# Tag Info

33

According to Pauling's famous The Nature of the Chemical Bond , 3rd edition, at page 73: In the hydrogen molecule a quantum-mechanical treatment has shown that the two ionic structures $\ce{H+H-}$ and $\ce{H- H+}$ enter into resonance with the extreme covalent structure $\ce{H-H}$ only to a small extent, each ionic structure contributing only about 2 ...

27

First of all, let me state the obvious: Phosphorus is awesome. After we got that out of the way we can focus on why. There are many different modifications of phosphorus in nature. With increasing thermodynamic stability they are $$\ce{P_{white} -> P_{red} -> P_{violet} -> P_{black}}.$$ Apart from this there are many low molecular different ...

22

$\ce{C^{4+}}$ ions: Single ions are really only observed in the gas phase. There is absolutely nothing that prevents a C4+ ion from being generated in the gas phase, given sufficient energy. Each successive electron removal requires additional energy though, so by the time you get to 4 electrons removed, you're looking at quite a lot of energy required. ...

22

The typical Si–Si single bond length in a silane is around 2.33 Å. This is much longer than a typical C–C single bond (~1.53 Å) and helps explain why silicon-silicon single bonds are so much weaker than carbon-carbon single bonds (bond dissociation energy: ~53 kcal/mol for silicon-silicon vs. ~83 kcal/mol for carbon-carbon single bonds). When we try to ...

21

Chemical structures are a tradeoff of several factors, including the conditions on how they were formed. The stability of any given chemical structure depends on the ease with which any specific reaction can turn it into something else. Both graphite and diamond are very stable structures which basically means they are hard to easily convert into something ...

19

The real issue is that no one has ever taken a picture (i.e. electron density) of genuine, unambigious, cases of a single, double, triple, quadruple??? bonds. And they never will, because these concepts are not based on quantum mechanics. Two atoms reside next to each other, and if they have a favorable electrostatic interaction, then a certain type of ...

19

Quartz and diamond are stronger substances because their molecules form network covalent structures. These structures form a lattice-like structure, much the same as ionic compounds. This molecular network is also the reason that diamond and quartz form a crystalline structures, just like you'd see in ionic substances such as NaCl. Some other structures you ...

19

tl;dr The ambiguity is due to an unfortunate incompleteness of the previously (prior to 2016) existing rules. The oxidation state of fluorine in $\ce{FNO3}$ is $-1$ according to the present rules. As a general reminder, it is important to understand, that oxidation states are a bookkeeping tool only, and that they hardly represent the general bonding in ...

18

Unfortunately, the arguments presented by buckminst and Uncle Al aren't completely right. The MO schemes are correct but the HOMO-$\sigma$ orbital ($s_{\sigma}^{*}(5\sigma)$ in buckminst's diagramm, $\sigma_{3}$ in Uncle Al's diagram) is not antibonding but slightly bonding in character because there is some mixing with the $\ce{p}$ atomic orbitals of the ...

18

The order of bonding, and so the valence state of Cl in $\ce{ClO_{x}-}, x>1$ compounds is very debatable. Generally, two models exist. $\ce{Cl}$ atom, just like $\ce{S, P}$ and some others has unoccupied $d$-orbitals in the valence shell. It is possible to move some electrons from $p$-orbitals to $d$-orbitals, producing half-occupied orbitals that can ...

18

Bonds can be completely covalent as in $\ce{Cl-Cl}$, $\ce{H-H}$, etc. In these cases the electron density is shared equally between the two atoms. Bonds can also be ionic as in $\ce{Na^{+} Cl^{-}}$ where much (~80%) of an electron has been transferred from the sodium atom to the chlorine atom. Between these two extreme cases exist a continuum of bonds that ...

18

$\ce{XeF8}$ is not known to exist though O.N is +8. Why is this so? At least 2 compounds have been reported that contain the $\ce{XeF8^{2-}}$ unit. See, for example: $\ce{(NO^+)2[XeF8]^{2-}}$ (reference) Metal salts of the form $\ce{(M^{+})_2[XeF8]^{2-}}$ where M is a metal salt such as $\ce{Cs, Rb}$ (see the above reference) or $\ce{Na}$ (see p. 62 in ...

18

What you call iron sulphide, in my opinion, is more appropriately referred to as or iron disulphide. If one were to assign oxidation states to each atom, an appropriate description would be $\ce{Fe^2+}$ and $\ce{S_2^2−}$. This formalism recognizes that the sulfur atoms in pyrite occur in pairs with clear $\ce{S–S}$ bonds. These disulphide units can be ...

17

Yes, they do exist and were characterised spectroscopically, see here (and there is a note on similar clusters for sodium): Blanc, J.; Bonačić‐Koutecký, V.; Broyer, M.; Chevaleyre, J.; Dugourd, P.; Koutecký, J.; Scheuch, C.; Wolf, J. P.; Wöste, L. Evolution of the electronic structure of lithium clusters between four and eight atoms. J. Chem. Phys. 1992, 96 ...

16

There is no sharp line between ionic, metallic or covalent bonds. Most of transition metal oxides and sulfates have strong covalent characteristics. For example, the whole reason of using ligand field theory instead of crystal field theory is that the ionic description (CFT) breaks down, and the covalent effects are important in all but the simplest cases. ...

15

Whether sulfur or phosphorous actually expand their octet is contested within the chemistry community. Another term for this octet expansion is "hypervalency." You can find many works of research regarding hypervalency. The consensus, according to Wikipedia, is that both can expand their octets, but not to a significant extent. In other words, d-orbital ...

15

There are examples of metal compounds that are regarded by a majority of chemists as covalent in organometallic chemistry. The problem is that not everyone regards some of these compounds as covalent, because the border between ionic and covalent isn't so strict and it doesn't matter. Alkyllithium compounds and Grignard reagents ($\ce{R-Li}, \ce{RMgX}$) are ...

14

Boron pentachloride is likely not stable except perhaps in extreme conditions, such as under very high pressures. Even then it may be possible that a description such as $\ce{[BCl4^{-}]Cl^+}$ containing a tetrahedral boron anion could turn out to be more accurate than any hypercoordinate structure (a boron atom surrounded by more than four ligand atoms). ...

14

Yes, $\ce{KHF2}$ is an ionic compound and a covalent compound. There is nothing exceptional about it; in fact, most compounds that we call ionic (not $\ce{NaCl}$, though) have covalent bonds in them. Consider $\ce{MgSO4}$ with ionic bonds between $\ce{Mg^2+}$ and $\ce{SO4^2-}$, and covalent $\ce{S-O}$ bonds within $\ce{SO4^2-}$. Consider $\ce{Na2CO3}$ with ...

13

The reason that we discovered more fluorides of xeon than helium, argon, and krypton is quite obvious and you might already know it. As lighter noble gases have more stable shell configuration, it is more difficult to make any kind of compound, including fluorides. Also, the higher amount of fluorine atoms in a molecule of the compounds means drawing more ...

13

Generally, you wouldn't describe the bonds in an ammonium ion as dative bonds. They are usually considered to be proper covalent bonds. The term "dative bond" is typically used for bonds in (transition) metal complexes between ligands and a metal center. But for those complexes Lewis structures have lots of weaknesses and are rarely a good ...

13

Recall the fact that the basic difference between an ionic bond and a covalent bond, is basically just charge separation. In ionic bonds, charges are well separated, and the bond arises due to opposite charge attractions. Covalent bonds, on the other hand, have the electrons shared in between them, and it is these electrons that hold the nuclei of the two ...

13

Bonds are usually used to describe how electrons are shared. "Partially ionic bonds" are polar covalent bonds. Which is not the same as saying that a compound has different types of bonds in the whole compound. I think that your teacher gave you a bad question and a bad answer. If the question had asked whether all the bonds in sodium nitrate are ionic, then ...

12

The answer relates to the strength of the interactions between the component units that make up a crystal or a solid. The reason why anything is a solid at a given temperature is, crudely, that the interactions between the units that make up the solid (atoms, ions or molecules) are stronger than the amount of thermal energy available at that temperature. ...

11

You seem to have fallen into the trap of thinking that ionic and covalent bonds are fundamentally different. They are not - they are just two ends of a spectrum, which has an arbitrary division somewhere in the middle into an ionic and covalent regime. This is explained in the answers to this question. In the case of pyrite we have a relatively hard cation, ...

11

Just to add to ron's answer with some information on the properties of the silicon analogues. Silicon "alkynes" These exotic compounds were first reported on in 2004 in this paper. Typically show bond order somewhat less than 3; the HOMO has some non-bonding/lone pair character which is also manifest in a small deviation from linearity. The silicon-...

11

There is a very relevant article by Gillespie Covalent and Ionic Molecules: Why Are BeF2 and AlF3 High Melting Point Solids whereas BF3 and SiF4 Are Gases? J. Chem. Educ., 1998, 75 (7), p 923. According to the article, the charge on Be is +1.81 and the charges on the Fs are -0.91. (citing to his earlier article Reinterpretation of the Lengths of Bonds to ...

11

There is no point in doubting the composition of these compounds, the measurements don't lie. ;-) The point is that S or C (or e.g. P) can not only form bonds to the metal, but also between themselves: $\ce{FeS2}$ for example does not have $\ce{S^2-}$ anions, but $\ce{S2^2-}$, with a covalent bond in it. Add one $\ce{Fe^2+}$ and everything is fine. Sulphur ...

11

Your conception of $\pi$-bonds is a bit too restrictive. A $\pi$-bond is a bond in which there is one node along the internuclear axis. So, yes, two side-on $p$-orbitals do form a $\pi$ bond because there is one node along the internuclear axis (the one separating the bottom lobe from the top lobe). A $\sigma$-bond is a bond with no nodes (i.e. cylindrically ...

10

This is essentially the same question as to why carbon concatenates through up to four bonds with binding energies low enough to be broken again yet stable enough to last given typical terrestrial environmental conditions. Great for entities like us. On the other hand, the binding energies of Silicon, Germanium, from the same group IV, for instance, are ...

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