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|The content of Electrovalency was merged into Ionic bonding on 6 April 2015. That page now redirects here. For the contribution history and old versions of the redirected page, please see ; for the discussion at that location, see its talk page.|
- 1 Electron affinity
- 2 Electronegativity
- 3 Caption
- 4 Reformatted
- 5 Difference with covalent bonds
- 6 Why does it happen?
- 7 Ion pairs?
- 8 DANGER!!! INACCURATE!!!
- 9 Strength of Ionic Bonds vs. Strength of Covalent Bonds
- 10 B class?
- 11 "Stock system Name" for ions
- 12 conflict of interest
- 13 List of common ions
- 14 Percent ionic versus covalent
- 15 Ionic interactions redirects here
- 16 Extra Ionic Bonding Examples
- 17 New image
- 18 Formation Section
- 19 Bond Strength
- 20 Disputed
Dear author, in the Formation section,second line......Electron affinity of an atom should be negative for it to accept electrons and form an anion. It is the energy released. ([User:Mani_g1|Dr. Manisha Jain]) —Preceding unsigned comment added by Mani g1 (talk • contribs) 13:41, 5 May 2011 (UTC)
I was taught that the difference in electronegativity between the atoms in an ionic bond is greater than 1.25, and anything less is Polar Covalent. The article mentions electronegativity but does not go into enough detail to explain what's going on. 22.214.171.124 (talk) 22:25, 18 January 2009 (UTC)
How about a nice caption to say something about the figure - what it represents, how it was made, etc.? -- Marj Tiefert, Wednesday, May 14, 2002
I've done a faily major reformat of this page. If it looks bad to anyone please reformat, or alternativly say what the problem is on this page along with their browser and screen resolution. I'm a bit worried about the equation overlapping the table on the right, which I think it might do with a small enough screen size Theresa knott
- I think the phosphorus anion is fairly unlikely to exist, the third ionisation energy is probably prohibitivly large. In addition to this, since P is a relatively large atom anyway, there will almost undoubtedly be considerable covalency in the bond. Good old NaCl, although for less exotic, is probably a better example. jwasey 16:16, 28 Dec 2004 (UTC)
Difference with covalent bonds
- I think some sort of general bonding page needs to be created which explains either end of the covalent/ionic continuum but emphasises that nothing really exists in these theoretical bonding modes. I don't know how to do this, is it possible? jwasey 16:16, 28 Dec 2004 (UTC)
- Covalent is sharing of electrons and ionic is electrostatic forces of attraction resulting from the non metal having more electrons than protons (so being - ve) and the metal having less electrons than protons (+ve) Tourskin 05:27, 5 February 2007 (UTC)
Why does it happen?
Since the Lithium and Fluorine atoms are both electrically neutral prior to bonding, what causes an electron to "leave" the Lithium atom for the Fluorine atom in the first place? --RussAbbott 02:27, 17 Nov 2004 (UTC)
- Atoms always "try" to obtain the most stable electron configuration (that of a noble gas). As the Li athom has 1 electron in its last electron shield, it gives up 1 electron and loses its last shield: the next outer shield has 8 electrons, which is a stable configuration. The F has 7 electrons in its last shield: it catches the electron the Li gave up and then has 8 electrons in its last shield: another stable configuration. Maybe this explanation should be in the main article? -- Habbit 17:13, 17 Nov 2004 (UTC)
- Is it really the case that if there were a free electron floating about in the vicinity of a Fluorine atom, the Fluorine atom would capture it, thereby becoming electrically negative in the process? And is it really the case that a sole Lithium atom expels its sole outershell electron so that it can empty its outermost shell--becoming electrically positive in the process? Neither of these seem likely. Perhaps the two together make more sense. But what is the force that causes it to happen? It isn't an electrical attraction or repulsion (since the two atoms are neutral). In fact, the electron is going in the "wrong" direction with respect to both the Lithium and Fluorine atoms. So it must be something else. What is it? --RussAbbott 23:48, 17 Nov 2004 (UTC)
- This is an extremely important point in all of chemistry. How do quantum effects create "forces" that favor complete energy levels, at the expense of having a net charge -- and thus violating our intuition (correct in the classical world) that a net charge leads to electostatic instability?
- Keep in mind first that there is no stable electric configuration in the classical world -- in fact there is no equilibrium configuration -- and one of the the first things that quantum did was to fix that, without introducing any new forces. It's fermionic exclusion plus h-bar as embodied in the Schroedinger equation.
- But that's just the warm up. To explain the capture of the 10th electron, continue to think in terms of the Schroedinger equation.
- The most important point: the 9 electrons of neutral F are smeared out, not concentrated on top of the positive charge (for h-bar and fermion reasons). The electric force felt by a hypothetical 10th electron is the sum of a standard attractive inward Coulomb toward the nucleus, plus a smeared-out (convolved) integral combination of repulsive Coulombs from the all points of the 9-electron cloud. The latter force is radially outward (at least assuming the distribution of the 9 electrons is spherically symmetric), but it is weaker than the inward attraction when near the nucleus. It cancels the inward attraction (fully shields it) when the hypothetical 10th electron is far from the nucleus. So there is a zone near the nucleus where the attraction is toward the nucleus, and the potential energy for the 10th electron is lower than at infinity. So the 10th electron would like to go there, or at least would like to stay there if it gets there. Of course, for h-bar and fermion reasons it can't just squeeze in there, it also has to be diffuse, but it turns out that it can get close enough, without moving too fast, to have a total energy less than it would have at infinity. So it likes to be there, and that completes the shell.
- Now in reality, this was a simplification. Actually, the 9-electron cloud is not spherically symmetric, but is a combination of 9 quantum states whose sum effect is not spherically symmetric, and the 10th electron fits in as a 10th quantum state together with the first 9 and nicely completes a representation of SO(3). Indeed, the 10 together now have spherically symmetrical distribution. Furthermore, the interaction of the 10th electron is not with a 9-electron fluid spread out in a density (a mean-field), but rather it interacts via a 10-fold alternating tensor product that allows for all the electrical couplings, phase cancellation (interference) effects, and the exclusion principle to be handled all at once. Finally there is second quantization, QED, relativistic effects and so forth.
- And at the same time, it is clear that you can't just keep adding electrons this way. If you add enough electrons, the favorable "zone" for the next added electron to have lower net energy than infinity gets too small to be usable, both because there is more negative charge around, but also because more orbitals are filled and the new electron can't get close enough to the nucleus to profit much from the positive charge there. But what determines when you can't continue to add electrons?
- Indeed, it turns out that the average distance from the nucleus, or the net energy -- in fact, these are probably the same thing -- will jump upward a big leap once the energy level gets exactly filled. The next electron is forced to be suddenly a considerable distance further out and it is no longer favorable. Indeed, this is exactly when the representation gets full and the positional distribution gets spherically symmetric again.
- Update: see Slater's rules, effective nuclear charge, shielding effect. They don't give the shape of the potential, but gives some guide on the behavior of the shielding. In particular, Slater's rules says that (as an empirical rule of thumb, anyway) adding more electrons to a given d or f orbital does not increase the repulsive effect at all for further electrons in the same subshell. That is, the shielding they produce doesn't kick in until you get to a higher subshell. A modified version of this happens for s and p orbitals. This says, more or less, that the effective potential due to the combination of the nucleus and the already-present electrons is significantly more positive at a given radius, which I called a zone, than you would think just from computing the total charge of the atom. (In fact, the total charge only predicts the behavior of the potential much further out, near infinity.)
- I see that I've only explained why the fluorine atom would capture the electron (because shielding will fail to be complete), not why the lithium one would let go of it. Of course, the latter takes energy. To expel the electron you have to assume it happens because of an opportunistic, random "bump" either from radiation that creates an ion, or some kind of quantum jumping over an energy hump into the waiting fluorine "zone" if the fluorine atom is close enough. In either case, it comes from the dynamism of the situation and does not happen just passively by itself.
doesn't litium have three electrons? I don't think the diagram is right
It is energetically unfavorable to transfer an electron from Lithium to Fluorine to form separate ions. However, quite a lot of energy is released when the ions pack into the lattice, and this more than compensates for the energy required to form the ions.
I am somewhat suspicious of the statement that no bond is purely ionic. In the compound sodium borohydride, there are Na+ and BH4- ions. There are no lone pairs available on BH4- to share with sodium, so how can there possibly be any covalency?
- You are forgetting, that covalent bonds aren't just confined to the "sticks between atoms" model. If you think about it using the LCAO model, the borohydride anion has a molecular orbital that is a linear combination of the atomic orbitals; this can then engage in mixing with the remaining electrons floating around the sodium (and the unoccupied shells too). The net energy contribution of this mixing is probably inconsequential next to the "ionic" contribution, however.
Well the bond starts off as a covalent bond, with the two atoms sharing lithium's single valence electron. However, their electrongativities are 1.0 for lithium and 4.0 for fluorine. The electronegativity of fluorine is obviously much stronger, so it pulls the electron closer to itself, until it gets to the point that the electron has comletely left lithium and has completed fluorine's octet.126.96.36.199 03:35, 11 December 2005 (UTC)
Ionic bonding reminds me of the mooring lines that get cast into the harbour when the big ship arrives. A sodium atom sails close to a chlorine atom. It takes an interest in it, and then it casts over a mooring electron. The two atoms then haul themselves together. Clearly the bonding has begun before the official explanation for the bond has even come into existence. See 'Gravity Reversal and Atomic Bonding' at http://www.wbabin.net/science/tombe6.pdf Yours sincerely, David Tombe (188.8.131.52 12:22, 8 December 2006 (UTC)).
Re: Li and F reaction. I think it is a mistaken notion that they should not react. The most stable state for lithium and fluorine are as lithium cations and fluoride anions. That is the most common state you would find them in nature, i.e., the same electronic configuration as their nearest noble gas. You must supply energy to convert a lithium cation to lithium metal. That energy suggests lithium is sufficiently reactive to enable it to react with water. I understand the reason to ascribe elements as 0 kJ/mol in their elemental states, but that wrongly implies they are actually at some neutral state. Petedskier (talk) 23:03, 21 May 2012 (UTC)
'ion pair' redirects to this page, but there is no mention of it in the article. From what i understand it is a case when two ions which perhaps were formally attached are still next to each other, but not necessarily bonded. Does this sound about right? Do they influence each other significantly?
- It is virtualy imposible to separate ions enough that they are not bonded to each other any longer. Atempting to do so though the input of considerable energy (in, for example, an electrochemical cell) invariably produces the uncharged elemental atoms. The only instances where ions can exist without some sort of latice or solvent is under extreemly high temperatures (on the level of the corona of the sun) or under vacume conditions where ions can be produced via radiation bombardment. To get back to your quesiton, "Ion Pairs" are always bonded.
Hang on a a day!! Not a minute! This is wrong. The first paragraph states a nonpolar covalent bond is weaker than a covalent polar bond. Okay. But you CAN'T use that to determine if an ionic bond is stronger than a covalent. Because Covalent bonding is very very strong between the actual atoms involved and weak only as an intermolecular force. Ionic bonds are between molecules as in a lattice (intermolecular) and between the actual ions. To prove this, you only need to look at Diamond, Carbon-Carbon bonds, all covalent 100%, with no degree of ionic behavior. Diamond is the hardest material in the world. So I shall remove this very damaging bit at the first paragraph, unless someone with more real scientific knowledge (like a degree or something) rules me wrong. Tourskin.
Strength of Ionic Bonds vs. Strength of Covalent Bonds
There are some problems here of accuracy.
Polarisation effects - this para is pretty good until it says AlCl3 is covalent! Its an ionic solid (6 coordinate Al) and forms 4 coordinate Al dimers (covalent) only in the melt and gas phase.
Ionic structure - What is this paragraph saying- is there a confusion here with metal structures where some metals approximate closest packing with hcp and fcc structures being common (bcc isn't closest packed)? Simple cubic what is that? In the unit cell the weight is one of the atoms? In an ionic solid the unit cell has to contain at least one formula unit i.e 2 different atoms. Keeping it really simple and only considering binary ionic compounds - there are two factors - the ratio of positive to negative and the relative sizes of the two ions involved.
Ionic v Covalent - the assertion that covalent bonding is determined by VSEPR overstates the power of VSEPR - which best predicts molecular shapes of main group compounds (inert pair effect and inner shell polarisation effects ignored).
--Axiosaurus (talk) 09:41, 13 August 2008 (UTC)
- I agree there are some problems. I suppose AlCl3 has some tendency toward covalent bonding, which allows it to readily form molecules by melting at a low temperature or dissolving in organic solvents. That statement could be improved by specifying under which conditions AlCl3 (or Al2Cl6) is really molecular, or by using a different example (isn't Al2Br6 is molecular even in the solid?). Simple cubic is the name given to the structure of CsCl. Even if it has two different atoms, the atoms of each individual element are in a simple cubic lattice. But anyway, feel free to downgrade the rating. This could be a good use for the new "C-class" rating--more than a "Start" but less than a "B". --Itub (talk) 08:50, 14 August 2008 (UTC)
- I have made some remedial changes - the AlCl3 story is interesting but I feel irrelevant to this article- e.g. how conductivity changes as melting point approaches- AlCl3 article is very specific in saying that it is covalently bonded- and that article needs attention and more detail adding.--Axiosaurus (talk) 13:46, 20 August 2008 (UTC)
"Stock system Name" for ions
I changed the heading to "Common name" from the previous heading "Stock System Name" If there ever was a formal Stock system for ions, it has now, 80 years or so later, been superseded many times, most recently by IUPAC latest 2005 recommendations. --Axiosaurus (talk) 17:09, 14 April 2009 (UTC)
conflict of interest
The author of the papers is the author of the section in the wikipedia article:
Some charge transfer (i. e., ionic character) can exist even in pure elemental crystals, if atoms occupy symmetrically inequivalent positions - but usually such ionic character is too small to be important. Significant charge transfer was recently reported in a high-pressure phase of boron, which consists of two types of clusters: B12-icosahedra and B2-pairs, arranged as ions in a NaCl-type structure, with charge transfer of ~0.5 electrons from B2 to B12.
- Oganov A.R., Chen J., Gatti C., Ma Y.-Z., Ma Y.-M., Glass C.W., Liu Z., Yu T., Kurakevych O.O., Solozhenko V.L. (2009). Ionic high-pressure form of elemental boron. Nature 453, 863-867
List of common ions
I'm tempted to delete the entire section listing common ions, as (1) it is of limited usefulness, (2) it is incomplete, (3) it cannot be completed practically, (4) there is already Category:Cations and Category:Anions, and (5) there is already a list of representative ions at Ion. Any objections?—Tetracube (talk) 16:32, 11 June 2009 (UTC)
- Alright, since nobody objects, I will remove the section.—Tetracube (talk) 17:04, 13 July 2009 (UTC)
Percent ionic versus covalent
In the page there is a segment dicussing that all ionic bonds must have some covalent nature. It also dicusses the percent of ionicity. It would be useful to also include the equation by which percent ionicity can be calculated. % ionic character of a bond = 1 - exp (1/4 * delta chi ^2) where delta chi is the difference between electronegativity of the anion and cation. Also in response to the below comment about electronegativity - linking to the page on electronegativity is a legitimate explanation for electronegativity. —Preceding unsigned comment added by 184.108.40.206 (talk) 16:01, 10 December 2009 (UTC)
I agree with the above, but a reference would be necessary. Also, adding a diagram illustrating the ranges for the three types of electrochemical bonds (ionic, polar covalent, and non-polar covalent) would be useful. Mego (talk) 04:56, 30 December 2009 (UTC)
Ionic interactions redirects here
It shouldn't. Ionic interactions refers to the weak electrical attraction which occurs between atoms that are already covalently bonded, e.g. in globular proteins. There doesn't appear to be a page for this type of interaction. :S —Preceding unsigned comment added by 220.127.116.11 (talk) 18:23, 15 March 2010 (UTC)
Extra Ionic Bonding Examples
In the summary section, the example of Sodium and Chlorine seems rather incomplete and simple as the bonding of complex ions require new sets of rules (parentheses...) . In other words, eg) MgSO4 + AlCl3 => MgCl2 + Al(SO4)3 Such "more complex" examples, in my opinion, should be added. --Jjeong12 (talk) 17:43, 21 June 2010 (UTC)
The new gif is nifty, but I'm a little concerned that it shows um... perhaps the gas-phase reaction of sodium with fluorine, rather than simply the ionic bond. --Rifleman 82 (talk) 22:50, 17 June 2011 (UTC)
- Well yes, but I think it is the right image for the audience of this article. It's unlikely that PhD research students are likely to be looking at this fairly brief page on ionic bonding to get to grips with the subject. Either way, the previous image was just too awful for the year 2011. Wdcf (talk) 00:20, 18 June 2011 (UTC)
I have a problem with the Formation section, but I'm not brave enough to make a change in it. The sentences that seem incorrect are:
"The removal of electrons from the cation is endothermic, raising the system's overall energy." and "However, the action of the anion accepting the cation's valence electrons and the subsequent attraction of the ions to each other releases energy and thus lowers the overall energy of the system."
In both cases, they are not cations or anions until they lose or gain the electron. They are just elements that become either cations or anions.
- I think the reason the author used cation and anion even though they aren't ions yet (as you pointed out) is because they could not come up with a better/clearer label for which atom they are referring to. If you can come up with better labels, than be bold and change it! Wertyu739 (talk) 16:31, 31 May 2014 (UTC)
This is not the first discussion of bond strengths. In the discussion there are comments on ionic v covalent bonds. I ended up on this article still trying to understand the differences in bond energies. What or how are bond dissociation energies for covalent bonds different than the bond strength of NaCl? What is actually measured in determining bond dissociation energy vs bond strength? Petedskier (talk) 23:19, 21 May 2012 (UTC)
Ionic bonding exists but there is not really such thing as AN ionic bond. IMHO the article should therefore be renamed "ionic bonding". The term ionic bond should be avoided completely.
The term is wrong and very misleading because of what it suggests in analogy with covalent bonding: directionality and localized behavior. (For metallic bonding the same argument holds.) In an ionic compound like NaCl there is no single bond between one Na and one Cl. Each ion has six neighbors at angles of 90 degrees. This is completely incomprehensible from a covalent VSEPR point of view. Not everything is an organic molecule! Jcwf (talk) 16:19, 6 November 2013 (UTC)
- Agree. So just move it to Ionic bonding (easy part) and rewrite the whole thing (not so easy part). Vsmith (talk) 17:34, 6 November 2013 (UTC)
- Now that the article has been moved and rather thoroughly rewritten the dispute tag is also gone.