Chemical Forums
Chemistry Forums for Students => Inorganic Chemistry Forum => Topic started by: wacki on June 02, 2013, 07:10:32 PM
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All of group II are negligible but beryllium. Why?
Also, why isn't Flourine negligible like it's brothers Cl, I and Br?
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I'm just trying to understand why NO3, CLO4 & CLO3 are negligible. And a really good way to drill these into my head.
NO3 is nitrate cation
NO2 is nitrite anion
I'm not sure why the above switches from + to -
Wikipedia has a great chart on chlorate and perchlorate. I assume a minimum of 3 oxygens are needed to stabilize the negative charge via resonance. 2 doesn't cut it. Other than memorization, is there any way to tell why 2 oxygens aren't good enough?
http://en.wikipedia.org/wiki/Chlorite
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What do you mean by negligible? In what way are these elements 'negligible'?
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By negligible I assume you mean non-coordinating/nonreactive? Typically halogen anions form bonds with various metals and functional groups through mostly ionic means. This means in most polar solvents (especially water) these compounds are highly liable. Fluoride's smaller size allows it to bond with much more covalent character than the other halogens - in particular is bonds especially well with silicon (such as would be present in glassware) - thus fluoride is much less liable and not "negligible".
Group two elements with their typical 2+ charges are even more likely to bond through ion rather than covalent means, beryllium's small size makes it an exception in much the same way as fluoride.
For your nitrate and nitrite question, the -ate and -ite endings in general refer to the oxidation number of the central atom. -ate denotes high oxidation number (if there are several high oxidation states the prefix per- is also used i.e. permaganate) and -ite denotes low oxidation numbers (for several low oxidation states the prefix hypo- is used i.e. hypochlorite). The common positive oxidation states of nitrogen are +7 and +3 therefore we name the compound containing +7 nitrate and the one containing +3 nitrite. In these polyatomic ions oxygen's oxidation number is -2. Really the only difficulty is only in the naming convenients.
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+7 nitrogen? Didnt you mean +5?
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+7 nitrogen? Didnt you mean +5?
Ah yes. My bad I think I was going to use permaganate as an example at first then rewrote it and forgot to change everything. Good catch.
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NO3 is nitrate cation
NO2 is nitrite anion
No no no...they are both anions (minus-one charge).
You won't be able to see why unless you have knowledge of Lewis structures.
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What do you mean by negligible? In what way are these elements 'negligible'?
Negligible as in they are used to identify strong acids. HCl is a strong acid because Cl- is negligible. Na+ is neglible so NaOH is a strong base.
Negligibles don't change the pH of water.
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NO3 is nitrate cation
NO2 is nitrite anion
No no no...they are both anions (minus-one charge).
You won't be able to see why unless you have knowledge of Lewis structures.
I understand everything here:
Lewis Structure YouTube
http://www.youtube.com/watch?v=jv55tgoyM0g
It's not really the lewis structure it's the formal charge on the individual atoms. N is supposed to have 5 but only has 4 electrons and gets a formal charge of +1. The two single bonded oxygens have a formal charge of -1. So +1 for N and -2 for 2*singleBonded Oxygen sums up to -1.
The double bonded oxygen has neutral formal charge.
I get all of that. The problem is you can't draw the lewis structure without assuming 24 electrons as a starting point.
My ultimate goal regarding the polyatomics is this: Given just PO4, SO2 and SO3... what is the actual charge? How do you figure it out? You only have the periodic table as a guide.
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I really wish the mod hadn't merged these threads. These are two different questions.
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Leaving PO4 aside for a moment, SO2 and SO3 can both be either negatively charged or neutral compounds in common practice.
You can't figure out the charge from the formula. You have to state the formula and the charge and draw the structure from that.