Unravelling the Mystery of Transition Metals

[The_ArtofScience]

Is there is a reason why the elements in that 'family' tend to be more satisfied with having their higher s orbitals half filled before their d orbitals and another thing, why is it that from Sc to Zn its size remains static (that is strange considering its Z eff is increasing)?

There's also got to be a reason why some transition metals, ie, Cu, Ag, Au happen to appear shinier than others and I suspect it has a lot to do with its electron config. Cu has 1s22s22p63s23p64s13d10. Silver has 1s2s22p63s23p64s23d104p65s14d10. What is it that makes light behave the way it does when it backscatters from these elements?

 

[Borek]

Dull appearance will ba most likely caused by the oxide layer, so I suppose you refer to chemical properties, not electron configuration.

Not that these are separate things

 

[Renge Ishyo]

Is there is a reason why the elements in that 'family' tend to be more satisfied with having their higher s orbitals half filled before their d orbitals and another thing, why is it that from Sc to Zn its size remains static (that is strange considering its Z eff is increasing)?

My hunch is that these weird effects are due to the closeness in energy of the 4s and 3d orbitals. The way these orbitals can shift as electrons are either added to them or removed from them accounts for all types of weird affects. For example, the 4s orbitals fill before the empty 3d orbitals at first, but once you move past iron and the 3d orbitals are at least half filled, it appears as if the 4s orbitals have now shifted upward in energy above the 3d orbitals, making it easier to remove electrons from the 4s orbitals from this point forward (there is some discussion of this orbital shift in Paulings "The Nature of the Chemical Bond" if you are interested; I actually read about it in an abbreviated version of this book called just "The Chemical bond," but nowadays the former is easier to find).

Btw this sort of weirdness increases even more once you go further down the periodic table, and start looking at the lanthanides for example (they all seem to have similar properties *across* the entire row which might have to do with even more intense orbital overlapping between the s,d,and f orbitals).

 

[mgb_phys]

I was wondering that aswell.
Mercury is a liquid because the outer electrons are so strongly held that there are none available to form bonds to other mercury atoms - so it is very soft. A similair reason applies to gold.
But these two are also very reflective and conductive - implying that there are are lots of free electrons available!

 

[Renge Ishyo]

Mercury is a liquid because the outer electrons are so strongly held that there are none available to form bonds to other mercury atoms - so it is very soft. A similair reason applies to gold.
But these two are also very reflective and conductive - implying that there are are lots of free electrons available!

As far as I know, Mercury is a liquid at room temperature not because it's electrons are so strong that they resist forming bonds (mercury does form dangerous oxides with the air, it is hard to see this because the product is a colorless vapor), but rather that the mercury atom itself is so large that the bonds that form between mercury atoms are relatively weak and easy to break.

I am assuming that the explanation for the appearance of mercury is analogous to the reasons why Cesium in group I is a gold, reflective liquid at room temperature; Cesium is actually the most reactive of the non-radioactive group I metals (easiest to remove its electrons; it ignites explosively if exposed to oxygen) and yet it forms a liquid at room temp when isolated because its bond strength in the elemental form is so low due to the increasing size of the cesium atom compared to the group I metals in the column above it. A general trend is that as atom size increases, the chemical bonds that are formed in the elemental form are more reactive/less stable than those of atoms that appear above it on the table, and a quick glance at the periodic table shows that mercury appears in the same period/row as cesium (although both are unique in the sense that they are so soft that they actually form liquid metals at room temp).

 

[Borek]

Trends won't explain mercury low melting point. If anything - Zn melts around 150 deg C, cadmium around 320 deg C, mercury should melt higher. It doesn't. Atoms size is not that important as well - thallium (closest neighbor to the right) is slightly larger, still solid and melts around 300 deg C. So the thing is much more subtle.

My take is that we are talking about system (atoms/electrons) that is complicated enough to show some occasional anomalies. Mercury is one of these.

 

[Renge Ishyo]

Zn melts around 150 deg C, cadmium around 320 deg C, mercury should melt higher

?? The number you have provided for Zinc disagrees with my source.

From The Elements by Elmsley, J.

Melting Point Data

Zinc - 692.73 K
Cadmium - 594.1 K
Mercury - 234.28 K

The melting point data for group I

Potassium - 336.80 K
Rubidium - 312.2 K
Cesium - 301.55 K

Although the point is taken that transition metals don't seem to follow any sort of logical connection with the behavior of the main block (anymore than the Lanthanides do). I'll do 20 pushups for forgetting that ^^

 

[Borek]

Sorry, my mistake. Error in source. Still, trend suggests melting point of mercury around 500 K, no way for it to be below 0 deg C.