Buffers...

[canadianofpeace]

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I know the typical rule is that a weak acid has a strong conjugate base and vice versa. A buffer can be made from HF and its conjugate base, for example using mixing it with NaF.

However, I am confused because F is not a strong base. HF is a weak acid so shouldn't its conjugate base be strong?. Is F an outlier to the rule?

 

[pharmolu]

weak acids do NOT have a strong conjugate base. weak acids have a weak conjugate base.

 

[Teleologist]

weak acids do NOT have a strong conjugate base. weak acids have a weak conjugate base.

Who told you that? That's wrong. Consider water, a weak acid. What's its conjugate base? Hydroxide ion. That's strong!

 

[Teleologist]

I know the typical rule is that a weak acid has a strong conjugate base and vice versa. A buffer can be made from HF and its conjugate base, for example using mixing it with NaF.

However, I am confused because F is not a strong base. HF is a weak acid so shouldn't its conjugate base be strong?. Is F an outlier to the rule?

Brah, it's a sliding scale. It's not categorical. The weaker the acid, the stronger its conjugate base. At some point you cross the threshold for "strong" in water solution. That's when you get to water or any acid weaker than water.

 

[pharmolu]

You have another example, other than the particular solvent of the system? In all aqueous solutions, by virtue of the leveling effect and autoionization of water, the conjugate of a weak acid is a weak base.

 

[Teleologist]

You have another example, other than the particular solvent of the system? In all aqueous solutions, by virtue of the leveling effect and autoionization of water, the conjugate of a weak acid is a weak base.

Brah, I don't think you can explain how the auto-ionization of water is relevant to this discussion.

Because it isn't.

 

[pharmolu]

Brah, I don't think you can explain how the auto-ionization of water is relevant to this discussion.

Because it isn't.

LMAO. Ka * Kb= 10 ^-14 of conjugate pairs, which is why if one is weak, the other cannot completely dissociate in water. BTW, 10^-14 is the self-ionization constant, since you don't seem to see how it meshes in this discussion.

 

[Teleologist]

LMAO. Ka * Kb= 10 ^-14 of conjugate pairs, which is why if one is weak, the other cannot completely dissociate in water. BTW, 10^-14 is the self-ionization constant, since you don't seem to see how it meshes in this discussion.

Brah, I don't think you understand causation.

For example, a patient is sick not because Google said so.

A patient is sick because of a virus (or whatever).

Kw =Ka*Kb is a side effect, not a cause, or a "why."

 

[pharmolu]

You're sounding crazy right now 'brah'. What does causality have anything to with the statement all weak acids have weak conjugate bases? You're wrong anyway, the autoionization is in fact the cause, because it exhibits the relationship for conjugate pairs, which is pivotal to the explanation; if pKa +Pkb didn't equal 14, then there would be no argument. Anyway, I'm done with this discussion; I hope that, if you're set to take the MCAT, you don't say that a weak acid has a strong conjugate base.

 

[Teleologist]

You're sounding crazy right now 'brah'. What does causality have anything to with the statement all weak acids have weak conjugate bases? You're wrong anyway, the autoionization is in fact the cause, because it exhibits the relationship for conjugate pairs, which is pivotal to the explanation; if pKa +Pkb didn't equal 14, then there would be no argument. Anyway, I'm done with this discussion; I hope that, if you're set to take the MCAT, you don't say that a weak acid has a strong conjugate base.

Brah, water's conjugate base. You already admitted you were wrong.

 

[Vybe]

Addressing the initial question, there are actually 3 classes of conjugate pairs that exist.

Very strong acids have very weak conjugate bases
Weak acids have weak conjugate bases &
Very weak acids have very strong conjugate bases.

From what I've learned, HF is a very weak acid, which results in F- being a strong base. You can look at this in terms of electron affinity.

When in solution, Fluoride (F-) has its full octet of electrons, surrounded by a shell of hydrogen atoms. This makes the F- extremely stable, which has no interest in binding with anything else (speaking in extreme terms). The more independent an atom can be, the more basic it is; referring to the concept of basicity.

We know from many resources that HF is a very weak acid because, like we mentioned of Fluorine, it is extremely electronegative; most in the periodic table. So it is hard for fluorine to get rid of its H, even when put in an (aq) solution of water. That is what makes HF very weak acid. It doesn't want to donate its H+ as opposed to Cl -, which gives away the H+ happily. Hope this perspective helps you a little better

 

[Teleologist]

From what I've learned, HF is a very weak acid, which results in F- being a strong base. You can look at this in terms of electron affinity.

Brah, you did not just say that.

First, if by "very weak acid" you mean weaker than water, then you're wrong.

Second, F- is not a strong base.

Third, elemental fluorine has a relatively low electron affinity due to its great charge density. Fluorine actually breaks the trend of electron affinity.

Fourth, Coulomb's law, are you really gonna apply the concept of electron affinity to a charge dense anion such as the fluorine anion? Last time I checked, opposites attract, and ... non-opposites ... don't attract ... unless you're Macklemore. In other words, the fluorine anion is not going to be attracted to electrons because 1) it's already isoelectronic with a noble gas and 2) it has a high negative charge density.

Fifth, why are we even talking about electrons here? Last time I checked, we were talking about the Bronsted-Lowry definition of acids. If anything we should be talking about proton affinity.

We know from many resources that HF is a very weak acid because, like we mentioned of Fluorine, it is extremely electronegative; most in the periodic table. So it is hard for fluorine to get rid of its H, even when put in an (aq) solution of water.

Sixth, you got that reversed. By nature of its electronegative nature, the fluorine in HF gives the H a partial positive charge, and this makes the hydrogen atom in HF electrophilic and thus reactive.

That is what makes HF very weak acid. It doesn't want to donate its H+ as opposed to Cl -, which gives away the H+ happily. Hope this perspective helps you a little better

Seventh, Cl- doesn't have a H+ to give away.

Eigth, I hope you get some perspective from this.

 

[Vybe]

Brah, you did not just say that.

First, if by "very weak acid" you mean weaker than water, then you're wrong.

Second, F- is not a strong base.

Third, elemental fluorine has a relatively low electron affinity due to its great charge density. Fluorine actually breaks the trend of electron affinity.

Fourth, Coulomb's law, are you really gonna apply the concept of electron affinity to a charge dense anion such as the fluorine anion? Last time I checked, opposites attract, and ... non-opposites ... don't attract ... unless you're Macklemore. In other words, the fluorine anion is not going to be attracted to electrons because 1) it's already isoelectronic with a noble gas and 2) it has a high negative charge density.

Fifth, why are we even talking about electrons here? Last time I checked, we were talking about the Bronsted-Lowry definition of acids. If anything we should be talking about proton affinity.



Sixth, you got that reversed. By nature of its electronegative nature, the fluorine in HF gives the H a partial positive charge, and this makes the hydrogen atom in HF electrophilic and thus reactive.



Seventh, Cl- doesn't have a H+ to give away.

Eigth, I hope you get some perspective from this.

I apologize, I may have used electron affinity in the wrong way.

Your 7th point, I should clarify, I was referring to the easiness in which HCl differentiates into Cl- and H+ in water, as opposed to how easily HF differentiates. F loves to hold onto its electrons, which is what I meant by electron affinity. Yes, Fluorine is electron dense, aka electronegative.

And thanks Teleologist, I had my concepts in reverse. HF is a weak acid, F- is simply a weak conjugate base. Which is the 2nd of the 3 conjugate acid-base pairs I had described.

 

[gettheleadout]

Who told you that? That's wrong. Consider water, a weak acid. What's its conjugate base? Hydroxide ion. That's strong!

To be clear to other readers, it's also incorrect so say that all weak acids have strong conjugate bases.

Brah, it's a sliding scale. It's not categorical. The weaker the acid, the stronger its conjugate base. At some point you cross the threshold for "strong" in water solution. That's when you get to water or any acid weaker than water.

Yep.

You're sounding crazy right now 'brah'. What does causality have anything to with the statement all weak acids have weak conjugate bases? You're wrong anyway, the autoionization is in fact the cause, because it exhibits the relationship for conjugate pairs, which is pivotal to the explanation; if pKa +Pkb didn't equal 14, then there would be no argument. Anyway, I'm done with this discussion; I hope that, if you're set to take the MCAT, you don't say that a weak acid has a strong conjugate base.

The dissociation constant for water is actually temperature dependent, so the value of Kw = 10^–14 or pKw = 14 is not inherently meaningful.

weak acids do NOT have a strong conjugate base. weak acids have a weak conjugate base.

In any case the problem with statements like these is the definition of terms like "weak acid." If we mean a species that does exhibit acidic reactivity but does so to a very low degree (e.g. Ka = 10^–5) then it will be correct to state that the conjugate of such a species will also exhibit a low degree of reactivity (in this case basic). This is the case for the great majority of water soluble acid/base species (e.g. acetic acid, hydrofluoric acid, etc.). If we instead mean, by "weak acids," species that are conjugates of basic species but that themselves exhibit virtually zero acidic activity (e.g. solvated sodium ion), then the statement that these species will have basic conjugate species classified as "strong bases" (where "strong" refers to ~100% hydrolysis in solution) will be correct. This is a problem of equivocation. TBR specifically delineates by terming the former, measurably reactive acidic species "weak acids" and the latter, nonreactive acidic species "very weak acids."

Edit: Accidentally put chloride in where I meant sodium above (now fixed). Solvated chloride ion is basic but not really; that's the point I was trying to make. It's the conjugate of an acid, so we call it a base, but it doesn't actually undergo hydrolysis to form HCl again. So calling it a "weak base" isn't clear enough because it's not comparable in terms of basic activity to something like acetate. Thus, we call it a "very weak base."

 

[Teleologist]

We know from many resources that HF is a very weak acid because, like we mentioned of Fluorine, it is extremely electronegative; most in the periodic table. So it is hard for fluorine to get rid of its H, even when put in an (aq) solution of water. That is what makes HF very weak acid. It doesn't want to donate its H+ as opposed to Cl -, which gives away the H+ happily. Hope this perspective helps you a little better

To be clear to other readers, it's also incorrect so say that all weak acids have strong conjugate bases.

Yep.

The dissociation constant for water is actually temperature dependent, so the value of Kw = 10^–14 or pKw = 14 is not inherently meaningful.

In any case the problem with statements like these is the definition of terms like "weak acid." If we mean a species that does exhibit acidic reactivity but does so to a very low degree (e.g. Ka = 10^–5) then it will be correct to state that the conjugate of such a species will also exhibit a low degree of reactivity (in this case basic). This is the case for the great majority of water soluble acid/base species (e.g. acetic acid, hydrofluoric acid, etc.). If we instead mean, by "weak acids," species that are conjugates of basic species but that themselves exhibit virtually zero acidic activity (e.g. solvated chloride ion), then the statement that these species will have basic conjugate species classified as "strong bases" (where "strong" refers to ~100% hydrolysis in solution) will be correct. This is a problem of equivocation. TBR specifically delineates by terming the former, measurably reactive acidic species "weak acids" and the latter, nonreactive acidic species "very weak acids."

Brah, you probably got a 45 on the MCAT. You are God.

 

[gettheleadout]

Brah, you probably got a 45 on the MCAT. You are God.

Not quite and not quite haha. 😉

Plus I typoed that post lol, just fixed it.

 

[pharmolu]

The dissociation constant for water is actually temperature dependent, so the value of Kw = 10^–14 or pKw = 14 is not inherently meaningful.

Of course that's true, I misspoke when I said 14 but meant that in an aqueous solution, the two add up to the auto ionization constant of water( which is 14 at ~25 degrees).

In any case the problem with statements like these is the definition of terms like "weak acid." If we mean a species that does exhibit acidic reactivity but does so to a very low degree (e.g. Ka = 10^–5) then it will be correct to state that the conjugate of such a species will also exhibit a low degree of reactivity (in this case basic). This is the case for the great majority of water soluble acid/base species (e.g. acetic acid, hydrofluoric acid, etc.). If we instead mean, by "weak acids," species that are conjugates of basic species but that themselves exhibit virtually zero acidic activity (e.g. solvated sodium ion), then the statement that these species will have basic conjugate species classified as "strong bases" (where "strong" refers to ~100% hydrolysis in solution) will be correct. This is a problem of equivocation. TBR specifically delineates by terming the former, measurably reactive acidic species "weak acids" and the latter, nonreactive acidic species "very weak acids."

I guess I was assuming the absolute definition in each case; by strong I meant completely dissociation in water, by 'weak' I meant not strong, but stronger than the solvent(water, in most cases)and negotiable would be virtually no dissociation in water (weaker than water). However, I have never seen any book use a definition that deviates from this.