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Recently, we covered metallic bonding in chemistry, and frankly, I understood little.

I understand that:

  • Metals bond to each other via metallic bonding
  • Electricity can flow via free or delocalized electrons

But, I do not understand why the metal atoms turn into ions and delocalize the electrons, why don't the metal atoms stay as atoms?

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When electricity flows, the electrons are considered "free" only because there are more electrons than there should be, and because the transition metals, such as iron, copper, lead, zinc, aluminum, gold etc. are willing to transiently accept and give up electrons from the d-orbitals of their valence shell.

Transition metals are defined in part by their stability in a wide range of "oxidation states"; that is, in several combinations of having too many or too few electrons compared to protons. This is thought to be because of the d orbital in their valence shells. Compared to the s and p orbitals at a particular energy level, electrons in the d shell are in a relatively high energy state, and by that token they have a relatively "loose" connection with their parent atom; it doesn't take much additional energy for these electrons to be ejected from one atom and go zooming through the material, usually to be captured by another atom in the material (though it is possible for the electron to leave the wire entirely). This impetus can be caused by many things, from mechanical impact to chemical reactions to electromagnetic radiation (aka light, though not all of it visible); antennas work to capture radio frequencies, because the light at those frequencies induces an electric current in the wire of the antenna. Now, in the absence of a continuous force keeping the electron in this higher energy state, the electron (and the metal atoms) will naturally settle into a state of equilibrium. Electricity is generated when just such a force is acting on the metal, giving energy to the electrons in the d orbital and forcing them to move in a certain direction.

This impetus can come from many sources, as discussed, be it the movement of a magnet within a coil of wire, or a chemical redox reaction in a battery creating a relative imbalance of electrons at each of two electrodes. The end result is that the electrons, given additional energy from this voltage source, are ejected from their "parent" atom and are captured by another. The "holes" left behind by these electrons are filled by other electrons coming in behind them from further back in the circuit. Thus, the energy provided by the voltage source is carried along the wire by the transfer of electrons.

The analogy typically made is to the flow of water, and it generally holds in many circumstances; the "voltage source" can be thought of as being like a pump or a reservoir, from which water flows through pipes, and the amount of water and the pressure it's placed under (by the pump or by gravity) can be harnessed to do work, before draining back to a lower reservoir. The pipes are similar to wires in many ways; the larger the diameter, and the smoother the inside of the pipe, the more and the faster water can flow through it (equivalent in many ways to the thickness and conductivity of the metal wire), and when under enough pressure (high enough voltage), the pipes will actually expand slightly and hold more water than they would at low pressure (this is a property of wires and other electrical conductors called "capacitance"; the ability to store a charge while under voltage and to discharge it after the voltage is released).

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  • $\begingroup$ If it loses an electron, "usually to be captured by another atom in the material (though it is possible for the electron to leave the wire entirely)," where does it go? Has it been "captured" by some other element we just don't know which one at that time? $\endgroup$
    – johnny
    Sep 23 '16 at 20:32
  • $\begingroup$ Hard to say; it's difficult but not impossible for the electron to leave the Earth entirely and go zooming out into space. That would be just fine; the Sun bathes the Earth in bajillions of charged particles every second. Much more likely, our ejected electron will be captured by other materials within a rough line of sight of the atom from which it was ejected. That will affect the relative electron balance of that material alongside everything else, creating a static charge, but sooner or later the charges will equalize and the excess energy is released as a photon, likely heat. $\endgroup$
    – KeithS
    Jan 25 '18 at 14:32
  • $\begingroup$ KeithS's explanation works well with transition elements. But it does not explain why non-transition metals like aluminum or magnesium are good conductors. $\endgroup$
    – Maurice
    Dec 16 '21 at 21:10
  • $\begingroup$ This doesn't answer the question. It explains why electrons might flow but not why why metals contain "free" electrons which was the question. $\endgroup$
    – matt_black
    Dec 17 '21 at 19:50
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Molecular orbital theory gives a good explanation of why metals have free electrons

The best way to explain why metals have "free" electrons requires a trek into the theory of how chemical bonds form. Molecular orbital theory, or, at least, a simple view of it (a full explanation requires some fairly heavy quantum stuff that won't add much to the basic picture) can explain the basic picture and also provide insight into why semiconductors behave the way they do and why insulators, well, insulate.

The first step in getting to a useful intuition involves picturing how small molecules form and how their bonds work. If you start from isolated atoms, the electrons form 'orbitals' of different shapes (this is basic quantum mechanics of electrons). There are plenty of pictures available describing what these look like. But, when atoms come together to form molecules, the simple view of what the clouds of electrons look like gets a lot more complex. If the two atoms form a molecule, they do so because the energy levels of the orbitals in the molecule are lower than those in the isolated atoms for some of the electrons. This is what causes chemical bonding.

The important insight from this picture of bonding is that molecular orbitals don't look like atomic orbitals. When a bond forms, some of the orbitals will fill up with electrons from the isolated atoms depending on the relative energy levels. There may also be other orbitals (some might, were there enough electrons to fill them, form anti-bonding orbitals, weakening the strength of the bond). In some molecules those orbitals might cover a number of atoms (archetypally, in benzene there is a bonding orbital that is shared by all the atoms in the six-membered ring occupied by two electrons and making benzene more stable than the hypothetical hexatriene with three isolated double bonds).

In some solids the picture gets a lot more complicated. In graphite, for example, the bonding orbitals are like benzene but might cover trillions of fused hexagons. What makes the solid hold together is those bonding orbitals but they may cover a very large number of atoms. In metals it is similar.

It is also worth noting that in small molecules you can often get a good idea of the shape of the discrete molecular orbitals, each containing two electrons, when you start dealing with large networks of atoms joined together, the simple, discrete, picture of individual two-electron orbitals becomes pretty useless as there are too many similar ones to make reasonable distinctions.

So solid state chemists and physicists start thinking of the picture as consisting of "bands" of orbitals (or of the energy levels of the orbitals). And this is where we can understand the reason why metals have "free" electrons. When metal atoms come together in a solid, the bonds between the atoms form lower energy orbitals than the isolated atoms. But the orbitals corresponding to the bonds merge into a band of close energies. And those orbitals might not be full of electrons. In metals these orbitals, in effect, form a bond that encompasses the whole crystal of the metal and the electrons can move around with very low barriers to movement because there is plenty of free space in the band.

In insulators, the orbitals bands making up the bonds are completely full and the next set of fillable orbitals are sufficiently higher in energy that electrons are not easily excited into them, so they can't flow around. In semiconductors the same happens, but the next set of orbital bands is close enough to the bands filled with electrons that thermal energy is enough to excite some of them into a fairly empty orbital where they can move around.

Wikipedia give a good picture of the energy levels in different types of solid: source Nanite, CC0, via Wikimedia Commons. The picture shows both the spread of energy levels in the orbital bands and how many electrons there are versus the available levels. In reality there is a continuum of band widths and gaps between insulators and metals depending on how the energy levels of all the bonding orbitals work out in a particular solid and how many electrons there are to fill them up.

In short, metals appear to have free electrons because the band of bonding orbitals formed when metals atoms come together is wide in energy and not full, making it easy for electrons to move around (in contrast to the band in insulators which is full and far away in energy to other orbitals where the electrons would be free to move).

This is, obviously, a very simple version of reality. But it links the easier theory or chemical bonding and molecular orbitals to the situation in network solids from insulators to metals.

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Re: Why the metal atoms turn into ions and delocalize the electrons, why don't the metal atoms stay as atoms?

This is sort of asking why is water wet?

By definition if the atoms in an elemental sample have delocalized electrons (so that the sample will conduct electricity) then the element is a metal. If there are no delocalized electrons, then the sample won't conduct electricity and the element is a nonmetal.

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  • $\begingroup$ I'm more asking why Salt doesn't give up its electrons but steel does. (I know Salt is an Ionic compound and behaves differently to a metal, it was just an example, but the point still stands) $\endgroup$
    – Deep
    Oct 22 '15 at 20:25
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The valence electrons in the outermost orbit of an atom, get excited on availability of energy. They overcome the binding force to become free and move anywhere within the boundaries of the solid. Thus they contribute to conduction

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Though a bit different from what is asked, few things are worth noting:

  • Electrons barely move in metal wires carrying electricity.

  • Electrons do not carry energy, the electric and magnetic fields around it (outside the wire) carry and transfers energy.

A great video to explain it: https://www.youtube.com/watch?v=bHIhgxav9LY

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  • $\begingroup$ Rather, the electron net velocity during flowing electrical current is very slow. Their random momentary thermal velocity, causing resistor thermal noise, is not so small. $\endgroup$
    – Poutnik
    Dec 19 '21 at 16:43
  • $\begingroup$ If you want to comment rather than answering, I recommend you use a comment. I agree that the video is great. $\endgroup$ Dec 19 '21 at 16:58

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