Solve Mystery of Titanium Trichloride's Electron Structure

[Dual Op Amp]

It seems we're not going anywhere with this. If only sub-shells were filled to stabilize an atom, then why does the octec rule work? If that were true, then carbon could simply make two bonds, and then it would be stable. All it would have to do is empty it's P subshells. That's not what happens, instead it tries to fill the entire second shell. So, why it titanium so different. Like I said, I've learned that an entire SHELL, not sub-shell, but SHELL has to be filled, in order for it to stabilize. So, where am I wrong? Do you have any sites to where I can learn some of this?

 

[Astronuc]

Welcome to the world of Transitional Metal Chemistry. Unfortunately, I have not found a website that goes into sufficient detail, but I did find a number of university chemistry courses that a uniquely devoted to transitional metals and their chemistry.

The notion that a full shell provides stability is supported by the fact that the nobel gases, e.g. He, Ne, Ar, Kr, Xe are stable with completely full primary shells, and full p subshells - in fact they are monatomic gases - unlike diatomic N₂ and O₂, or F₂ or Cl₂.

Carbon cannot be compared to the transitional metals, since it does not have any 'd' orbitals. It is the presence of the 'd' and 'f' suborbitals (subshells) that allows for the multi-valent ions of the transitional metals.

In forming compounds, it is not just the cation (+ ion) that affects the structure of the molecule, but also the anion (- ion), and in complex molecules or coordination compounds, it is the type of ligand (atoms or molecules surrounding a central metal ion) that affects the metal ion valence.

I will see if I can find some appropriate references, and I will get back to you on carbon.

 

[Dual Op Amp]

No wait, let me rephrase my question, how is argon stable, when it still has the d orbital to fill. If I know this, I would be able to understand Ti.
argon is Titanium without the 4s and 3d.
According to what I've learned, an entire shell must be filled. Argon avoids this rule too.

 

[Dual Op Amp]

AHAA!
Argon can be stable, because the 4S orbital interupts the 3D orbital. This, for some weird reason, allows argon to be stable. Which allows titanium tetrachloride to be stable.
But, what I don't completely understand it how does a lower 4S make it stable?
Hmm...

 

[chem_tr]

Hello

Single electrons in a whole orbital tend to be thrown out to give a more stable electronic configuration; that is why potassium with electron configuration [Ar]4s¹ throws one electron to reach the most stable configuration, [Ar].

Titanium(III) chloride is a classical d¹ ion, thus very eager to give away its single (and somewhat troublesome) electron, to reach titanium(IV) ion, again [Ar].

So why is argon so stable, thus not reactive at all? Its basic electronic shells are COMPLETELY filled, with no half-filled orbitals. If you manage to bombard with an electron, you'll end up argon again, after a system interconversion, Ar^-\longrightarrow Ar+e^-.

Let me show the other alternative, namely, pulling out an electron; in this case, a cation like Ar⁺ will be formed, but this process is extremely energy consuming, so the negligible amount of ionization energy is very eagerly satisfied by the system, to give finally the most stable form, Ar: Ar^++e^-\longrightarrow Ar

You are okay to think that 3d orbital can be filled, and if you manage to do this (as this will be VERY hard, due to 3d orbitals of extremely higher level than that of 4s and 3p), with enough number of electrons, the super anion (probably be 7-) might be somewhat stable, but another problem occurs here; argon is argon, with its small radius will be very unstable with this number of extraordinary electrons; an unusual value of \displaystyle \frac {charge}{radius} will not possibly cause the anion to be stable. You must keep this anion at extremely low temperatures to avoid its decomposition (here, ionization energies will be very low to give away these extra electrons, finally meeting at the usual place, [Ar])

Hope this helps.

 

[Dual Op Amp]

Well, Mr. Chem tr, it's not going to be an anion, when it only covalently bonds, but this is a metal to non-metal, so it would be an ionic bond. But, the octec rule should only apply to atoms with no more than two shells. N = 2. But, it doesn't. In titanium, it tries to empty the 4S and the 3D. This doesn't make any sense, an atom is stable when the ENTIRE shell is filled. To do that, titanium would have to fill 3S, 3P, and 3D. But, it doesn't, that's where I'm stuck.

 

[Dual Op Amp]

Okay, I found an even better way to phrase it.
An atom is stable, when the entire valence shell is filled.
So, why would an atom with three shells be stable, when it still has it's D orbital to fill?
Example: Argon. 1S^2, 2S^2,2P^6,3S^2,3P^6
The D orbital is not filled, yet argon is stable!

 

[Gokul43201]

An atom is stable, when the entire valence shell is filled.

Where did you read this sentence ?

That statement is not correct.

It is commonly used because it applies to all atoms/ions up to [Ne], and these include many of the common elements like H, C, N, and O.

The correct statement is that an atom is most stable when it has 8 electrons (2 for atoms/ions with 2 electrons or less) in its outermost shell. This is known as the octet(/duet) rule... which I believe you're aware of. So how is it that you misquoted the rule ?

Additionally, keep in mind that stability is not a binary condition, and is always relative to something.

 

[Dual Op Amp]

Well, I've read it everywhere.

The association between stability and electronic structure is a general phenomenon. It is summarized by the filled shell rule: Atoms are most stable when they have a filled valence shell. The electron shell with a principal quantum number of 1 is filled when it contains 2 electrons. In this case the filled shell rule is sometimes called the duet rule. The only atom we will be concerned with in terms of the duet rule is hydrogen. When the principal quantum number equals 2, the filled valence shell contains eight electrons, and the filled shell rule is referred to as the octet rule. We will use the octet rule as a guideline when we consider valence bond theory. Bonding rules provides a summary of the implications of the filled shell rules on bonding in organic molecules.
Hmm...Okay, now I am confused.

 

[Gokul43201]

So basically, what you've quoted says that the "filled shell rule" works (only) for n=1, 2. Beyond that, you have to use the Octet Rule (which is identical to the FS Rule for n=2). And even this is not an absolute rule, but can definitely be used to judge relative stability.

 

[Astronuc]

Bonding rules provides a summary of the implications of the filled shell rules on bonding in organic molecules.

Caution - organic chemistry (based on carbon and group 2 elements) is different from inorganic chemistry, and particularly transition metal chemistry.

Carbon (Z=6) has a relatively simple structure compared to the transition metals.

One point about a full P-subshell - all the electrons are paired - which makes for stability. This is consistent with the observation that each of the noble gases has the highest ionization potential of its respective period.

Another observation - the elements of the second group of the periodic table never form compounds in which they have more than 8 electrons in their valence shell. But elements in the third group are not so restricted.

Regarding the octet rule - take a look at http://wine1.sb.fsu.edu/chm1045/notes/Bonding/Except/Bond08.htm

Regarding:

Example: Argon. 1S², 2S²,2P²,3S²,3P².
The D orbital is not filled, yet argon is stable!

There is no 2d orbital (quantum mechanics, and just like there is no 1p orbital) and the 3d orbitals start filling with the fourth period elements, not in the third period.

Also, I think Gokul43201 made an important point that must be kept in mind - stability is relative.

 

[Dual Op Amp]

Well, it's almost been a year. For two, maybe three years I searched for the COMPLETE answer to covelant bonding. Only to find it was in hybridization. I thought this was some stupid theory, and ignored it for the longest time. I've known the answer for about two to three months now, but am now remembering this forum.
I FINALLY found the answer here...
www.chemguide.co.uk/atoms/bonding/covalent.html#top