# Why are noble metals more electronegative then most metals?

I was researching about electronegativity when I looked up what a graph of electronegativity within the periodic table is. And, this appeared. I scanned it, matching up everything I knew about the elements up to that table. Yep, Group 1 is all light colored; yea, the nonmetals are all really dark blue culminating in fluorine; and yes, the noble gases are mostly zero (with the exception of the traitors xenon and krypton, which coincidentally are also the only electronegative noble gases!)

But while I was scanning the page from Wikipedia on noble gas compounds, I suddenly realized that there was a compound called xenon hexafluoroplatinate, meaning that xenon could bond with...platinum? Checking the table, I was astonished to find that the entire noble metal group was actually more electronegative then the metals surrounding them! Why?! Aren't they extremely non-reactive? How then can they crave electrons even more then regular metals?

• Are you asking about noble gases or metals? You can calculate electronegat. also for other noble gases. – Mithoron Jul 17 '15 at 19:10
• Metals have tendency to lose electrons so high electroneg. means it's smallest for them - perfectly reasonable. – Mithoron Jul 17 '15 at 19:14
• I'm asking why are the noble metals actually even more electronegative then the other metals surrounding them. It is kinda counterintuitive to me. – HyperLuminal Jul 17 '15 at 22:37
• It shouldn't be - actually elements with highiest and lowest electronegativity are usually highly reactive - those in the middle like gold aren't. – Mithoron Jul 23 '15 at 16:00

The noble metals are defined by a resistance to oxidation and corrosion, and this should not be interpreted as a lack of reactivity, but instead an aspect of their high EN. So, there is no contradiction as you seem to be thinking. Basically, they hold onto their electrons better than other metals, so it is harder for acids and oxygen to steal electrons from these metals.

The noble metals are generally considered to be ruthenium, rhodium, palladium, silver, osmium, iridium, platinum, and gold, so I'll mainly be addressing these particular elements. There are a few reasons why they should have a higher electronegativity:

• The lanthanide contraction causes these atoms to have a higher than expected $$Z_{\text{eff}}$$. This means they hold on to their electrons tightly, so have higher electron affinities and ionisation energies which corresponds to a high EN. This doesn't carry into group 12 because accepting electrons for these elements would result in adding another energy level, so it's not as favorable.

• The filling order: By period 6, the 6s, 4f, and 5d orbitals are all so close in energy that the filling order changes, which affects the properties and chemistry of the elements. Notice that almost all of these metals, apart from Os and Ir, break the typical filling order. Adding electrons to these atoms therefore does not have the same effects that are dictated by normal periodic trends.

• Relativistic effects - Somewhat related to filling order. By period 6, nuclei are so heavy that the core electrons are moving close to the speed of light. This causes a contraction of the s orbitals, the inert s pair effect, and other things that have a big effect on the electronic structure and chemistry.

Also, the noble gases start becoming reactive around Kr, and more so for Xe, because they are so big with such high shielding that electronegative atoms are able to take electrons to form bonds. Some people get confused about this because they think it's just Kr, and Xe being weird, but it's actually a trend and continues with Rn, but there isn't much data on Rn compounds, or much use for them, since Rn-222 is the longest lived isotope with a half-life of ~3 days.

Periodic Trends of Transition Metals could help explain some of this.

• @orthocresol Based on periodic trends, the expected order to fill would be 6s, 4f, 5d, which is observed in most period 6 elements (Cs, Ba, Pr to Ir). The "4f<5d<6s" was confusing to me based on the expected filling order, so I removed it. – Mecury-197 Oct 10 '15 at 22:26
• Eek that was a typo on my part, my bad. – orthocresol Oct 10 '15 at 22:27
• Note that lanthanide contraction is irrelevant for Ru, Rh, Pd and Ag. – Gimelist Oct 11 '15 at 3:40

Noble metals are close to filling both $s$ and $d$ subshells, so there is a certain stability in gaining electrons. Gold atoms in complexes form bonds with each other similar in strength to hydrogen bonds and can form stable $\ce{Au^-}$ salts with cations like $\ce{Cs^+}$. Platinum similarly forms $\ce{Pt^{2-}}$. There are also relativistic effects that change the characteristics of the $d$ subshell.

https://en.wikipedia.org/wiki/Aurophilicity

http://www.sciencedirect.com/science/article/pii/S129325580500230X?via%3Dihub