What is the best way to draw Lewis structures for transition metals?

[devanlevin]

if i am told to draw a lewis diagram for
PCl[tex]^{+}_{4}[/tex]
does this mean a melocule of 1P atom and 4 Cl atoms, the whole molecule with a charge of +1 or is it 1 P atom and 4 Cl[tex]^{+}[/tex] ions, giving the molecule a charge of +4

what does the lewis structure look like, is it P in the middle with Cl on each of its 4 sides, all in square brackets with a + sign after the brackets.

does anyone know of a site that has a tool where i can plug in the molecule and get the structure?

 

[Borek]

1P atom and 4 Cl atoms, the whole molecule with a charge of +1

That's it.

what does the lewis structure look like, is it P in the middle with Cl on each of its 4 sides, all in square brackets with a + sign after the brackets.

Sounds OK.

 

[devanlevin]

what are the lewis structure for these ions?

F[tex]^{-}[/tex]
is it [F][tex]^{-}[/tex]

Fe[tex]^{3+}[/tex]
is it [Fe][tex]^{3+}[/tex] with 5 dots around the Fe, (5 valence electrons) Fe[tex]^{3+}[/tex] : [Ar]4s[tex]^{2}[/tex]3d[tex]^{3}[/tex]

Co[tex]^{3+}[/tex]
is it [Co][tex]^{3+}[/tex] with 6 dots around Co

generally how are the transition metals displayed in lewis structure since they arent in any of the 7 colums of valance electrons

 

[Joskoplas]

PCl₄+:

Write P in the center. Draw a single bond to each of four Cl atoms.

P normally has five electrons. With only four bonds, it has a formal charge of +1. Good thing that fifth electron took a hike, huh?

Each of the four chlorines is stuck at the end of a single bond, with three lone pairs. The formal charge is zero for each.

Therefore, the entire molecule has a charge of +1.

F-:

Draw F in the center. Surround it with 4 lone pairs. The formal charge is now -1. Done.

Iron, Cobalt:

Applying the rules for lewis dot structures to transition metals really isn't the best idea. Even atoms in the non-metals get sort of snarky as you go further down. Nitrogen has to stop at four bonds, but phosphorous can have five (using not just its s and p orbitals, but its brand new d-orbitals as well). Carbon is stable in methane with CH₄, but Silicon in silane (SiH₄) burns immediately on contact with air -- the extra d orbitals again. Sort of like extra parking spaces for reaction intermediates.

And the transition metals? Worst of all. They mostly ignore the octet rule (though some do kind of like the 18 electron rule, which you may hear about later). So trying to make them follow Lewis structure rules is kind of like tying a knot in water -- it's too easy to break the rules for the rules to make much sense.

 

[chemisttree]

what are the lewis structure for these ions?

F[tex]^{-}[/tex]
is it [F][tex]^{-}[/tex]

Fe[tex]^{3+}[/tex]
is it [Fe][tex]^{3+}[/tex] with 5 dots around the Fe, (5 valence electrons) Fe[tex]^{3+}[/tex] : [Ar]4s[tex]^{2}[/tex]3d[tex]^{3}[/tex]

Co[tex]^{3+}[/tex]
is it [Co][tex]^{3+}[/tex] with 6 dots around Co

generally how are the transition metals displayed in lewis structure since they arent in any of the 7 colums of valance electrons

Why are you using brackets in a Lewis dot structure? Don't!

In your examples, first you determine the oxidation state of the element(s). If the oxidation state is positive, you subtract electrons (or dots, 'x's', ^o's, etc...) from the neutral element. For example, NCl₃ would have an oxidation number of +3 for nitrogen and -1 for each chlorine. The lewis dot structures for those would be an 'N' with two dots (5-3=2) and three 'Cl' with eight 'o's'(7+1=8). Together it would be a central 'N' surrounded by with two 'dots' and 6 'o's'; three sides would have a chlorine with 8 'o's', two of those being shared with the nitrogen.

Transition metals follow the 18 electron rule, not the 8 electron rule.