Metallic bonds result from the electrostatic attraction between metal cations and delocalized electrons. The nature of metallic bonding accounts for many of the physical properties of metals, such as conductivity and malleability. Created by Sal Khan.
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- Why is a metallic bond of the same atom different from a non-polar covalent bond? If iron is sharing electrons with another iron atom, they both have the same charge. How is that different from something like two oxygen atoms?(17 votes)
- In a metallic bond, each metal atom is surrounded by lots of other metal atoms, and they all share their valence electrons. When two oxygen atoms bond, they become a molecule and don’t interact much with other molecules. A metallic bond behaves more like one big molecule (except that unlike diamond or graphite, it’s malleable because there aren’t technically any covalent bonds forming a specific crystal structure.)
Let me know if you need more help!(25 votes)
- Are there any tricks to identify if a bond is covalent, ionic, or metallic? And can someone explain metallic bonds to me? I still am having trouble after watching the video.(5 votes)
- Ionic Bonds - A bond between metal and nonmetal elements. Involves transferring electrons.
Covalent Bonds - Also known as molecular bonds. A bond between two nonmetals. Involves sharing electrons.
Metallic Bonds - A bond exclusively between metals. It creates a bulk of metal atoms, all "clumped" together. An example of this is a copper wire or an aluminum sheet.
Hope this helps.(33 votes)
- Can metallic bonds only happen between atoms of the same metal? Or can two different metals have a metallic bond?(13 votes)
- Two different metals can technically have metallic bonds, but that would become an alloy, which is not a compound but a mixture, since you can change its composition. Whereas the metallic bonds between two atoms of the same metal form a compound.(7 votes)
- How do we name metallic bonds?(3 votes)
- I'm going to assume you mean chemicals which engage in metallic bonding. In which case if it's a pure lump of metal with atoms form the same element, then we just call it by it's element name. So a lump of sodium would just be called sodium metal or pure sodium. In a chemical equation you would see it written as Na(s) where the 's' means its solid. If you gave multiple metal elements engaging in metallic bonding then you have an alloy which have no straightforward systematic nomenclature. They all have their own unique name. An alloy of tin and copper is bronze while an alloy of copper and zine is brass. So alloys are more a memorization trick than a set of nomenclature rules like other chemicals. Hope that helps.(8 votes)
- Considering the structure of an atom, how are metals different from nonmetals, like why can some atoms share their valence electrons with the whole group and others cannot?(3 votes)
- The main difference between the two groups on the atomic level is that nonmetals are more electronegative than metals. Electronegativity essentially being a measure of how well atoms can hold onto electrons.
The nonmetals occupy the upper-right corner of the periodic table where electronegativity values are the highest. Electronegativity increases as you move left to right across a period because the effective nuclear charge increases. Having a higher effective clear charge means that negative electrons 'feel' more of the proton's positive charge from the nucleus. And electronegativity also increases as you move from down to up along a group because the number of electron shells decreases. Having more electron shells means that the valence electrons are farther from the nucleus and 'feel' less attraction to the protons.
This is why metals are more able to lose their electrons in ionic bonds and delocalize their electrons in metallic bonds, since they don't have as strong of a pull on them as non-metals.
Hope that helps.(7 votes)
- Why the metals loses its electrons in metallic bond?
Do the metals either lose or gain electrons in metallic bond?(6 votes)
- Why don't positive ions formed in mettalic bond repel each other?(3 votes)
- they are held together by the delocalised electrons between them. And they do repel. If they didn't, they would be much closer than they already are. But thanks to the repulsion of two cations and the strong attraction of the anions, metals are allowed to have a strong structure(6 votes)
- @1:49what did he mean by overlapping valence electrons?(2 votes)
- It means that the electrons are not localized to a single atom as with a conventional covalent bond, but rather they are delocalized over several atoms. Individual metal atoms in a metallic bond do not have sole 'ownership' of electrons, instead they are shared between several other metals atoms.
Hope that helps.(4 votes)
- I did not really get the meaning of the covalent bonds(1 vote)
- Imagine that you were holding on to a handlebar, and someone facing away from you decided to hold on to your handlebar. This is essentially the same idea as atoms "holding on to" or "sharing" electrons to one another. This is a super basic analogy, and there is so much more to covalent bonds. I recommend looking at the covalent bond videos Sal has which looks much more in depth into covalent bonds and will help you understand it better.(5 votes)
- Why are metallic bonds weaker than ionic and covalent bonds?(0 votes)
- Not sure who told you that, but metallic bonds can be just as strong as any ionic or covalent bonds.
In reality all three of these bonds exist on a spectrum. Certain covalent bonds are stronger than ionic bonds (the carbon-carbon bonding in diamond is an example of an exceptionally strong covalent bond), and certain ionic bonds are stronger than most covalent bonds. The same idea applies to metallic bonds. No one group is always stronger than the others, instead there is a spread of strengths.
Hope that helps.(7 votes)
- [Instructor] Now the last type of bond I'm going to talk about is known as the metallic bond, which I think I know a little bit about because I was the lead singer of a metallic bond in high school. I'll talk about that in future videos. But let's just take one of our metallic atoms here. Iron is a good example. Iron is maybe one of the most referred to metals. Let's say we have a bunch of iron atoms. So Fe, Fe, Fe, Fe, hope you can read that. These are all iron atoms. And if they're just atoms by themselves they're going to be neutral. But when they are mushed together, they will form a metallic bond. Makes sense because they're metals. And what's interesting about metallic bonds, I'll draw it down here, is that metals like to share their electrons with the other metals. It kinda forms this sea of electrons. So what it can look like is, each of the irons lose an electron, I'll draw it a little bit bigger. So let's say this is Fe plus, so it has a positive charge. Fe plus has a positive charge. Fe plus, these are all iron ions, you can imagine. Fe plus, and we're imagining that they have this positive charge because they've all contributed an electron to this sea of electrons. So you have an electron here which has a negative charge. And electrons are not this big, but this is just so that you can see it. Electron here that has a negative charge. And so you can imagine these positive ions are attracted to the sea of negativity, the sea of negative electrons. Another way to think about it is, is that metals, when they bond in metallic bonds, they will have overlapping valence electrons. And those valence electrons are not fixed to just one of the atoms, they can move around. And this is what gives metals many of the characteristics we associate with metals. It conducts electricity because these electrons can move around quite easily. It makes them malleable, you can bend it easily. You can imagine these iron ions in this pudding, or this sea of electrons. So you can bend it, it doesn't break. Well if you were to take a bar of a salt right over here, if you were to try to bend it, it's very rigid. It is going to break. So there we have it, the types of bonds. It's important to realize that you can view it as something of a spectrum. At one end, you have things like ionic bonds where one character swipes an electron from another character and says, "Hey, but now we're attracted to each other," and you get something like salt. Or you have covalent bonds where we outright share electrons. And then you have things in between covalent bonds and ionic bonds where the sharing is not so equal and you get polar covalent bonds. And then another form, I guess you could say, of extreme sharing is the metallic bonds where you just have this communal sea of electrons.