Can you PLEASE tell me what this is all about; I'm sure we never learned it in class.
An aqueous solution of a Group II metal chloride,XCl2, forms a white precipitate when dilute
aqueous sodium hydroxide is added. A separate sample of the solution of XCl2 does not form
a precipitate when dilute aqueous sodium sulphate is added.
An aqueous solution of a different Group II metal chloride, YCl2, does not form a precipitate
when dilute aqueous sodium hydroxide is added. A separate sample of the solution of YCl2
forms a white precipitate when dilute aqueous sodium sulphate is added.
Suggest identities for the Group II metals X and Y. Write equations, including state symbols,
for the reactions which occur.
OK, don’t panic!
Before I go into more detail, I hope I don’t embarrass you if I go over really simple things first. I find a lot of students struggle to get these concepts down, but are too scared to ask, so it’s good to picture some terminology, first.
Remember where your metals are. Most of the periodic table is metals--you should memorize where they are! Look for the staircase line dividing the metals from the nonmetals cutting off the northeast corner of the periodic table. Some periodic tables show this better than others. All elements to the right of this line are nonmetals, and so is the oddball hydrogen. Everything else is metals. (There are some oddballs on the staircase line that have intermediate properties called metaloids, but your question isn’t dealing with them.) It is really helpful to know whether you are dealing with a metal or a nonmetal, since they do different things.
Remember that metals come in two forms, generally. They are either uncharged, and then look like the shiny metals we know, and we call those elements. Or, they lose some number of electrons, and get a positive charge. I have never seen one with a negative charge. It is really helpful to know that if a metal is an ion, it will be positive, not negative! So, a metal will most likely either be an element, or a positive ion (a cation).
Some, like groups I and II (first and second columns of the P. table) always have predictable charges of +1 and +2, thank God (when they are charged and not elements, that is). Many of the rest of the metals have more than one charge possible, so in doubt look them up or try to figure them out based on what the charge of the anion they are with and the formula of the salt. (The charges of anions are fortunately always predictable--so use them to figure out the charge of the cation they are with.) Your question has us distinguishing between group II metals, which narrows it down considerably. It talks of them as chloride salts, meaning they are joined in rock like fashion to Cl-, so it makes sense that their formulae are XCl2 and YCl2, because X and Y will both have +2 charges.
The term “salt” is slang for ionic compound, so we have a positive ion and a negative ion stuck together.
When text books say “chloride salt” or “group II salt” they are assuming that you know that this is an ionic compound containing chloride, or a salt containing a group II metal ion, respectively. You would be surprised how many people don’t get that right off that bat, and I blame textbooks for not being more clear on that point.
All salts all solids when they are dry (OK, some people have invented some weird organic ones that are weird liquid solvents, for those nit pickers out there, but those are very, very esoteric!) So, I like to tell my students to memorize that:
“All ionic compounds are rocks”. Solid. Solid solid. Picture that?
That is not to say that all solids are ionic compounds! No no! Just because all cats have tails does not mean all things with tails are cats. That is called reductionist logic, and it leads people into illogical horrors. Just remember that all ionic compounds--”salts”--are solid when NOT in water.
OK, next thing to picture:
When you dunk a salt in water, SOME of them fall apart into positive and negative ions, and since these ions are too small to see with your eyes, they seem to disappear, and we say they “dissolve”.
SOME ionic compounds, however, do NOT fall apart into separate ions in water. They remain like a rock in water. Picture a rock in water. This is what we call insoluble. Sometimes they can fool you, and the water looks like milk or there is scum on top or the bottom of the container--this is still not “dissolved.” If something is truly dissolved, we can call the liquid a true “solution”. So the term “solution” means something special.
If there is a precipitate (solid thing) in the liquid, or it looks milky or if there is solid dreck on top or the bottom, that is not a “true solution”. It might be called a colloid or slurry or something else.
First, we have to guess which metal is coupled with the chloride, but the question says “group II metal” thank goodness, so that narrows it down quite a bit (to Be, Mg, Ca, Sr, Ba, and I doubt you would be dealing with the radioactive Ra-Radium...but that’s the whole list.)
The bottom line is that this question expects you to use your knowledge of solubility of ionic compounds, to sort out whether the unknown metal is Mg++ or Ca++ or whatever, since of of these is soluble when coupled with OH-, one is not, and one is soluble when coupled with sulfate, while the other is not.
So how do you know which will dissolve in water, and which won’t? Well, you have to memorize, or at least look up, the famous “Solubility rules”. I make my students memorize them (after an apology), and they grumble, but I wouldn’t do it to them if I didn’t think it wasn’t helpful to them (and their patients) in the long run.
It’s always a pain to memorize solubility rules, but the reason there are these rules is because predicting what is soluble and what is not is too tricky because there are too many competing factors for solubility, which I will describe later, below.
But before we go into the nitty gritty, let’s just look at these damn rules: You can look them up yourself by googling "solubility rules".
Thanks to Professor Kenneth W. Busch from whose Web page these data were extracted.
1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are rare. Salts containing the ammonium ion (NH4+) are also soluble. 2. Salts containing nitrate ion (NO3-) are generally soluble. (Holly’s addition: also it is helpful to note that acetate salts CH3CO2-are usually soluble.) 3. Salts containing Cl-, Br-, I- are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble. 4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver; virtually anything else is insoluble. 5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4, and CaSO4. 6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble. 7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble. 8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and Ba) are insoluble. Some other insoluble carbonates include FeCO3 and PbCO3. 9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO4 10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag2PO4 11. Fluorides are frequently insoluble. Examples: BaF2, MgF2 PbF2. This offering of solubility rules is in the public domain and may be copied without restriction. The user is encouraged to download it for private use and public distribution in any form, including that of making it available on other Web servers.
Your question just has us dealing with sulfate and hydroxide salts. Oh, first we deal with the metal ion stuck to a chloride, and the chloride ion isn’t very sticky in water, (unless it encounters postively charged Hg, Ag, or Pb). You can see this looking at rule 3 above.
So we know for sure that as soon as the metal stuck to the chloride (XCl2 or YCl2) hits the water it will fall apart. That is very convenient, otherwise, our group II metal would be stuck in a chloride salt and unable to stick to anything else we might test it with!
We can then forget about the chloride ion, thank goodness, don’t give it another thought. The question is, what will the metal ion stick to, next?
Hmm, I see now that the rules above for hydroxide salts (rule 6) aren’t very detailed. If you look up different versions of these rules on the internet you will get different levels of details. What I know from being an old lab hag, and from looking at “milk of magnesia” (Mg(OH)2) is that magnesium hydroxide does not dissolve in water. It makes a milky not-a-real-solution called a colloid, which you have to shake before you drink it up, as an antacid.
I also know that calcium hydroxide DOES dissolve in water, and creates a basic solution that people call lime water. More detailed solubility rules will tell you that solubilities of group II metals with OH- is not an either/or thing, but increases as you go down the group.
OK, let us look at the sulfates (rule 5) (don’t confuse them with sulFIDES!)
I also recall that Mg sulfate dissolves in water, and the rule above says Ca does not. What more detailed rules will tell you is that the solubility of sulfates with goup II metal ions DECREASES as you go down the group.
So, based on your info, I would say that X is Mg and that Y is Ca.
So why are the solubility rules so hard to predict, thus making generations of grumbling students memorize them?
The stickiness of two ions (or any positive and negative charge for that matter) depends on two things, as a consequence of something called Coulomb’s law.
The first is the distance they are separated. Like gravity, the force of attraction decreases with the square of their distance. So if you have two big fat fluffy ions trying to get close, they will not stick very well, because of their size. Small ions can get closer, so they may be stickier.
But wait, here’s another factor. The second thing determining the stickiness of a charge is the magnitude of the charge. (In my mind this is analogous to mass in gravitational attraction.) For example, a +2 charge is more affected by charge than a +1 charge, and a -2 charge is more affected by charge than a -1 charge. In short, the closer a charge is to zero, the less tugged around it will be by electrical forces.
But we have even more to consider than just the force of the two ions sticking to each other in the salt, which is called the lattice energy. We also have to consider the “energy of hydration”, which is the attraction these separate ions will feel for water. Yes, water does not have a net charge, but its hydrogen has a partial (fractional, if you will) positive charge, and its oxygen has a partial negative charge. It has uneven charge distribution over its surface, or you could say it is “polar”. Even though it’s overall charge is zero, parts of a water molecule can attract either a positive ion or a negative ion, by Coulomb’s law.
So, you see, there are all these competing factors for dissolving--the less sticky two ions are with each other and the more attracted they are to water, they more likely they will dissolve.
I hope that helps!