One of the things that trips me up in chemistry is getting my head around which element is oxidised and which is reduced and which one is the oxidant and which one is the reductant in a redox reaction. They're not terribly difficult concepts really, but it's easy to get them mixed up. It's also pretty easy to get all mixed up between which process takes place at which electrode when a substance is undergoing electrolysis. So here I'm going to compile as many different ways as I can think of as remembering which one is which. Having a variety of different methods might just make you more confused, though, so read on with caution!
If you have absolutely no clue what I'm talking about, then look at my pages on redox equations and electrolysis.
Anyway. Here we go with all the mnemonics and whatnot:
AN OIL RIG CAT
This is probably the one you'd have been taught in your Chem class. OIL stands for Oxidation is Loss [of electrons], and RIG stands for Reduction is Gain [of electrons]. So, to find out whether something is being oxidised or reduced, just work out of the substance is losing or gaining electrons and then apply this mnemonic. Another nifty thing this mnemonic does is it helps you remember where oxidation and reduction takes place in electrolytic cells or whatever they're called. You see, AN stands for anode, and it's where Oxidation (which Is Loss) takes place. CAT stands for cathode, and it comes directly after RIG as the cathode is where Reduction (which Is Gain) takes place.
Reduction = Reduced [oxidation number]
This is the way I remember it. A substance is being reduced if its oxidation number is reduced. A substance is being oxidised if its oxidisation number is increasing (higher oxidation number = more oxidation?). Then you just have to remember that the reducing agent is the thing being oxidised and vice versa. (I find that once I work out what's being oxidised and reduced, it's much easier to work out the oxidant and reductant.)
Cathode = Cations
One way that I remember that stuff gets reduced at the cathode is that cations (positive ions) go to the cathode. Since the ions are already positive, they'd probably be wanting to gain electrons, meaning that there's a reduction taking place. Similar reasoning can be applied for the anode: anions (negative ions) go to the anode, and they probably would be wanting to lose some electrons rather than gain them. Hence oxidation takes place at the anode.
Monday, July 29, 2013
Saturday, July 13, 2013
Titrations- Basics of Titrations and Acid-Base Titrations
Argh. Titrations. I. HATE. TITRATIONS. But titrations are part of studying chemistry, so I guess I'll just have to get over it.
Basically, in titrations, you have one solution with a known concentration, and another solution of an unknown concentration that you want to find the concentration of. To do so, you basically get a known quantity of one solution, like 20mL, and then you add gradually increasing quantities of the other solution until you get some kind of indication that the reaction has gone to completion. Then you use all the stoichiometry stuff learned from before to work out the concentration of the unknown substance.
Now here comes the kicker. Titration lab work. Ugh.
As I mentioned, you need to have a solution with a known concentration. And the best way to know what the exact concentration of your solution is is to make it yourself so that you know exactly what goes inside it.
The problem? A lot of caution is needed to make sure nothing gets contaminated. Also, you need to decide on a good solution to make up. Not all solutions are equal for this purpose. The best substances are those known as primary standards. Primary standards have characteristics that make them good for making up into solutions of known concentrations: they are pure and aren't deliquescent (i.e. don't absorb water). They are anhydrous and don't react readily with stuff in the atmosphere. It's also good if the substance has a relatively high molar mass, which will make it easier to measure out one mole (or however much you need) with a greater accuracy.
Making a Solution of a Primary Standard
The first step in making a primary standard solution is working out how much solid needs to be dissolved. To do this, take a look at my post on simple calculations involving moles. Measure this amount out in a petri dish (remember to put the petri dish on the scales and tare before putting any solid on) and then tip the solid into a beaker (make to rinse the beaker with distilled water first to avoid contamination). Dissolve it with a little bit of distilled water in here. Remember, you must use distilled water, otherwise you'll have other random ions floating around in your solution which you don't want. After dissolving, first rinse a volumetric flask of whatever size you need with distilled water, then pour in the solution and make it up to the mark with some more distilled water. Remember that the bottom of the meniscus must be level with the mark (when read from eye level). Then put a stopper on and swirl the solution around a bit so that everything's all nice and mixed around. Finally, get a storage bottle, and rinse a couple of times with distilled water, and then with a little bit of the solution. Now that's done, pour in the rest of the solution. You have now made up a solution with a known concentration. Yay.
The Titration Part: Acid/Base Titrations
Now that you have a solution of a known concentration, it's time to titrate it with something to make all that hard work worth it. There are different types of titrations- I'm just going to start off with acid/base ones because we've been doing acid/base titrations in class.
What you'll need: a pipette (a reasonably large one that can hold 20-25mL, not one of those puny ones), a burette (kind of like a long plastic pipe with measurements and a little tap thing on the end so you can control how much liquid comes out), some conical flasks, a retort stand and burette holder. You'll probably want one or two small beakers so that you can pour stuff into there before transferring it to the burette or conical flask. Oh, and you'll also need some indicator for acid/base titrations. The type of indicator depends on what you're titrating- more on that later.
I really should draw a picture of something of how to set it up. Better yet, I could take a photo so that you can see all the equipment properly. It would probably help if I could be bothered to sync my phone up to my computer though. Ah well.
Anyway... basically what you do is first you rinse the pipette with whatever you're going to put into the conical flasks, the burette with whatever you're going to put in the burette, and the conical flasks with distilled water. Also you should rinse the little beakers with whatever you're planning to put in there. (I'll explain the reasons behind all the rinsing later.) Then you use the pipette to put a certain amount of one solution into the conical flask, and a small funnel to put the other solution into the burette (you don't have to fill all the way to the 0mL mark, but you can't go over). Write down the initial reading. Put a few drops of indicator into the conical flask.
Now you're ready for your rough titration. Turn on the tap of the burette and swirl the flask around. Occasionally, stop pouring stuff in and stop swirling the flask to see if there's a permanent colour change throughout the whole liquid. Stop titrating when you get a permanent change. Write down the reading on the burette and work out how much of the solution was consumed. Now, do it again, except stop titrating a few mL before you use up the amount that you used before (e.g. if you needed 20mL last time, stop titrating once you've used up 17 or 18mL). Now go more slowly, drop by drop, until you work out the exact drop needed to get a permanent colour change. Do this a few times. When you have more or less the same amount consumed for multiple titrations, choose this amount for your calculations.
Now I've explained what a titration is, I'm going to explain the finer points in a bit more detail!
Rinsing
What do you rinse with what, and why? Here's a guide:
When preparing a solution
Basically, in titrations, you have one solution with a known concentration, and another solution of an unknown concentration that you want to find the concentration of. To do so, you basically get a known quantity of one solution, like 20mL, and then you add gradually increasing quantities of the other solution until you get some kind of indication that the reaction has gone to completion. Then you use all the stoichiometry stuff learned from before to work out the concentration of the unknown substance.
Now here comes the kicker. Titration lab work. Ugh.
As I mentioned, you need to have a solution with a known concentration. And the best way to know what the exact concentration of your solution is is to make it yourself so that you know exactly what goes inside it.
The problem? A lot of caution is needed to make sure nothing gets contaminated. Also, you need to decide on a good solution to make up. Not all solutions are equal for this purpose. The best substances are those known as primary standards. Primary standards have characteristics that make them good for making up into solutions of known concentrations: they are pure and aren't deliquescent (i.e. don't absorb water). They are anhydrous and don't react readily with stuff in the atmosphere. It's also good if the substance has a relatively high molar mass, which will make it easier to measure out one mole (or however much you need) with a greater accuracy.
Making a Solution of a Primary Standard
The first step in making a primary standard solution is working out how much solid needs to be dissolved. To do this, take a look at my post on simple calculations involving moles. Measure this amount out in a petri dish (remember to put the petri dish on the scales and tare before putting any solid on) and then tip the solid into a beaker (make to rinse the beaker with distilled water first to avoid contamination). Dissolve it with a little bit of distilled water in here. Remember, you must use distilled water, otherwise you'll have other random ions floating around in your solution which you don't want. After dissolving, first rinse a volumetric flask of whatever size you need with distilled water, then pour in the solution and make it up to the mark with some more distilled water. Remember that the bottom of the meniscus must be level with the mark (when read from eye level). Then put a stopper on and swirl the solution around a bit so that everything's all nice and mixed around. Finally, get a storage bottle, and rinse a couple of times with distilled water, and then with a little bit of the solution. Now that's done, pour in the rest of the solution. You have now made up a solution with a known concentration. Yay.
The Titration Part: Acid/Base Titrations
Now that you have a solution of a known concentration, it's time to titrate it with something to make all that hard work worth it. There are different types of titrations- I'm just going to start off with acid/base ones because we've been doing acid/base titrations in class.
What you'll need: a pipette (a reasonably large one that can hold 20-25mL, not one of those puny ones), a burette (kind of like a long plastic pipe with measurements and a little tap thing on the end so you can control how much liquid comes out), some conical flasks, a retort stand and burette holder. You'll probably want one or two small beakers so that you can pour stuff into there before transferring it to the burette or conical flask. Oh, and you'll also need some indicator for acid/base titrations. The type of indicator depends on what you're titrating- more on that later.
I really should draw a picture of something of how to set it up. Better yet, I could take a photo so that you can see all the equipment properly. It would probably help if I could be bothered to sync my phone up to my computer though. Ah well.
Anyway... basically what you do is first you rinse the pipette with whatever you're going to put into the conical flasks, the burette with whatever you're going to put in the burette, and the conical flasks with distilled water. Also you should rinse the little beakers with whatever you're planning to put in there. (I'll explain the reasons behind all the rinsing later.) Then you use the pipette to put a certain amount of one solution into the conical flask, and a small funnel to put the other solution into the burette (you don't have to fill all the way to the 0mL mark, but you can't go over). Write down the initial reading. Put a few drops of indicator into the conical flask.
Now you're ready for your rough titration. Turn on the tap of the burette and swirl the flask around. Occasionally, stop pouring stuff in and stop swirling the flask to see if there's a permanent colour change throughout the whole liquid. Stop titrating when you get a permanent change. Write down the reading on the burette and work out how much of the solution was consumed. Now, do it again, except stop titrating a few mL before you use up the amount that you used before (e.g. if you needed 20mL last time, stop titrating once you've used up 17 or 18mL). Now go more slowly, drop by drop, until you work out the exact drop needed to get a permanent colour change. Do this a few times. When you have more or less the same amount consumed for multiple titrations, choose this amount for your calculations.
Now I've explained what a titration is, I'm going to explain the finer points in a bit more detail!
Rinsing
What do you rinse with what, and why? Here's a guide:
When preparing a solution
- Before making a solution up to the mark, rinse everything with distilled water. The petri dish should be dried before weighing so that you don't have water stuffing up your calculations, but otherwise it doesn't matter so much for everything else as you're going to be dissolving the solid and diluting the solution anyway.
- When you're storing the solution, you'll first have to rinse the storage bottle a couple of times with distilled water to get rid of any other contaminants already in the bottle before rinsing with a little bit of the solution. Rinsing with the solution gets rid of the last little bits of distilled water and ensures that the concentration stays more or less the same. After all, you wouldn't want to do all that hard precise work and then have it all crushed by something stupid like not rinsing the storage bottle correctly, would you?
When doing a titration
- The pipette and burette need to be rinsed with the solutions that are about to go in them so the concentration stays constant and thus allows you to calculate the number of moles of each substance in whatever amount of solution that you end up using. Ditto with the little beakers if you're using them to store solutions before putting them in the pipette or burette. (Make sure you know which beaker's which!)
- The conical flasks need to be rinsed with distilled water to keep the number of moles of the substance constant. If you rinse with solution, you increase the number of moles of the substance in the conical flask, which will stuff up your calculations.
Indicator Choice
Okay, firstly I have to talk about the "end point" and "equivalence point" of a titration.
The "end point" is the point at which you stop titrating because of a colour change, and the "equivalence point" is the point where the reaction has actually stopped. Naturally, you want to choose an indicator that will change colour close to the equivalence point.
For example, let's look at the reaction between NaOH and HCl. This reaction is complete when the pH is around about 7. One indicator that you might want to look at is phenolphthalein, an indicator that is colourless for low pH and pink at high pH. The colour change occurs at around about pH 8, which is relatively close to pH 7. You might ask why not use an indicator that changes colour at pH 7, but the thing is, with strong acids and bases, one small drop can cause a massive change in pH when you're close to the equivalence point, so it's generally better to have an indicator that changes slightly early. (I think.)
There's different kinds of indicators which are useful for different titrations but you have to pick your indicator carefully, otherwise you'll stop titrating too early or too late and your calculations will be stuffed up. Don't get too hung up on this, though, as there are lots of times where several different indicators might be appropriate. For example, I think that it's also acceptable to use methyl orange for the above experiment, which changes colour at around pH 4.
And that's all for now on titrations! I have a practical test on titrations when school starts again... argh
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, O2- 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:
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, O2- 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 NH3 (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...
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