increases and the covalent character decreases so its something like H-O has 90% covalent and 10% ionic while H-F has 70% covalent and 30% ionic. So why should H-F have a stronger bond? I mean if we explain that the ionic character increases, there's also the covalent character decreasing.

Also, for this question, shouldn't the reason whether the H-X is more acidic or not depends only on strength of the bond? Because if an ionic bond is very strong, i don't think the solvent-ion interaction is strong enough to dissolve it. So in that sense if the bond is very strong, it would also be very hard to seperate the H from the X to create new bonds?

Thanks so much for the reply

[9-92-6-19]

If we were to look at an electronegativity chart we would see that H is 2.1, O is 3.5, and F is 4.0

We also know that:

A compound is ionic if and only if ΔEN > 2.0
A compound is polar covalent if and only if .5 < ΔEN < 1.7
A compound is non-polar covalent if and only if ΔEN < .5

Yet there is a small gap between 1.7 and 2.0. The rule for this gap is very important to acidity; if a metal is involved in the compound, then the compound is said to be ionic—if not then polar covalent.

HO bonds have a ΔEN = 1.4 so they follow rule number two, but HF is 1.9; so strong, in fact, that it is on the brink of being an ionic substance. Therefore, HF is a much stronger acid compared to H₂O because of its much greater capability to attract the electron from the hydrogen.

Yet, because F^- can also act as a Lewis base shown by the following formula:

F^- + H₂O  HF + OH^-

The acid can, in effect, partially cancel itself out slightly.

Furthermore, HF is capable of strong intermolecular forces: hydrogen bonding; this characteristic inhibits complete theoretical disassociation.

If we were to move down a group to Chlorine then we can see that Cl EN = 3.0, therefore HCl  ΔEN = .9, which means it is a polar covalent compound; thus, because the bond between H⁺ and Cl^- is much weaker (but still strong compared to a non-polar covalent acid) than that in H⁺ and F^- then the aforementioned oddities with HF no longer occur. For example, Cl^- will not act as a Lewis base and attract a H⁺ from water.

Cl^- + H₂O  No reaction

Similarly, HCl does not qualify for hydrogen bonding, therefore making it favorable for the molecule to completely disassociate.

[Needaask]

If we were to look at an electronegativity chart we would see that H is 2.1, O is 3.5, and F is 4.0

We also know that:

A compound is ionic if and only if ΔEN > 2.0
A compound is polar covalent if and only if .5 < ΔEN < 1.7
A compound is non-polar covalent if and only if ΔEN < .5

Yet there is a small gap between 1.7 and 2.0. The rule for this gap is very important to acidity; if a metal is involved in the compound, then the compound is said to be ionic—if not then polar covalent.

HO bonds have a ΔEN = 1.4 so they follow rule number two, but HF is 1.9; so strong, in fact, that it is on the brink of being an ionic substance. Therefore, HF is a much stronger acid compared to H₂O because of its much greater capability to attract the electron from the hydrogen.

Yet, because F^- can also act as a Lewis base shown by the following formula:

F^- + H₂O  HF + OH^-

The acid can, in effect, partially cancel itself out slightly.

Furthermore, HF is capable of strong intermolecular forces: hydrogen bonding; this characteristic inhibits complete theoretical disassociation.

If we were to move down a group to Chlorine then we can see that Cl EN = 3.0, therefore HCl  ΔEN = .9, which means it is a polar covalent compound; thus, because the bond between H⁺ and Cl^- is much weaker (but still strong compared to a non-polar covalent acid) than that in H⁺ and F^- then the aforementioned oddities with HF no longer occur. For example, Cl^- will not act as a Lewis base and attract a H⁺ from water.

Cl^- + H₂O  No reaction

Similarly, HCl does not qualify for hydrogen bonding, therefore making it favorable for the molecule to completely disassociate.

Hi thanks for the reply

Actually why would the bond be stronger if the electronegativity fifference is large in magnitude? Because if the ionic character increases the covalent character would decrease. So saying that since the ionic character is greater, it would suggest that ionic bonds are stronger than covalent bonds. however i heard that in some cases covalent bonds could be stronger and vice versa.

Going into the question though what do you mean by "HF is a much stronger acid compared to H2O because of its much greater capability to attract the electron from the hydrogen". I'm not very familiar with Lewis acids and bases because my course/book that I'm reading covers only Bronsted-Lowry's acid and bases. Also, I don't quite understand how the presence of hydrogen bonding would affect the solubility of the H-X species.

For the two equations I think i understand why F- would react with water to form an OH- and HF. Is the equilibrium tilted to the left because floruine is much more electronegative than oxygen. So as a result F- is much more stable than the O- such that the equilibrium is shifted to the left?

Then for the second equation if Cl^- +H₂O  HF+OH^- I would guess that this equation would be pushed to the left even more than the first because Cl- is larger than F- so it the charge density is smaller making it more stable. Hence the equilibrium would be so much to the left that the Cl- ion can't undergo hydrolysis with water? So coupled with this fact and that HCl bond is weaker, HCl would form ions more easily than HF. However, i have an opposing factor to this which is F being more electronegative than Cl but that factor is easily outweighed by the charge density.


I'm thinking if i compare HF vs NH3, NH3 having weaker bonds should be easier to form NH2- and H3O+ as compared to water to form H3O+ and OH-. But in my book they explained that the conjugate bases OH- is more stable than NH2- due to O having a higher electronegativity than nitrogen. So the only think i can think of now is that the reaction NH2- +H2O->NH3+OH- proceeds more than the F-+H2O->HF+OH-. So for the first reaction, OH- being more stable the reaction would tend to the right so all the NH3 will reform and OH- will cancel the effect of the H3O+ (maybe not all but to bring my point) so we'll get back to NH3 and H2O with a little NH2- and H3O+ only. While in the second reaction since less of the reaction proceeds I would end up with more H3O+ and F- making it more acidic than the NH3.

So its like the 2 factors 'challenge' each other and unlike going down the group whereby they supported each other (weaker bond+more stable anion by charge density+but more unstable by electronegativity) here they oppose each other (weaker bond+less stable anion and no charge density factor as all the anions are about the same size) so in this case, the electronegativity factor outweighs the bond strength factor making it the consideration we take when determining which species is more acidic?