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Chemical Bonds and Valence Electrons · Acedra
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Chemical Bonds and Valence Electrons
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0:00
Hi, I'm Professor Melko, chemistry PhD and your university chemistry professor. Today we're going to talk about the chemical bond. I know what you're thinking. The chemical bond? That sounds complicated. Well, it's really not that bad. But I have to yell at you, I will. Lecture is about to start pay attention and we're off. The chemical bond is all about valence electrons. It turns out when two atoms get close to one another, their valence electrons overlap and form a chemical bond.
0:33
We need to know the difference between core and valence electrons. Core electrons are these closest to the nucleus and valence electrons we've talked about before in previous videos, feel free to go back and check those out. But these are the outermost and for these lectures were really only considering the outermost s and p, d and f electrons behave a little strange. Really we're just concentrating on these outermost s and p electrons. These are the furthest ones from the nucleus.
1:00
They're the ones responsible for forming those chemical bonds because as atoms get close to other atoms, the furthest out electrons are the first ones to meet up and approach the other atoms for this out electrons. Overall, we have this octet rule, octet meaning eight. This rule states that all elements want to achieve this octet configuration and they'll do so by transferring electrons from one atom to another or sharing electrons between atoms. Transferring is ionic,
1:35
sharing electrons is covalent. These are the two types of bonding we're going to talk about today. First, ionic. That's what we see here. One atom is giving an electron to another atom. They're not shared, but a total transfer of electrons. That's the nature of the ionic bond. Now, what forms ionic bonds? We've learned previously, it's usually elements on the left combining with elements on the right and that's because these metals tend to lose electrons to approach their octet or
2:10
noble gas configuration and these elements on the right tend to gain electrons becoming anions to achieve their noble gas configuration and this tendency is described by the octet rule. These halogens over here are just one step away from becoming a noble gas and achieving this octet configuration. They all need one more electron to move to the right one spot and become a noble gas. These first column alkali metals need to lose one electron moving one to the left,
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which takes them back to the right side of the periodic table over here, moving backwards, for instance, from lithium to helium, achieving their octet. That's these electrons being added or removed to make these elements achieve their noble gas configuration. Here we've written out the electron configuration for magnesium and when it loses these two s electrons, its configuration now for Mg^2+ is 1s^2, 2s^2, 2p^6. In other words, identical to this neon noble gas.
3:19
That's what all of these elements are trying to do. These metals tend to lose electrons and become cations so you can have something like sodium will lose an electron and becomes sodium plus and nonmetals, maybe chlorine, this halogen gains one electron to become Cl^- and it's really this transfer of the electron from sodium to chlorine. Sodium gets rid of its electron, but that electron goes somewhere. It goes on to chlorine. The bond that is formed is then this attraction between sodium
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plus and chlorine minus and this type of attraction is not the same, but we can think of it as the same attraction between magnets. These opposites attract, there's positive charge and there's negative charge. That is what forms the chemical bond. They transferring this electron separates these charges and now I have a plus charge that's going to be attracted to a minus charge and it turns out we can arrange a bunch of these and pair them up. In this fashion, there's now
4:29
two sodium chloride molecules that are forming this square or this lattice, where all these charges benefit by being attracted to their neighbors. Now, in Chapter 5, earlier videos, we learned about trends in atomic radii. How big is an atom? Well, it depends what atom you're talking about. As you go down the periodic table, they get larger. Hydrogen to cesium. Cesium is really large. As you go left to right, they get smaller. Helium, smaller than hydrogen.
5:01
But what happens when we add or remove electrons? Add or remove electrons, I should say. If we're taking sodium, down here with this sphere, and I remove an electron. Well now this three s electron that sodium had, I've removed and so I don't have a third energy level because the third energy level and sodium was just that one electron removing that one electron removes the whole energy level and so now I only have two energy levels so my atom is smaller.
5:29
Remember these energy levels we can think of as orbits. One s maybe the second energy level is out here, maybe the third energy levels out here. Removing this third energy level, it makes the atom overall smaller. As things get positively charged or lose electrons, they become smaller and the opposite trend is true and anion gains electrons and that's larger than it starting atom. There's something like chlorine that is maybe this size. When it gains an electron,
6:00
it gets bigger, just a little bit, but it gets larger.
6:07
We can have these charged particles here's sodium. Here's chlorine, it's larger. Why? Because chlorine is now a negative ion, it's gotten bigger. Sodium is a positive ion. It has gotten smaller and chlorine is further to the right on the periodic table, so it's larger than sodium anyways. Now we can arrange these things plus, minus, plus, minus into this what's called a lattice or crystal structure. The bonds shown in red here are very strong and form this really rigid structure.
6:43
This is what it looks like at the molecular level or the microscopic scale here. All of these ions forming this long extended network solid and so you can imagine sodium chloride, the table salt you used to flavor French fries. At the atomic level, it looks like these little cubes of positive and negative charges. Let's take a moment and try some quiz questions about this ionic bond concepts. First, let's try to answer this one. You can pause the video here and try to answer it yourself.
7:15
I'm going to proceed to solve it. How many valence electrons are in a N^3- ion? Valence, remember, is the outermost s and p. We can try to write the configuration of nitrogen, 1s^2, 2s^2, 2p^3 and if you're confused about this, go back and watch my previous video on electron configurations. This three minus means there's three extra negative charges, three extra electrons and so this 2p^3, when I add three more electrons, I add them here because a p sublevel can hold six electrons.
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But now this is for N^3- 1s^2, 2s^2, 2p^6. What are the outermost s and p? 2s and 2p, s has two, p has six. 2+6 is eight. This N^3- ion 1, 2, 3 has filled up its octet. It has eight valence electrons. Let's try another problem. Which of the following does not contain an ionic bond? Again, pause the video here. Take 30 seconds or so and try to answer this one on your own.
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Ionic bond is going to be something between a cation and an anion and cation ions are generally forming from these metals like this, row hydrogen is not a metal. The alkaline metals, the alkaline earth metals, aluminum, gallium, any of these transition metals combining with one of these non-metals over here, or hydrogen. Let's look at this magnesium metal. Nitrogen non-metal, it'll form an anion. This will be an ionic bond. Sulfur that's over here.
9:13
That's a non-metal, oxygen nonmetal that will not be an ionic bond. Calcium metal, fluorine nonmetal, that will be an ionic bond. Aluminum metal. It's over here on the right, but it's still a metal. It's still left of this staircase. The semi-metals that separates the metals from the nonmetals and oxygen a non-metal. This is ionic as well. Which does not contain an ionic bond? That would be this one. What type of bond is in this sulfur dioxide molecule?
9:44
It's covalent, and that's the next type of chemical bond we're going to learn about. Just to summarize this section, ionic bonds form when metals lose electrons and become cations and they transfer that electron to a non-metal, which becomes an anion and then the cation and anion are opposite charges which are attracted to one another and form this ionic bond. It's this transfer of electron. That's different than covalent bonds shown here, where you're sharing electrons.
10:13
Here you're both putting some electrons into some common pool that both atoms now have access to. That's the nature of covalent bonds. This is when two or more non-metals share electrons. The sharing really only happens between elements that are nonmetals, these over here, plus hydrogen. They're the ones that are able to share these electrons and form covalent bonds. For example, the bond between hydrogen and chlorine, both are non-metals, that is a covalent bond.
10:44
Hydrogen shares its one electron with chlorine, giving chlorine eight electrons. Remember chlorine is over here, it has seven valence electrons, it wants eight, so hydrogen's able to give it to that one electron, which becomes eight. Likewise, fluorine is sharing one of its seven electrons with hydrogen, which gives hydrogen two electrons. Now most things for the octet rule, we assume it's eight valence electrons. Hydrogen, though it's special, it only needs to move one place to get to
11:12
helium to have two valence electrons for its full shell. Hydrogen doesn't need eight electrons for a full shell, it just needs two, it needs its 1S energy sublevel to be filled. Hydrogen just needs two. It has one, so if it can use one of chlorine's, they can get two and achieve this noble gas configuration. Whereas chlorine is trying to achieve its noble gas configuration of argon. Chlorine has seven, it uses hydrogen's one electron to make eight. Simultaneously, hydrogen is borrowing one of
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fluorines shared electrons to get two valence electrons, making it isoelectronic, same number of electrons as helium. Now both things are happening to the hydrogen here and the chlorine here now have happy full valence shells. That's the nature of the covalent bond. Now, the distance between two atoms that are covalently bonding will be smaller than the sum of the radii. That's a lot of technical words. All it means is that if I look on the periodic table,
12:17
hydrogen is this size and chlorine is this size. When they bond, they morph into one another. The size is not going to be hydrogen plus chlorine size, it's going to be something smaller. Here's the size that I drew, and it's definitely smaller here. That's the real point of this slide, is just knowing that when these things share electrons, they are hugging it out here, getting closer to one another. We call this distance between the covalently bound atoms,
12:50
the bond length, and this is a term we'll often use. The chemical bond length and hydrochloric acid, HCl seen here, is this small length between the center spots of these two atoms. That's going to be smaller than if you just stack these two spheres on top of one another since they morph into one another when they're sharing their electrons. Now, you can form different numbers of bonds or types of bonds, single, double or triple. This depends on how many electrons you are sharing.
13:21
For the hydrogen chlorine example, hydrogen shared one, chlorine shared one, each sharing one means it's a single bond. But for other types of elements, you might need two electrons to get to your noble gas configuration or octet, or you might need three. But you have to find a sharing partner that is capable of giving you two or three electrons. You can't form a double bond if there's two pairs of electrons that are being shared or four electrons total, two from one atom, two from another,
13:54
so now you have four electrons total. Each atom is sharing one pair for a total of two pairs of electrons, so that's a double bond. Remember single bonds like HCl, there's two electrons total. Double bonds, there's four electrons total. Triple bonds, now, each atom is contributing three to this shared pool for a total of six electrons and it's all about these electron pairs. One pair is single, two pair, double, three pairs, six electrons, a triple bond.
14:24
It turns out the more electrons that atoms give to this pool forming double or triple bonds, the stronger that bond becomes, so a double bond is stronger than a single bond, a triple bond is stronger than a double bond because there's more of these shared electrons that forms this stronger attraction. This stronger attraction also makes the atom sort into one another more and more. Here's a single bond. If this was a double bond, these two would morph into each other even more,
14:58
and this bond length would get shorter and shorter. That's what's happening here. The stronger the bond, the more shared electrons, the shorter that bond becomes. Now, covalent bonds are sharing. Ionic bonds were about complete transfer. Turns out though the sharing is not always equitable or equal, covalent bonds that are nonpolar are completely equal. But most covalent bonds are actually polar, meaning there's not an equal sharing.
15:30
Although they're shared, it's not necessarily fair to both atoms. Here's HCl again, hydrogen and chlorine. One atom may hold onto that shared pool of electrons more. It might pull the electrons towards itself more strongly, and so we call these polar covalent bonds. The electrons are shared, but not equally. Here it's showing oxygen attracting electrons more strongly than hydrogen.
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This ability of the atom to attract electrons to itself and hog some of these shared electrons is called electronegativity. This governs how unequal that sharing will be in the covalent bond. On this periodic table here, I've written down some of these electronegativity values. For hydrogen, it's 2.1. For oxygen, it's 3.5. The larger the number here, the more the electrons are attracted to that element or the more strongly that element holds onto electrons.
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Oxygen's number 3.5 is higher than hydrogen's number 2.1. When H and O get together and share electrons, oxygen hogs most of those electrons. This is a polar covalent bond because these numbers aren't equal. They want the electrons different amounts and oxygen will hog some of those electrons a little bit more tightly to its chest than hydrogen will, so this is a polar covalent bond. We'll generally signify this with these Delta symbols. Now that we have some idea of looking at
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this periodic table and the electronegativity values and figuring out, okay, oxygen wants the electrons more, we generally associate that with this Greek lowercase Delta symbol. The one that has the higher electronegativity value, like 3.5, gets a negative sign. It's hogging the electrons. Hydrogen is a positive sign. Because in that shared pool of electrons, it doesn't really have the electrons very often so most of the time it's slightly positive,
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whereas the oxygen is slightly negative because it has those negative charge electrons more often. Here, this slide is comparing hydrogen and chlorine. Chlorine is higher electronegativity wise, so it hogs the electrons. You can imagine this as the electrons that are orbiting the atoms might spend more time around chlorine than they do the neighboring hydrogen. Both atoms have access to the electrons, but they're spending more time closer to chlorine than they are hydrogen.
18:12
We get this slightly negative charge on chlorine that's shown here with this Delta signal or symbol at the bottom as opposed to hydrogen, we use this Delta positive symbol. The higher electronegativity value, it's delta minus, the lower number, it's Delta plus. We're comparing the things that are bound to one another forming this polar covalent bond. Nonpolar covalent bonds will form if the two atoms have the same electronegativity value. This will always be the case in a diatomic.
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Homonuclear, same nucleus, diatomic to be specific. Basically when the same atom like H2, when the same atom is bound to itself, that is a covalent bond because these are both nonmetals, but it's non-polar because both have the same 2.1 electronegativity value. Neither the left hydrogen or the right hydrogen wants the electrons any more than the other, so it's a completely fair equitable situation. Both atoms equally want the electrons, so it's non-polar.
19:25
The electrons are not spending time around the right hydrogen anymore than the left hydrogen. We call these nonpolar covalent bonds. Here's the one shown for fluorine.
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Let's stop here with covalent bonds and try to answer this quiz question. Pause the video here and try to answer it for yourself.
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Using this Delta notation indicates the partial charges in a P-F molecule. Here's P, its bound to F. What are the partial charges? First, what kind of bond is it? P is over here, F is over here. Both are non-metals. This is definitely going to be covalent. Now, is it polar covalent or non-polar covalent? Well, 2.1 versus 4.0. Fluorine is the most electronegative element on the periodic table and also in this compound, this P-F bond, so fluorine is going to get this Delta negative,
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whereas phosphorus is going to get this delta positive. Most of the time the electrons in this bond that is shared between these atoms will be spending time with fluorine closer to it. That does it for this lecture. I'll present another lecture where we go over Lewis dot structures and molecular geometry as part of this same unit here, and also a practice problem video. As always, if you'd like these videos, feel free to subscribe to my channel on YouTube and you can
21:01
also find me over on Twitch where I do office hours every week answering whatever questions you have about chemistry. That's all for this video. I'll catch you next time.
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