Non-electrolytes do not affect the pH of a solution?

[Alfred001]

Could you tell me whether my understanding here is correct:

In order for something to affect the pH of a solution, it needs to dissociate. Non-electrolytes do not dissociate in water and would therefore not affect its pH. Would they, however, dissociate in gastric juice?

[exchemist]

33 minutes ago, Alfred001 said:

Could you tell me whether my understanding here is correct:

In order for something to affect the pH of a solution, it needs to dissociate. Non-electrolytes do not dissociate in water and would therefore not affect its pH. Would they, however, dissociate in gastric juice?

Think of ammonia. It does not dissociate appreciably in solution but accepts H+, causing water to dissociate, and thereby raises pH.

Stomach acid has a pH ~2, so has plenty of free H+. I can't see why this would in itself cause other substances to dissociate, though protein dissociation is catalysed by enzymes in the acid environment of the stomach, I believe.  

[Alfred001]

23 minutes ago, exchemist said:

Think of ammonia. It does not dissociate appreciably in solution but accepts H+, causing water to dissociate, and thereby raises pH.

Stomach acid has a pH ~2, so has plenty of free H+. I can't see why this would in itself cause other substances to dissociate, though protein dissociation is catalysed by enzymes in the acid environment of the stomach, I believe.  

So a substance does not necessarily need to dissociate to affect the pH of a solution? But can lowering of pH happen only when dissociation happens (because the substance needs to donate a H)?

[exchemist]

44 minutes ago, Alfred001 said:

So a substance does not necessarily need to dissociate to affect the pH of a solution? But can lowering of pH happen only when dissociation happens (because the substance needs to donate a H)?

Not necessarily. For example, some cations form complexes with OH-, thereby causing water to release H+, acidifiying the solution.

Both in the case of ammonia and in the case of metal cations such as Al 3+, water is caused to split, by abstracting either H+ in the case of ammonia or OH- in the case of  Al 3+, leaving behind extra OH- or H+ respectively and thus altering the pH of the solution.

[Alfred001]

I see. Ok, let me ask a more specific question, and this is a really multifactorial one, so I know this might be tough to answer:

I know that after a meal gastric pH rises due to buffering effect of food, now suppose you were to add dextrose on top of that, would you expect the addition of dextrose to affect the pH in either direction, bearing in mind the dextrose might be broken down by enzymes into, I think, maltose and fructose? And let's say no further gastric acid will be secreted, to simplify things.

[MigL]

Always wondered about that @exchemist

NH3 in water solution becomes NH₄⁺ and OH^-, making the solution strongly basic.
But pure Ammonia, by itself, is not basic, just a strong reducer ?

I never did understand a Lewis acid as an electron pair acceptor.
( Grade 13 was a long time ago )

 

[John Cuthber]

5 minutes ago, Alfred001 said:

, would you expect the addition of dextrose to affect the pH in either direction

No, or at least, only very slightly.
More sugar would draw water in from the surrounding tissues (and from the blood) and that would dilute the acid a bit, raising the pH.

[exchemist]

1 hour ago, MigL said:

Always wondered about that @exchemist

NH3 in water solution becomes NH₄⁺ and OH^-, making the solution strongly basic.
But pure Ammonia, by itself, is not basic, just a strong reducer ?

I never did understand a Lewis acid as an electron pair acceptor.
( Grade 13 was a long time ago )

 

Actually not that much becomes converted to ammonium and hydroxide. Most of it is still NH3(aq) : https://en.wikipedia.org/wiki/Ammonia_solution   

As I recall, ammonia is not much of a reducing agent as it can't give up electrons easily, only share them in a covalent bond, as in H3N:->H+ . (If it actually reduced H+, hydrogen would be evolved.)

But yeah it's easy to get in a muddle with Lewis acids and bases. It's one of those concepts I tended to dodge, not finding it awfully helpful on most occasions.  

1 hour ago, Alfred001 said:

I see. Ok, let me ask a more specific question, and this is a really multifactorial one, so I know this might be tough to answer:

I know that after a meal gastric pH rises due to buffering effect of food, now suppose you were to add dextrose on top of that, would you expect the addition of dextrose to affect the pH in either direction, bearing in mind the dextrose might be broken down by enzymes into, I think, maltose and fructose? And let's say no further gastric acid will be secreted, to simplify things.

No I don't think so. Sugars have hydroxy groups on them that can be protonated in acid conditions, i.e.  R-OH + H⁺ <-> R-OH₂⁺ , but they are no more liable to be protonated than a water molecule, i.e. H₂O + H⁺ <->H₃O⁺, so that won't raise the pH. 

(While sweet substances make acidic substances taste less acid, they do not do so by neutralising the acidity. It's just a trick of the taste buds.) 

[chenbeier]

On 3/5/2024 at 10:23 PM, MigL said:

But pure Ammonia, by itself, is not basic, just a strong reducer ?

Ammonia is a gas at standard conditions, we cannot speak it has an pH, like water steam has no one.

But liquid ammonia has also its autopotolysis like water it has.

2 NH3 => NH4+ + NH2-

2H2O => H3O+ + OH-

So could have a pH scale based on ammonium theoretical.

 

Edited March 7 by chenbeier

[joigus]

Is this (admittedly rough) understanding that I've acquired through the years correct?:

The currency of red-ox reactions is electrons

The currency of acid-base reactions is protons

Now, in a manner of speaking,

Both oxydisers and reductors can be understood in terms of "soaking up" and "giving off" electrons

Both bases and acids can be understood in terms of "soaking up" and "giving off" protons

That's the reason why so much of chemistry hinges around these two dual concepts

Other cations, even the smallest ones, like Li⁺, are "monsters" in comparison to H⁺. Orders of magnitude so much so. So even though the mean free path of a proton is sizeably higher than that of an electron, it's bound to be gigantic as compared to that of even such a small thing as Li⁺. That would qualitatively account for an extraordinarily high mobility of protons, thereby the reactiveness of anything that either gives them off or soaks them up. That's the key to the concept of Lewis acids. Is it not?

Then, for something to be a base, in its most general sense, it must be able to soak up protons. But for it to display this character, there must be some protons around to soak up. Wouldn't something like this be at the root of NH₃ not "behaving as a base" just by itself, or in the presence of chemicals that cannot give off protons?

Wouldn't it behave as a base in the absence of water, but in the presence of acids (neutralisation) like,

NH₃+A --->NH₄⁺+A^-

with A being any acid?

[John Cuthber]

With a lot of trouble and care (and a mass spectrometer) you can measure acid base equilibria in the gas phase with no solvent present.
Many people have got PhDs by doing this, but I'm not sure it's had much wider use. (By wider, I mean outside the lab)

You don't get "H⁺" ions in water. You get a hydrated version, typically modelled as [H₉O₄]⁺

So you are almost always looking at how strong an acid is compared to water.
In principle, if you have just liquid ammonia, there are always some protons around to soak up.
They arise from ammonia acting as a base.
2 NH3 --> [NH₂] ^-  + [NH₄]⁺
The extent of that reaction is tiny, but not zero. (It's about 1 in 10 ^33)

[KJW]

3 hours ago, joigus said:

The currency of red-ox reactions is electrons

The currency of acid-base reactions is protons

I prefer to think of it this way:

Oxidation is the loss of electrons.

Reduction is the gain of electrons.

Acids accept lone-pairs of electrons.

Bases donate lone-pairs of electrons.

It would seem that there is a correlation between being an oxidising agent and being an acid, and a correlation between being a reducing agent and being a base. However, they are distinct notions. The hydride anion, H^–, is an example of both a base and a reducing agent. Indeed, the reaction:

H^– + H⁺ —> H₂

is both an acid-base reaction and a redox reaction. By contrast, diborane, B₂H₆, is a Lewis acid that is a reducing agent.

 

 

46 minutes ago, John Cuthber said:

You don't get "H⁺" ions in water.

Whereas the hydroxide ion, OH^–, is not really much different from any other base, the hydrogen ion H⁺, is singularly unique as a Lewis acid in having no electrons at all. As a proton, a subatomic particle, its charge density is so great that, within the context of chemistry, it will never be free. It will always be bound to some electron density, regardless of how reluctant that electron density is to bond to anything. Therefore, a Brønsted acid is conceptually distinct from a Lewis acid in that the hydrogen ion is always bound to a base, and the measure of the Brønsted acidity is really a measure of the basicity of that base (the weaker the base, the stronger the Brønsted acid).

 

Edited March 7 by KJW

[KJW]

On 3/6/2024 at 7:23 AM, MigL said:

But pure Ammonia, by itself, is not basic, just a strong reducer ?

One thing I find intriguing is that by an oxidative process, ammonia can be converted to the strongest known reducing agent (azide ion, N₃^–).

 

 

On 3/6/2024 at 7:23 AM, MigL said:

I never did understand a Lewis acid as an electron pair acceptor.

Consider boron trifluoride, BF₃: A boron atom has three electrons in its outer shell. The three fluorine atoms each provide one electron, giving the boron atom of boron trifluoride six electrons in total in its outer shell. But boron atoms want to have eight electrons in their outer shell to complete the octet. Thus, boron trifluoride will react with a fluoride ion, F^–, which supplies a lone-pair of electrons, to produce the tetrafluoroborate anion, BF₄^–, which has eight electrons in the outer shell of the boron atom, completing the octet.