23

These species usually do not exist in nature, but they can be synthesized. Silver has been reduced in liquid ammonia to give $\ce{Ag-}$. A lot of anionic metal carbonyl complexes $\ce{M(CO)_{n}^{m-}}$ have been synthesized: -1 $\ce{[V(CO)6]-}$, $\ce{[Nb(CO)6]-}$, $\ce{[Ta(CO)6]-}$, $\ce{[Mn(CO)5]-}$, $\ce{[Ir(CO)4]-}$, $\ce{[Co(CO)4]-}$, $\ce{[Rh(CO)4]-}$...


7

When talking about resonance it is important to know that this is only a concept to visualise certain bonding features. It is very important to understand, that the actual bonding situation is a mixture of all possible resonance structures. The Molecule itself exists at all times as this mixture, hence neither of these resonance representations can be ...


7

A strategy different from the one described by Yusuf Hasan would be to look at the neutral isoelectronic molecule. The oxygen atom is isoelectronic to the fluorine cation, so $\ce{OF+}$ is isoelectronic to $\ce{O2}$. It is difficult to write a nice Lewis structure for $\ce{O2}$, see https://chemistry.stackexchange.com/a/15061. Experiments show that dioxygen ...


6

This is a misprint. Here is a similar problem (OpenStax Chemistry, retrieved from https://opentextbc.ca/chemistry/chapter/7-4-formal-charges-and-resonance/) that makes sense and has a correct answer: As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular ...


6

A good way to think about stuff like this, is to imagine a possible synthesis for the given intermediate. One way of making this species would be to cleave the H-OF bond of HOF against it's electronegativity, that is, heterolytically cleaving the bond to make H- and [OF]+ As we are going against electronegativity, you get the idea that the intermediates are ...


5

First, keep in mind that Lewis-structures and formal charges are huge simplifications. In reality, at the quantum level, charges and bonds are spread out and it's somewhat unrealistic to say "Atom X has a charge of Q" or "Atoms X and Y have an Nth order bond between them". There's merely a cloud of electron density, and it's somewhat arbitrary to say which ...


5

In fact the unpaired electron belongs neither to Cl nor to O, but to a molecular orbital spread over both atoms [but not equally: see below]. The Lewis structure is simply not a complete representation of what is going on. (And for what it's worth wikipedia gives the same Lewis structure you drew: http://en.wikipedia.org/wiki/Chlorine_monoxide) Edit: ...


5

A good description appears in the book, "Chemistry A Molecular Approach". In the molecule of hydrogen fluoride, we know that it has a dipole moment, and fluoride is slightly negative. Ignoring this information (electronegativity difference) and we share the bonding electrons equally, the formal charge can be calculated, and both of them are determined to ...


5

Yes, ozone has a dipole moment because it is bent and has resonance structures that involve formal charge separation. Resonance can affect the polarity and dipole moment of a molecule, but only if at least one of the resonance structures involves formal charge separation, and even then there are cases where the overall molecule will not have a dipole moment....


5

Partial charges are like resonance structures: they are a convenient notation which captures the most important aspects of a phenomenon, but they are not what is "really" going on. When we say that a molecule contains partial charges, what we mean is that the electric field surrounding a bond is polarized as if some fractional number of elementary charges ...


5

Both are. Any isolated object has integral charge, as any elementary particle has integral charge. However, as we know, elementary particles are not truly located in one point. When we talk about stationary states in physics, we can assume, that particles inhabit some volume with special density function indicating how much of the particle in question is ...


5

tldr - atomic partial charges can be non-intuitive Formal charges in a Lewis structure don't always match up with partial charge assignment methods. Typically, one wants 'accurate' partial charge models, for some definition of 'accurate' - but what that means can vary considerably. I explained some of this in a related question. While I haven't worked ...


4

It doesn't have to do as much with the combing as it does with surface tension. Water on the hair clumps them together, because smaller clumps have more surface area (net surface area) and thus more surface energy. Breaking these large clumps, therefore, requires energy input, which gives us the impression that they "stick" together. As for sticking to the ...


4

The negative charge is due to the unpaired electron in the oxygen atom, remained after all bonds(molecular orbitals) are filled. Oxygen has 2 unpaired electron in its valence shell and is singly bonded to carbon in this molecule. Hence being un-bonded, the other electron exhibits the charge. Check this link, hover the cursor on the $\ce{O}$-atom, ...


4

I would like to know how this has a positive charge given that it has 3 bonding pairs of electrons around it, which would give it 6 valence electrons and 8 electrons overall. If carbon has 6 protons, how does this give an ion with a positive charge? You need to count ALL the protons and electrons in the ion. What about the atoms other than the central ...


4

First off, in neither triiodide nor periodate have the structures you drew. They all assume octet expansion on iodine which has been disproven. It was traditionally explained by the ‘participation’ of iodine’s empty 5d orbitals but these are energetically too far removed and do not take part in bonding to any noticable extent. Second off, it is wrong to ...


4

Your proposed structure is wrong. Nitrogen does not exceed the octet in any of its known compounds (and even if $\ce{NF5}$ will be discovered it will not exceed the octet according to everything we know now). However, if you have a formal negative charge that means an additional electron added to the 5 nitrogen usually has; if four of those six electrons are ...


4

TL;DR: $\ce{He^2+}$ is the only preferred notation. Notations $\ce{He^{++}},$ $\ce{He^{+2}}$ or $\ce{He^{1+}}$ are obsolete and should be avoided. From IUPAC “Green Book” [1, p. 49], section 2.10.1 Other symbols and conventions in chemistry, subsection (i) The symbols for the chemical elements: The ionic charge number is denoted by a right superscript, ...


3

The ion is not one S atom and 4 O atoms together. The ion on the whole is having two negative charges, Thus it is somewhat like 1Sulphur + 4 Oxygen+2 electrons. It is not as though there was one sulphur and four neutral oxygen atoms, and they became negatively charged during bond formation. The bond must have formed in some reaction, and in that reaction the ...


3

Well I get what your problem is. The thing is that earlier C had 4 electrons in its valence shell. Now it is bonded to 3 other atoms giving it a total of 6 electrons in its shell. The octet rule says that it should have 8. Since the other carbon took away 2 electrons this means that our carbon should have a +2 charge. Well this is where the confusion ...


3

In simple terms electrons repel each other, so distributing them over a geater number of bonds will reduce the repulsive force between them so resulting in a lower energy state. In the case of phenol the oxygen p orbital lone pair at right angles to the ring can overlap with the delocalised p orbital above and below the ring enabling a degree of ...


3

Simply put: there are more electrons than protons. Protons: Hydrogen x 3 = 1 x 3 = 3 Oxygen x 1 = 8 x 1 = 8 Carbon x 1 = 6 x 1 = 6 3 + 8 + 6 = 17 protons. Electrons: Looking at the electron shells and covalent bonds in the Lewis diagram provided by blackSmith, you have the following: 4 single covalent bonds = 4 x 2 = 8 3 lone pairs on oxygen = 3 x 2 = ...


3

Yes, the formal charge at such atoms would be different. Such atoms, typically found at the at the edge of some crystalline or regular solid, are known as "defects" in the solid. Because they have an incomplete outer shell and lack a stable bonding configuration (bonding vacancies), they usually exist as a radical, anion or cation. They define an area of ...


3

This is best explained in an example. Let’s take methyl isocyanide: $$\ce{\overset{-}{C}#\overset{+}{N}-CH3}$$ This structural formula already includes the formal charges. How do we arrive at these? First, we draw a Lewis structure with the correct, required connectivity. In this case, we have to know that a $\ce{CN}$ fragment is connected to the methyl ...


3

We describe an object, which does not exist, but is very real, called and 'orbital'. It is a region of space where an electron is 'likely' to be. It is the result of charge field interactions between positively charged nuclei. Normally we only discuss orbitals in the context of 1-2 Angstroms distance between atomic nuclei. Some orbitals have a character ...


3

Net charge is the sum of all formal charges of the atoms in a molecule. Net charge is the charge of the molecule. Formal charge is the charge of an atom in a molecule. Formal charge varies when you look at resonance structure. See this post of the nitrate resonance structures. It can be obtained through: \begin{equation} Formal\ charge\ =\ Valence\ ...


3

The hydrogen atoms are not shown, so this is really $\ce{CH3CH2CH2CH2-}$ where the rightmost carbon atom has a lone pair and a formal negative charge. The species is the n-butyl anion:


3

Ah, the confusion that layman’s terms can cause. You are assuming sharing to mean ‘both have the same electrons as before, they just share a little bit of them’ — but that’s not really the nature of a chemical bond let alone the idea behind formal charges. Rather, once a molecule forms the electrons are delocalised across the entire molecule and the Lewis ...


3

The oxidation state is a form to distribute the charge of atoms in molecules considering to have only ionic bonds. In contrast, the formal charge is the distribution of the electrons considering to have only 100% electron-pair covalent bonds. Actually, the formal charge is not very realistic and is more or less a bookkeeping tool. Now there exists ...


3

As Jan hinted in the comments, there are multiple electron bookkeeping methods which serve different uses. We want to keep track of electrons because it gives us an idea of what type of chemical behavior might be attributable to an atom or molecule. Octet Rule The Octet Rule (and the related 18 electron rule for some transition metal compounds) serves to ...


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