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Author: Subject: Are all s-block elements metals?
r15h4bh
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[*] posted on 25-12-2012 at 19:30
Are all s-block elements metals?


The question is: "Are all s-block elements metals?"

My answer would be no, since Helium isn't and neither is Hydrogen and they're both s-block elements. But the answer given in the book is "Yes, all s-block elements are metals."

Who's correct? :)
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bfesser
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[*] posted on 25-12-2012 at 19:48


If I recall correctly, <a href="http://en.wikipedia.org/wiki/Metallic_hydrogen" target="_blank">solid hydrogen</a> <img src="../scipics/_wiki.png" /> has metallic properties. No clue about <a href="http://phys.org/news137255843.html" target="_blank">helium</a> <img src="../scipics/_ext.png" />.

[Edited on 7/9/13 by bfesser]




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Oscilllator
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[*] posted on 25-12-2012 at 20:06


can you eve get solid helium? Pretty sure it just goes superfluid and all sort of other weird stuff.
So I vote no, helium is not a metal.




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elementcollector1
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[*] posted on 25-12-2012 at 20:07


Helium can't be solid because of zero point energy, or something.



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r15h4bh
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[*] posted on 25-12-2012 at 20:14


But at room temperature it isn't a metal and I think that's what they're asking for. They weren't asking whether all s-block elements can show metallic properties they just asked are they metals. So I think I'll go for "No." Thanks :D
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bfesser
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[*] posted on 25-12-2012 at 20:40


A metal is defined as such by it's metallic (electronic) properties. I would say it's just a matter of opinion.

[Edited on 12/26/12 by bfesser]




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r15h4bh
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[*] posted on 25-12-2012 at 20:52


So what would you say it is?
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platedish29
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[*] posted on 29-12-2012 at 20:39


I'll take a position by bfesser: as an element, He2+ ejected from a metal's radioactive decay is a metal once in its life time, until it leaves the metallic surroudings.
Maybe if yu bring a huge ball of solid metal inside which a radioative sample is put, I don't think Helium will be kindly able to escape and invariably would have to stand his way. If you can't win, you lose lol
Semiconductors start to show more of the metallic properties upon applying certain conditions, so why helium should not?
Between, He been continuously held in this state is anotehr story.
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[*] posted on 30-12-2012 at 04:34


Helium can be made into a solid. It was used to examine claims of supersolidity.

As to the OP, I would say "No, not all s block elements are metals", but include the caveat that it depends on the temperature ranges and time scale one considers.
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DJF90
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[*] posted on 30-12-2012 at 15:49


For a start, I'd like to point out helium isnt in the S block, and neither is hydrogen! (although H is commonly placed atop group I).

S block consists of groups I & II - in periods; Li Be, Na Mg, K Ca, Rb Sr, Cs Ba, Fr Ra. All of these elements are metallic at STP (thats Standard Temperature and Pressure).

There, that wasn't too hard, was it.
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Vargouille
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[*] posted on 30-12-2012 at 17:09


You could very well consider helium and hydrogen to be part of the s block. Their electronic configurations are analogous to the other members of the s block (ie their valence electrons occupy the 1s orbital).
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[*] posted on 30-12-2012 at 18:06


The problem with what you say is that elements are grouped based on their chemical reactivities.Helium is certainly a group 8 element in the p block, not because its 1s shell is occupied, but because it has a stable electron configuration (a closed shell) akin to its sibling elements. Any periodic table that puts it in the s-block is pure waffle. Hydrogen is unique in its electronic configuration, and no other element resembles its chemistry, and as such doesn't belong to any group in the periodic table. It often floats alone at the top somewhere. Helium always sits atop the rest of the noble gases.
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[*] posted on 30-12-2012 at 18:26


It isn't, strictly speaking, true that elements are grouped based on chemical reactivities, helium being the obvious exception. Lead and oxygen are in the same group, and they have quite disparate chemistry. Moreover, the most recent elements have not been produced in quantities large enough to demonstrate their chemical reactivities, thus their places on the periodic table are chosen to be in line with their predicted electronic configurations. It is, however, strange to me that while it would be acceptable to include lead with the p block due to its highest energy orbital being a p orbital, it is not similarly acceptable to include hydrogen and helium in the s block for analogous reasons.
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[*] posted on 31-12-2012 at 04:49


And what periodic table are you looking at? Lead (Pb) is an element in group IV (sometimes called group 14); Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), and Lead (Pb). Oxygen is two columns to the right of Carbon in Group VI. I agree that new elements are laid out in terms of electronic structure, but back in Mendeleev's day he constructed the table based on the similarities between the elements, their physical and chemical properties and also on appearance, and from this he could predict missing elements and their properties quite well (I believe Germanium is a common example of his predictive work).
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[*] posted on 31-12-2012 at 05:44


Oh my, I have made a grevious error. I should have said that the chemistry of lead and of carbon are quite different. In any case, one should look at Mendeleev's periodic table, and then the modern periodic table. While Mendeleev got many groups correct, he made a few flaws. Lead, tellurium, gold, mercury, and probably a few others. I'll grant that he didn't know about a large number of elements, and had he known, he would have rearranged the periodic table, but he put lead under the alkaline earth metals for a reason. My guess is that he recognized that it, like the alkaline earth metals, formed an insoluble sulfate. His later one, while moving it out of the alkaline earth metals, did not put it with carbon. I think. His 1871 version is a little hard to parse. The placement of the actinides in the f block is also from electronic configuration (of which thorium and uranium were known to Mendeleev).
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[*] posted on 31-12-2012 at 06:37


The chemistry of lead and carbon are not so different; though it appears so if you only consider the salt forming ability of the element. Both form monoxides and dioxides, CO/CO2 and PbO/PbO2 respectively, though lead forms the mixed valency oxide Pb3O4. Both form tetrahalides (CX4, PbX4, although the latter is hydrolytically unstable), as well as what can be considered dihalides (transiently in the case of carbon, :CCl2 (dichlorocarbene)). Both elements can form hydrides containing an E=E double bond (where E = C, Si, Ge, Sn, Pb), i.e. Ethene H2C=CH2 and the corresponding Lead compound H2Pb=PbH2. There are also alkyl/aryl substituted diplumbenes (http://onlinelibrary.wiley.com/doi/10.1002/%28SICI%291521-37...). Lead even goes as far as forming a triple bond to itself, although in comparison to Carbon, the substituents are not colinear with the Pb-Pb triple bond (http://pubs.acs.org/doi/abs/10.1021/ja993346m). Tin and Silicon also form such compounds, although I'm sure about Germanium.

You can't expect Mendeleev to have gotten everything right, and his attempts at organising the elements (especially considering much of them were undiscovered) was formidable. In the first link you provide, it makes sense that he put lead with the alkaline earths. The nitrates all decompose to give the metal oxide and NO2. The sulfates as you mention are insoluble. Stable dihalides are formed, although those of lead are mostly insoluble, contrasting the rest of group II. Notice he also put Thallium with Group I; this is probably due to the similarity of K+ and Tl+ (which is partly what makes Tl salts so toxic!) although the 3+ oxidation state of thallium contrasts that of the alkali metals and reveals its place in group III. Also notice how he placed Bismuth in the right place, probably due to the Bi (III) and Bi (V) oxidation state: Such valencies are repeated oft in group (V); in fact all elements in group V have a selection of compounds in the +3 and +5 oxidation states. It is not suprising that the rare earth elements caused him some confusion. He placed Uranium and Gold in group three, probably because they both commonly form trivalent cations (as do most of the rare earths, actually). I'm not going to rationalise the rest of his decisions. The 1871 version you link to clearly shows lead in group IV, underneath carbon. Titanium and Zirconium are also included in the group, probably due to the prevalence of the +4 oxidation state (e.g. Oxides, chlorides etc.). All in all I think he did a good job of it, based on how little chemistry was actually known at the time.
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[*] posted on 31-12-2012 at 07:47


Regarding the chemical similarities between lead and carbon, I will accede that I was ignorant of the bonding of lead in organometallic compounds. In any case, it is somewhat of a digression from the point in question: Can hydrogen and helium be considered s block elements? The separation of the transition metal elements into the d block is due to their highest energy orbital being a d orbital. That of the lanthanides and the actinides is the f orbital, placing them in the f block. The elements of the s block, whether or not you decide to include hydrogen or helium, have an s orbital as the highest energy orbital. The elements of the p block have a p orbital as their highest energy orbital. By this definition of the blocks, hydrogen would be considered an s block element. As for helium, it is placed as a noble gas due to its inertness, however, because it does not have electrons in the p orbital at its ground state, it would be odd to consider it in the p block.
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[*] posted on 31-12-2012 at 08:36


The question is not can they be considered to be, but are they s-block elements. The answer is no. Your reasoning does not take into account a couple of things.

You say that Helium should be in the s-block because its highest occupied orbital is an s-orbital. However, for a ground state configuration, Helium does not have an accessible p-orbital, and neither does hydrogen, which makes them very unique in that respect. However, one thing that all group VIII elements have in common is not the full p-orbital, but a full atomic energy level, whether this is 1s2 (in the case of Helium, n=1 is filled) or 1s2 2s2 2p6 (in the case of Neon, n=2 is filled). This is why they are so inert, because the atomic level is full. Reactivity increases down the group (such that Xenon has quite extensive chemistry) because of the decreasing effective nuclear charge; it's easier to remove electrons from the outer shells when they're further from the nucleus and sheilded from its positive charge by the other electrons in lower orbitals.

Similarly, Group I elements lose one electron to attain an electronic configuration in which the appropriate atomic energy level is full, e.g. for Na+ the n=2 atomic level is full. Now consider hydrogen. When hydrogen loses its electron, it has the configuration 1s0, i.e. it has NO full atomic energy levels (in fact it has no electrons whatsoever- what exists is a proton). This is unique only to hydrogen (and hydrogenic species like He2+ etc.), and so it cannot be considered alongside the alkali metals. It deserves its own group, but is usually chucked in with Group I, erroneously.
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[*] posted on 31-12-2012 at 08:47


Place of hydrogen in the periodic table is quite vague, many textbooks solve this by not placing hydrogen in any of the groups! Hydrogen has only one 1s electron, so it should be in the group 1. On the other hand hydrogen is only a one electron short from a full 1s orbital so what about group 17... And wait hydrogen has also a half filled outer orbital so it might also belong to group 14. And ... yeah, this is really a question of defining stuff.

[Edited on 31-12-2012 by kavu]
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[*] posted on 1-1-2013 at 11:03


Quote: Originally posted by kavu  
Hydrogen has only one 1s electron, so it should be in the group 1. On the other hand hydrogen is only a one electron short from a full 1s orbital so what about group 17... And wait hydrogen has also a half filled outer orbital so it might also belong to group 14
This is the beginning of a sensible answer to this question. First, though, a reductio ad absurdum. If hydrogen really belonged in Group I, it would have similar enough chemistry that you'd see alkalicarbon molecule analogous to hydrocarbons. And you don't see that.

Hydrogen forms covalent bonds readily and Group I alkali metals do not, because of hydrogen's unique electron configuration. With principal quantum number n=1, the only possible angular momentum number is L=0, and so a full shell is just two electrons. Thus adding a single electron to the vicinity of the hydrogen enables covalence. Such is not the case for n &geq; 2, since you have the possibility of L=1 electrons. So while hydrogen has similarities in terms of ionic states (+1), that's not all of hydrogen's chemistry. The s-block should properly be defined with the additional condition that n &geq; 2.
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[*] posted on 2-1-2013 at 15:58


Ah Watson, trust you to spoil the fun by invoking quantum mechanics. Its been a couple years since I did any so I tried to avoid it, especially given this is an amatuer forum (and "Beginnings"!).
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[*] posted on 3-1-2013 at 07:54


Quote: Originally posted by DJF90  
Ah Watson, trust you to spoil the fun by invoking quantum mechanics.
Sometimes a simple question doesn't have a simple answer. All that really means is that it's an interesting question. It's particularly relevant for Beginnings, where the lesson should be learned early that while many things are predictable and form clear patterns, nature as a whole is not that way, and one should expect that one's ordinary expectations to be periodically confounded.
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[*] posted on 26-3-2013 at 15:16


this is a great table for placement of elements.
they actually place H/He twice !

http://www.perfectperiodictable.com/default.html

I really like how the P/S orbitals work on this table.



[Edited on 26-3-2013 by morganism]
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