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69

Tetrahedral complexes Let's consider, for example, a tetrahedral $\ce{Ni(II)}$ complex ($\mathrm{d^8}$), like $\ce{[NiCl4]^2-}$. According to hybridisation theory, the central nickel ion has $\mathrm{sp^3}$ hybridisation, the four $\mathrm{sp^3}$-type orbitals are filled by electrons from the chloride ligands, and the $\mathrm{3d}$ orbitals are not involved ...


60

That's a good, concise statement of Bent's rule. Of course we could have just as correctly said that p character tends to concentrate in orbitals directed at electronegative elements. We'll use this latter phrasing when we examine methyl fluoride below. But first, let's expand on the definition a bit so that it is clear to all. Bent's rule speaks to the ...


48

I'm not sure about the existence of molecules with bridges through rings. However, there are several publications of synthesis of molecules mimicking wheels and axles ([2]rotaxanes; The “[2]” refers to the number of interlocked components) as one shown below (Ref. 1): (The diagram is from Reference 1) This specific molecule (8; an “impossible” [2]rotaxane)...


45

TL;DR: The $\ce{O-O}$ and $\ce{S-S}$ bonds, such as those in $\ce{O2^2-}$ and $\ce{S2^2-}$, are derived from $\sigma$-type overlap. However, because the $\pi$ and $\pi^*$ MOs are also filled, the $\pi$-type overlap also affects the strength of the bond, although the bond order is unaffected. Bond strengths normally decrease down the group due to poorer $\...


42

Here are the $\ce{H-X-H}$ bond angles and the $\ce{H-X}$ bond lengths: \begin{array}{lcc} \text{molecule} & \text{bond angle}/^\circ & \text{bond length}/\pu{pm}\\ \hline \ce{H2O} & 104.5 & 96 \\ \ce{H2S} & 92.3 & 134 \\ \ce{H2Se}& 91.0 & 146 \\ \hline \end{array} The traditional textbook explanation would argue that the ...


40

Starting point: 2s orbitals are lower in energy than 2p orbitals. The $\ce{H-N-H}$ bond angle in ammonia is around 107 degrees. Therefore, the nitrogen atom in ammonia is roughly $\ce{sp^3}$ hybridized and the 4 orbitals emanating from nitrogen (the orbitals used for the 3 bonds to hydrogen and for the lone pair of electrons to reside in) point generally ...


40

Yes, there are coordination complexes of large elements which have coordination numbers greater than eight. Some examples are: $\ce{[ReH9]^2-}$ with a tricapped trigonal prismatic structure. The nine hydride ligands are small enough to fit around the relatively large rhenium atom fairly easily. This ion can be isolated as a potassium salt $\ce{K2ReH9}$. $\...


39

The meaning of covalent bonds being directional is that atoms bonded covalently prefer specific orientations in space relative to one another. As a result, molecules in which atoms are bonded covalently have definite shapes. The reason for this directionality is that covalent bonds are formed by sharing electrons between atoms, or, in other words, as you ...


35

Mathematical Explanation When examining the linear combination of atomic orbitals (LCAO) for the $\ce{H2+}$ molecular ion, we get two different energy levels, $E_+$ and $E_-$ depending on the coefficients of the atomic orbitals. The energies of the two different MO's are: $$\begin{align} E_+ &= E_\text{1s} + \frac{j_0}{R} - \frac{j' + k'}{1+S} \\ E_- &...


35

Yes, it has a lot to do with mass. Since deuterium has a higher mass than protium, simple Bohr theory tells us that the deuterium 1s electron will have a smaller orbital radius than the 1s electron orbiting the protium nucleus (see "Note" below for more detail on this point). The smaller orbital radius for the deuterium electron translates into a shorter (...


35

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 ...


34

All credit to Zhang et al. "Real-Space Identification of Intermolecular Bonding with Atomic Force Microscopy" Science Vol. 342 no. 6158 pp. 611-614. Yes, direct images of bonds, not only covalent bonds but also intermolecular hydrogen bonds have been recorded. It is the electron density that is being observed, covalent and hydrogen bonds involving high ...


33

Introduction The bonding situation in $\ce{(AlCl3)2}$ and $\ce{(BCl3)2}$ is nothing trivial and the reason why aluminium chloride forms dimers, while boron trichloride does not, cannot only be attributed to size. In order to understand this phenomenon we need to look at both, the monomers and the dimers, and compare them to each other. Understanding the ...


32

The difference between snow and ordinary ice cubes is mainly about the size of the particles. Snow is made from small, irregular crystals with many edges at a very small scale. Light is refracted or scattered by the edges (or the interface between air and the edges). Snow is white because the scattering effect of those edges dominates what happens to light ...


31

14 coordination is claimed in $\ce{U(BH4)4}$ (ref_1, p. 268). The molecule exists as a polymer in the solid state. Six hydrogens from two of the $\ce{BH4}$ groups bond between the boron and uranium (a bridge bond). Two hydrogens from each of the two remaining $\ce{BH4}$ groups also bridge bond to uranium; the other two hydrogens bond to an adjacent ...


30

A variation on this theme is Ice VII, in which two cubic ice structures are intertwined with hydrogen bonds from each component structure passing through the hydrogen-bonded rings formed by the other component. Known to occur naturally on Earth as a high-pressure phase trapped in diamonds, Ice VII is a stepping-stone to the macromolecular and superionic ...


29

Liquid carbon does indeed exist, but perhaps surprisingly, relatively little is known about it. It only exists above around $4000\ \mathrm{K}$ and $100\ \mathrm{atm}$, which are not trivial conditions to sustain and probe. There certainly are many theoretical studies into the properties of liquid carbon, though. You can find a phase diagram for carbon here, ...


29

Noble gases usually do not form strong bonds between their atoms - it takes a fair amount of energy to dimerise them into excimers, but those are short-lived excited molecules. Thanks to excitation, shells of the atoms aren't closed and they react, but very quickly they lose energy and become separate atoms. On the other hand there are many stable molecules ...


29

Here is a picture of a "classical" carbocation, there is an electron deficient carbon bearing a positive charge. There are many examples of "non-classical" carbocations, but the 2-norbornyl carbocation is among the best known. Labeling experiments have shown that the positive charge resides on more than one carbon in the 2-norbornyl ion. Early on, the ...


29

Because $\ce{PCl5}$ does something which is not immediately obvious from its molecular formula: it autoionizes and becomes an ionic solid $\ce{PCl4+PCl6-}$. As such, it has much stronger interactions than $\ce{PCl3}$ with its mere dipole-dipole attractions, hence the higher melting point. If not for that fact, you deduction should have worked just fine.


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 ...


26

You're right in that bond length, and therefore bond strength does affect acidity (see: $\ce{H2S}$, $\mathrm{p}K_\mathrm{a} = 7$ and $\ce{H2O}$, $\mathrm{p}K_\mathrm{a} = 15.7$). If we defined acidity with the following equation $$\ce{HX -> H + X}$$ then the bond strength would indeed be the only deciding factor in the acidity of $\ce{HX}$, since the ...


24

tl;dr The next in the series is called φ bond. There is even a tiny Wikipedia article about it. Nicolau pointed me to the Wikipedia article, that had at the time a tiny section about the φ symmetry of the bond. Ben also kindly agreed with my naming proposition. I'd like to back up just a little bit an quote one sentence of this article: The type ...


24

Consider the auto-ionization of water : $\ce{ 2H_2O->H_3O+ + OH-}$ The first oxygen has three bonds, the second only has one. You can think of the reaction taking place by a lone pair on the oxygen of one water molecule ripping off the proton only of the hydrogen of another water molecule to form a covalent bond between them using just the lone pair. ...


24

$\ce{C + O2}$ is awfully complicated, so let's just pretend you've asked this: In a single act of the reaction $\ce{H. + H .-> H2}$, how is momentum conserved? That's a legitimate concern all right. After all, we are taught that this reaction does happen instantly, once given a chance, and that's in fact true. Also, we know that it releases a lot of heat. ...


23

Unfortunately, nothing in the bonding situation in carbon monoxide is easily explained, especially not the dipole moment. According to the electronegativities of the elements, you would expect the partial positive charge to be at the carbon and a partial negative charge at oxygen. However, this is not the case, which can only be explained by molecular ...


23

Karsten's answer is excellent, but here is a figure that shows the mathematics involved: The central atom (green) is at the center of the cube, the four other atoms (purple) are at alternating vertices and the geometry should be clear. Alternatively, if you orient the molecule so one peripheral (purple) atom is directly "above" the central (green) atom, ...


22

Due to symmetry constraints ($D_\mathrm{3h}$) in $\ce{SO3}$ there 6 electrons in $\pi$ type orbitals. In a wider sense of the term this molecule is Y-aromatic, but the HOMO actually represents two electrons in in-plane lone pair orbitals of oxygen. The LUMO is an antibonding $\pi$ orbital with respect to all $\ce{S-O}$ bonds. The following valence molecular ...


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. ...


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