Acids and Bases part 3

Yes, we're back to good ol' acids and bases. A lot of it continues from Year 11, but there's some new bits to learn too.

Part 1: http://year11misadventures.blogspot.com.au/2012/10/acids-and-bases-part-1.html
Part 2: http://year11misadventures.blogspot.com.au/2012/10/acids-and-bases-part-2.html

Reactions, Equations and Stoichiometry

The main points that I need to go through here are writing equations and stuff for neutralisation and the hydrolysis of salts of weak acids and bases, and calculating pH.

Neutralisation equations can be found in part 1, and pH can be found in part 2.

It says here that you also need to be able to calculate the pH of the resulting solution when strong acid-base solutions are mixed. I think in this case it means those limiting reagent questions where you have a certain quantity of an acid mixing with a certain quantity of a base, and you have to work out whether the acid or the base is in excess. Then you have to work out how many H+ or OH- ions you have in excess and then use this to work out the hydrogen ion concentration of the solution, and finally the pH.

Whew. That now leaves us with writing equations for the hydrolysis of salts of weak acids and bases!

In part 1 I mentioned conjugate acids and bases. What I didn't mention, though, is that the strength of the conjugate acid/base depends on the strength of the acid or base that formed it. Strong acids tend to produce weaker bases while weaker acids tend to produce stronger bases. For example, HCl is a strong acid, so its conjugate base, Cl-, is a weak base. Meanwhile, CH₃COOH is a weak acid, while CH₃COO^- is a relatively strong base.

Note that I said a relatively strong base. While CH₃COO^- is a stronger base than Cl-, it comes nowhere close to NaOH.

The opposite holds true for bases too- strong bases tend to have weaker conjugate acids while weak bases tend to have stronger conjugate acids.

Note that really strong bases tend to also produce bases as conjugate "acids," and really strong acids produce acids as conjugate "bases." For example, O^2- is such a strong base that OH^- is also a base, and H₂SO₄ is such a strong acid that HSO₄^- is also an acid.

Anyway, my point being was that, for a salt to hydrolyse with water, the salt must contain an ion that is a strong enough acid or base. Cl- won't hydrolyse as, being the conjugate base of a strong acid, it simply isn't a strong enough base. CH₃COO^-, will, however, hydrolyse. The whole conjugate acid/base thing is just one way that might help you decide whether something is strong enough to hydrolyse with water or not.

So, what exactly is hydrolysis? Well, it's basically an acid-base reaction with water. For example, our good old ethanoate/acetate ion can serve as a base in a reaction with water to produce acetic acid and hydroxide ions:

CH3COO^- _(aq) + H₂O _(l) ßà CH₃COOH _(aq) + OH^- _(aq)

Similarly, the ammonia ion, being the conjugate acid of a weak base, can undergo hydrolysis with water to produce ammonia and hydronium ions:

NH₄⁺ _(aq) + H₂O _(l) ßà NH_{3 (aq)} + H₃O⁺ _(aq)

To figure out whether a salt is acidic or basic, you first have to work out whether either of its ions will undergo hydrolysis with water. In NaCl, neither Na+ nor Cl- undergoes hydrolysis, so NaCl is a neutral salt. In NH₄Cl, however, the ammonia ion does undergo hydrolysis to produce hydronium ions, as in the above equation, so ammonium chloride is an acidic salt. Similarly, salts with one ion that hydrolyses to produce hydroxide ions are basic salts.

Now, what happens when you get something like NH₄CH₃COO which has one ion that will hydrolyse to produce hydronium ions and another that will hydrolyse to produce hydroxide ions? I don't know. That's why I'm asking you. I have a hunch that it probably has to do with the extent that the two ions will hydrolyse with water, and the one that hydrolyses the most is the one that will determine whether the ammonium acetate is acidic or basic.

One person on Yahoo Answers says that it's neutral as both hydrolysis reactions have the same equilibrium constant and thus equal amounts of hydronium ions and hydroxide ions will be produced. I just need to find stuff to back that up now.

This other page just gives the vague "aqueous solutions can be acidic, basic or neutral."

I could search up a bit more but I can't be bothered right now because I'm lazy like that. Onto the next part...

The Next Part

There's no sub-heading on my course outline that I can name this section after, damnit!

Anyway. Some of the stuff here is the basic stuff in parts 1 and 2, such as the Arrhenius and Bronsted-Lowry models which I outlined in part 1. And then there's also the self-ionisation of water, which I should've explained in one of my previous sections, but I didn't, so here's a brief explanation.

Basically, water is amazing and can be an acid and a base. (Well, actually, the correct word is "amphoteric," but meh.) Water is also an electrolyte- i.e. it can split up into ions that carry an electrical charge which is why you shouldn't drop your hairdryer into a bathtub. I think that the way that it does this is mainly through self-ionisation, which is where some water molecules act as acids while others work as bases and then they react and then they all live happily ever after. (Not really. The reaction doesn't exactly go to completion. Probably a good thing too.)

H₂O _(l) + H₂O _(l) ßà H₃O⁺ _(aq) + OH^- _(aq)

In Year 12 we also take this one step further. Just like other equilibrium equations, this equation also has an equilibrium constant. At 25 degrees Celsius, this equilibrium constant is 1 x 10^(-14) (i.e. very little of the water becomes ions). However, this equilibrium constant can change with temperature.There's a little table on Chem Guide that displays the water constant at 0 degrees, 10 degrees and so on, along with corresponding pH changes.

Wait, what's that? you may ask. The pH of water changes with temperature? Sure does. If the equilibrium constant is higher, it means that more hydronium (and hydroxide) ions are being produced. This then results in a change in pH. From the table on Chem Guide, we can see that the pH seems to decrease with temperature.

One important thing to remember is that, in fact, pH does not necessarily determine acidity: it's more about the proportion of hydrogen ions in comparison to hydroxide ions. If the concentration of hydrogen ions is equal to the concentration of hydroxide ions, then water is neutral, regardless of pH. pH 7 is only neutral at 25 degrees Celsius.

Hmm... what else... Buffer solutions! Back in my post on equilibrium, I said that I had no idea what a buffer solution was. Well, now I do, so now I can explain!

Buffer solutions basically make use of solutions at equilibrium and Le Châtelier's principle in order to resist changes in pH. Adding H+ ions will shift equilibrium position to cause one reaction to speed up to partially counteract this change, and adding OH- ions will shift equilibrium position in the opposite direction to speed up the other reaction to partially counteract that change. Because the equilibrium position can shift positions like this in order to counteract changes, the H+ ion concentration doesn't change a lot and thus the pH doesn't change a lot either. Buffer solutions, however, do have a maximum limit of how much they can hold, so chucking 100L of 12 mol/L hydrochloric acid in one will still reduce the pH (unless it is a seriously strong buffer solution).

So, how do you go about making a buffer solution? Well, you simply add roughly equal amounts of a weak acid or base and a solution containing its conjugate base or conjugate acid. For example, you could form a buffer solution by mixing together acetic acid and a solution of sodium acetate:

CH3COO^- _(aq) + H₂O _(l) ßà CH₃COOH _(aq) + OH^- _(aq)

In this equilibrium solution, when OH- is added, the reverse reaction speeds up to partially counteract the increase in OH-, partially restoring the pH. When H+ is added, it reacts with OH- to form water, reducing the concentration of OH- ions and causing the forward reaction to speed up to reverse that change. Thus, pH is mainly kept the same. Buffer solutions reach their maximum capacity when one or more of the reactants are fully consumed.

Some other factors that affect the capacity of buffer solutions are the relative concentrations of the acid and conjugate base, or the base and conjugate acid. Having roughly equal amounts of the two is the best way to go. Also, higher concentrations of the substances is best for absorbing more hydronium or hydroxide ions.

And that's the main stuff on Acids and Bases! Next up is titrations. Argh. I hate titration labs. Actually I hate the theory part too but at least it's not as annoying as the labs...